PASADENA CITY COLLEGE

CHEMISTRY 22EXPERIMENTS AND STUDY ASSIGNMENTS

Experiment 1Experiment 2Experiment 3Experiment 4Experiment 5Experiment 6Experiment 7Experiment 8Experiment 9Experiment 10Experiment 11Experiment 12Experiment 13Experiment 14Experiment 15Experiment 16Experiment 17Experiment 18

Study Assignment 1Study Assignment 2Study Assignment 3Study Assignment 4Study Assignment 5Study Assignment 6Study Assignment 7Study Assignment 8Study Assignment 9Study Assignment 10Study Assignment 11Study Assignment 12

Measurements DensitySeparation of Mixtures Salt/Sand Separation Paper ChromatographyChemical and Physical Changes Chemical Reactions and Equations Empirical FormulaHydrated Salts Stoichiometry Hydrogen Gas Oxygen Gas Charles' Law Ion Identification Solutions TitrationAcids and Bases Preparation of Ethanol

Significant Figures/Math Review Unit ConversionsDimensional AnalysisElectron Configurations/ Periodic Properties Ionic CompoundsVSEPRShapes and Polarity Nomenclature/ Formulas Balancing EquationsReaction Types/ Writing Equations Solutions ProblemsNet Ionic Equations

Appendices Periodic TableList of Ions to KnowSolubility Table, Vapor Pressure Table

2019

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CHEMISTRY 22Experiment 1, Introduction to Laboratory Equipment

Introduction

The purpose of this experiment is to become more familiar with the metric system, learn how to use the centigram balance, practice conversions, learn more about Bunsen burners, and in some cases practice glass manipulations.

Procedure

PART 1. Length

A. Measure the length of this lab page in centimeters using a meter stick. The measurement should have two places after the decimal. Convert the measurement into mm and km. Also convert the length to inches and feet. Make sure you show the proper conversion factors.

B. Measure the length of a large test tube in centimeters using a meter stick. Record the length to the nearest 0.1 cm. Convert this measurement to meters and inches.

PART 2. VolumeFill a "250 mL" beaker to the very top with top water. Using a graduated cylinder, measure the

volume of water contained in the beaker to the nearest 0.1 mL. Do this by carefully pouring the water into a 50 mL graduated cylinder. Measure the volume of the water in the cylinder. You do NOT have to fill the graduated cylinder to exactly 50.0 mL each time. Repeat this until no more water is left in the beaker. Repeat this with a 10.0 mL graduated cylinder to find the volume of a large and small test tube. Convert the actual volume of the 250 mL beaker to liters, quarts, and gallons. Convert the actual volume of the large test tube to liters, quarts, and microliters.

PART 3. MassUsing the balance, find the mass of 5 pennies to the nearest 0.01 g. Calculate the average mass of

one penny. Convert the average mass to pounds.

PART 4. Bunsen Burner

A. Draw a sketch of a Bunsen burner, which is used for heating items. Several types of Bunsen burners exist, but each one has a gas inlet, a barrel for mixing air and gas, an opening near the base of the tube for introduction of air (air entrance valve), and when lit, a flame with two regions (an inner and outer cone). Find these parts on your burner and label it on your sketch.

B. Light the burner by partially closing the air vent, turning on the gas part way, and lighting the burner with either a match or a striker. Close the air hole at the base of the tube. Using crucible tongs, hold an evaporating dish about 5 cm above the top of the tube.

C. Open the air hole about half way and adjust the flame with the gas until the flame is about 8-10 cm high. The flame should be blue. Figure out the hottest part of the Bunsen burner flame by placing a nichrome (an alloy of nickel, chromium, and iron) wire in various positions within the flame and seeing how long it takes for the wire to glow red. Using tongs hold the end of a piece

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of thin nichrome wire and first place it into the flame so that it is resting on the top of the metal barrel. Then place it inside the inner blue cone, just above the inner cone, and at the tip of the whole flame. Record how much time it takes for the wire to glow red in each position. If it takes longer than a minute for the wire to tum red, just write down longer than 1 minute.

D. Hold first a copper wire and then a nichrome wire in the hottest part of the flame for 20-30 seconds. After taking the wire out, determine if each one has melted by seeing if the wire has changed shape at the end that was inside the flame. At lease one of the wires will melt. If neither of the wires melt, you need to adjust your flame to make it hotter.

PART 5. (Optional) Glass ManipulationsPerform as many of the following glass manipulations as are requested by your instructor.

A. Cutting glass tubing. At the point where you wish to cut the tubing, make a deep scratch by placing the tubing on a flat surface and draw a triangular file firmly across the glass once. Pick up the tubing with both hands and grasp it firmly so that the thumbs are opposite the scratch, then firmly and quickly press the thumbs against the tube and at the same time pull back with the fingers so as to break the tubing at the file mark. Practice cutting glass using scrap pieces of tubing.

B. Fire polishing. Newly-cut edges of glass are very sharp and should be fire-polished before being inserted into a rubber stopper or otherwise used. Fire polishing is accomplished by rolling the end of the tube in the burner flame to heat uniformly, until the glass softens and the sharp edges become rounded. CAUTION: Place hot glass on a wire gauze to cool, not directly on the bench top. BE CAREFUL NOT TO HANDLE THE HOT GLASS.

C. Bending tubing. Using a 12 to 15 cm piece of tubing, hold the tubing in the flame where it is to be bent, heating about 5 cm of its length. Rotate the tubing constantly to heat all sides uniformly. When the glass becomes noticeably soft, remove it from the flame and slowly bend it to a right angle. Fire polish the ends.

D. Drawing glass tubing. Use a short length of tubing (about 10 to 12 cm). Hold the center of the piece of tubing in the hottest part of the flame. Rotate the tubing as it is being heated until it is very soft. Remove it from the flame and pull the two ends apart until the softened region is as small as desired. When cool cut to the desired length and fire polish both ends, being careful not to seal the jet end. You have just made a dropper!

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CHEMISTRY 22 NameExperiment 1, Introduction to Laboratory Equipment Report Sheet

PART 1. LengthA. Length of this lab page in centimeters cm

Length of this lab page in millimeters using scientific notation mm Show work (include conversion factors):

Length of this lab page in kilometers using scientific notation km Show work (include conversion factors):

Length of this lab page in inches m Show work (include conversion factors):

Length of this lab page in feet ft Show work (include conversion factors):

B. Length of a large test tube in centimeters cm

Length of a large test tube in meters mShow work (include conversion factors):

Length of a large test tube in inches m Show work (include conversion factors):

PART 2. VolumeVolume of 250 mL beaker Volume of Large Test

TubeVolume of Small Test Tube

Volume 1 mL Volume 1 mL Volume 1 mLVolume 2 mL Volume 2 mL Volume 2 mLVolume 3 mL Volume 3 mLVolume 4 mL Volume4 mLVolume 5 mLVolume 6 mL

Total Volume mL Total Volume mL Total Volume mL

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Total volume of 250 mL beaker in liters-----------------------------LShow work (include conversion factors):

Total volume of 250 mL beaker in quarts qt Show work (include conversion factors):

Total volume of 250 mL beaker in gallons gal Show work (include conversion factors):

Total volume of large test tube in liters L Show work (include conversion factors):

Total volume of large test tube in quarts qt Show work (include conversion factors):

Total volume of large test tube in microliters µL Show work (include conversion factors):

PART 3. MassMass of 5 pennies g

How many significant figures are in the mass of 5 pennies? _

Average mass of one penny in grams g Show work:

Average mass of one penny in pounds lb Show work (include conversion factors):

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PART 4. Bunsen BurnerA. Sketch of Bunsen burner. Label the base, gas flow, air entrance valve, barrel, flame showing

inner and outer cone.

\

B. What is the appearance of the flame when the air intake is covered (make sure to mention the color)?

What is the deposit on the bottom of the evaporating dish?

Where did the deposit come from?

C. Time (in minutes or seconds) for the nichrome wire to glow red:Resting on the top of the metal barrelInside the Inner blue coneJust above the inner coneAt the tip of the whole flame

The hottest region of the Bunsen burner flame is the position where it took the shortest time for the wire to glow red. From your data, where is the hottest portion of the flame.

Make a sketch of the burner flame showing both the inner and outer cone. Label the hottest portion of the flame, the inner cone, and the outer cone.

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D. Did the copper wire melt? (circle one) Yes No

Did the nichrome wire melt? (circle one) Yes No

Melting point of copper (from handbook) _

Melting point of nichrome (from handbook) _

Based on which wire(s) melted and the melting points of copper and nichrome, what is the approximate temperature range of the Bunsen burner flame?

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D

CHEMISTRY 22Experiment 2, Density

Introduction

Density is a physical property of matter. It is defined as the mass of a substance divided by itsvo1ume. The mathematl.ca1 equat1.0n .IS:

dens1.

ty

mass= --- or = -

volume VTo experimentally determine the density of a substance you must measure both the mass and the

volume of the same sample, then you can divide the mass by the volume to compute the density. Density is always expressed in units of mass per volume, such as g/cm3 or g/L.

The mass of a sample is measured using a balance. How you measure the volume of a sample depends on the nature of the substance. If it is a liquid, the volume is measured using a graduated cylinder, pipet, buret, or volumetric flask. If the sample is a solid with a regular geometric shape, such as a block of wood, a ruler can be used to find the linear dimensions and then the volume can be computed using the appropriate geometric formula. If you have an irregularly shaped solid object, you can measure the volume by the displacement of water (or any other liquid). To do this, you put some water into a graduated cylinder and record the volume. Then drop the object into the water (it must sink) and record the new volume. The volume of the object is equal to the difference between the two volume readings because the object displaces a volume of water equal to its own volume.

In this experiment you will measure the density of four different things: water, an unknown liquid, a regularly shaped solid, and an irregularly shaped solid. After you measure the mass and volume and compute the density, you can get the correct density from your instructor. You can then compare how close your density is to the accepted value by calculating the percent error:

% error= ( experimental valu-e accepted value)( lOO%)accepted value

Procedure

PART 1. Density of Water.Weigh a m:y 50 mL graduated cylinder to the nearest 0.01 g. (Record this mass, and all weights and

volumes you measure, in the spaces provided on the lab report sheet as soon as you make the measurement.) Remove it from the balance and add 40 to 45 mL of distilled water. Record the actualvolume of water used to the nearest 0.1 mL. Weigh the graduated cylinder with the water in it. Calculate the density of the water. Measure the temperature of the water and look up the correct value for the density using the table on the next page. Calculate your percent error. (Note: express your percent error as a positive number, even if your answer comes out negative.)

PART 2. Density of an Unknown Liquid.Check out a 10 mL pipet from the stockroom. Take a clean, dry, large test tube with your name on it

to your instructor to obtain a sample of an unknown liquid.Weigh a clean, dry crucible with the lid to the nearest 0.01 g. Remove the crucible from the balance

and use the pipet to deliver precisely 10.0 ml of the unknown liquid into the crucible. Replace the lid and weigh the crucible with the liquid in it. Calculate the density of the liquid. You can dispose of your unknown liquid in the sink.

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PART 3. Density of a Rectangular or Cylindrical Solid.Obtain a regularly shaped solid sample from your instructor. Weigh it to the nearest 0.01 g. Use a

ruler to measure the linear dimensions of your solid and calculate the volume of the solid using the appropriate geometric formula. Calculate the density of the solid. Obtain the correct value of the density from your instructor and calculate your percent error.

PART 4. Density of an Irregular Solid.Obtain a sample from your instructor. Weigh an empty beaker or evaporating dish to the nearest 0.01

g. Place all the pieces of the sample into this container and record the mass. Place about 35-40 mL of water into your 50 mL graduated cylinder. Record the initial volume to the nearest 0.1 mL. Carefully slide all the pieces of your sample into the water in the graduated cylinder. Make sure they all sink and the volume is still on the scale. If not, you may have to repeat, adjusting the initial amount of water to get it to work. Record the volume reading for the water plus the sample. Calculate the density of the sample. Dry off your sample with paper towels and return to your instructor. Obtain the correct value of the density from your instructor and calculate your percent error.

TABLE OF THE DENSITY OF WATER AT DIFFERENT TEMPERATURES

Temperature (°C) Density (g/mL)0 0.99994 1.0000

15 0.999116 0.998917 0.998818 0.998619 0.998420 0.998221 0.998022 0.997823 0.997524 0.997325 0.997126 0.996827 0.996528 0.9962

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CHEMISTRY 22Experiment 2, Density Report Sheet

PART 1. Dens1·ty of Water

Name ------------

PART 2. Density of Unknown Liquid #mass of empty crucible plus lid g

mass of crucible plus lid plus liquid g

mass of liquid g

volume of liquid mL

density of liquid (show work)

g/mLaccepted value of density (from instructor) g/mL

percent error (show work)

%

mass of empty, dry graduated cylinder g

mass of graduated cylinder + water g

mass of water g

volume of water mL

density of water (show work)

g/mLaccepted value of density (from table on page 2-2) g/mL

percent error (show work)

%

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PART 3. Density of Rectan2:ular or Cylindrical Solid #

PART 4. Density of Irregular Solid #

dimensions of rectangular solid

length

width

height

cm

cm

cm

or dimensions of cylinder (V = nr41)

length

diameter

radius

cm

cm

cm

volume of solid (show work)

cm3

mass of solid g

density of solid (show work)

g/cm3

accepted value of density (from instructor) g/cmj

percent error (show work)

%

mass of container g

mass of container plus sample g

mass of sample g

initial volume of water in graduated cylinder mL

volume reading after putting in sample mL

volume of sample mL

density of sample (show work)

g/mLaccepted value of density (from instructor) g/mL

percent error (show work)

%

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CHEMISTRY 22Experiment 2, Density Additional Problems

For all of the problems you must show a neat, complete and logical method of calculation. Use the conversion factor (dimensional analysis) method and make sure all numbers are labeled with the appropriate units. Pay attention to significant figures and put your answer in the box provided.

1. 315 g of a metal occupy a volume of 41.0 mL. What is the density of the metal?

2. A liquid has a density of 1.45 g/mL. What is the mass of 2.5 L of this liquid?

3. How many mL of water would have a mass of 6.600 kg at 25°C? (Use the table on page 2-2 for the precise density.) You may express your answer in scientific notation.

4. A spherical ball bearing has a diameter of8.5 mm and a mass of2.315 g. What is the density of the

ball bearing (in g/cm3)? The formula for volume of a sphere is V =± 3

.n;r 3

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5. A rectangular piece of plastic has a width of 4.2 cm, a height of 1.9 cm, and a mass of 64.6 g. If thedensity of this plastic is 0.92 g/cm3 what is the length of the piece?

6. What is the mass, in pounds, of 1.00 cubic feet of lead? The density of lead is 11.3 g/cm3•

7. You have three beakers, each filled with a different liquid at 20°C. One beaker contains hexane, which has a density of 0.66 g/mL; one beaker contains carbon tetrachloride, which has a density of 1.59 g/mL; and the other beaker contains water. A rubber stopper is dropped into all three liquids. The rubber stopper sinks in the hexane and the water, but it floats in the carbon tetrachloride. What is the best conclusion you can make about the density of the rubber stopper?

,

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CHEMISTRY 22Experiment 3, Separation of Mixtures

Introduction

A mixture can be either heterogeneous or homogeneous. A homogeneous mixture is uniform throughout and therefore has only one phase. For example, salt dissolved in water is a homogeneous mixture. A heterogeneous mixture is non-uniform and can have different phases. A mixture of oil and water is an example of a heterogeneous mixture. In mixtures, unlike compounds, the components are not chemically linked together and therefore can be easily separated. The separation makes use of differences in physical properties such as particle size, solubility, boiling point, density.

For example, sand can be removed from the sand water mixture by filtration. The large sand particles will not pass through the pores of the filter paper but the very tiny water molecules will pass through. On the other hand, chlorophyll, the substance responsible for green color of leaves and photosynthesis, cannot be separated from the rest of the leaves components by filtration. They are all very tiny and pass through the pores of the filter paper. However, chlorophyll can be separated (extracted) from the rest of leaf because it is soluble in cyclohexane whereas the other components in the leaf are not soluble in cyclohexane. Therefore, this separation method (extraction) makes use of the differences in solubility.

The alcohol in beer can be concentrated by distillation because it has a lower boiling point than water, therefore, the distillate (the liquid that distills over) will contain more alcohol molecules in it than beer. Salt water can be separated into water and salt by distillation as well. The salt molecules are not volatile (do not easily turn into gas) and in the distillate there will only be water and no salt.

The components of a compound, however, are chemically bonded together and can only be separated into elements by a chemical change (breaking and reforming of the bonds).

In this experiment we will examine separations of mixtures based upon differences in volatility, particle size, density and solubility.

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Procedure

Part I. Distillation

A. Obtain the distilling apparatus from the stockroom. Wash the flask and condenser thoroughly. Set up the flask and condenser as shown in the Figure 3-1.

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In a beaker dissolve one spoon full of solid colored salt in about 100 mL of distilled water. Transfer the solution to the distilling flask using a long stem funnel. Make sure that none of the solution gets into the side arm of the distilling flask. Start water flowing through the condenser. Heat the distilling flask until the solution begins to boil gently. Distill over about 10 mL of the solution. The solution that distilled over is called the distillate. Check the distillate to see if it is blue. If the distillate is colorless then it does not contain any of the colored salt. However, if the distillate is blue then you can assume that it does contain the colored salt. Dispose of the solution and the distillate in the sink.

B. Clean out the distilling flask by rinsing it several times with water, then doing a final rinse with distilled water. Clean the condenser and all the equipment used by rinsing them thoroughly with water. Place about 100 mL of water and about 10 mL of 6 M ammonium hydroxide (NH4OH) solution in the distilling flask (using the long stem funnel as before). Distill over about 10 mL of this mixture. Test to see if the distillate contains ammonium hydroxide in it by adding 1-2 drops of phenolphthalein indicator to it. Phenolphthalein turns pink in presence of the base (NH4OH). Also test the solution in the distilling flask to see if it contains ammonium hydroxide by adding 1-2 drops of phenolphthalein to about 10 mL of it. Clean the distilling apparatus and return it to the stockroom.

Part II. Filtration

Place a filter paper in a funnel by folding it according to your instructor's instructions. Place the funnel into a clay triangle on a ring stand. Then place a beaker below the funnel to collect the solution that filters through the paper. Mix about a eighth of a spoon full of the colored salt and about the same amount of solid ferric oxide (Fe2O3) , note the color of each, in a 150 mL beaker. Add about 20 mL of distilled water. Stir the mixture for several minutes to make sure that the soluble material has dissolved. Pour the mixture into the filter paper in several portions, stirring before each addition to re-suspend the material in the mixture. Never fill the filter paper more than 2/3 full. After the filtering is complete, examine the residue on the filter paper and the solution that filtered through the paper (which is called the filtrate). The filter paper can be placed in the waste container provided. The filtrate can be washed down the sink. If any ferric oxide is in the sink, please rinse it down the drain.

Part III. Centrifugation

The centrifuge spins samples very rapidly. The force created due to the spinning is greater than the force of gravity. Therefore, the more dense particles settle out more rapidly than they would due to force of gravity.

Place about 10 mL of 0.1 M ferric chloride solution (FeCh) in a small beaker and add about 5 mL of 6 M ammonium hydroxide solution. Stir the mixture. The chemical reaction will produce insoluble ferric hydroxide (which is called the precipitate). Observe the mixture for a few minutes and note how the precipitate will slowly settle. Stir the

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mixture again and pour the mixture into two small (10 cm) test tubes until each tube is about 2/3 full. Place the two tubes(with no stoppers) on opposite sides of the centrifuge and spin for about a minute. Remove the tubes and observe.

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CHEMISTRY 22Experiment 3, Separation of Mixtures Report Sheet

Part I

A. Did any of the colored salt distill over?

Name--------

Yes No

What substance did you see distill over?

Can water and the colored salt be separated from each other by distillation?

Yes No

B. Did any of the NH4OH distill over?

Were the ammonium hydroxide and water separated by this distillation?

Part II

Can the colored salt and water be separated from each other by filtration?

Yes

Yes

No

No

Yes No

Can ferric oxide and water be separated from each other by filtration? Yes No

What evidence did you see to support your answers to the above questions?

Part III

What is the color of the substance at the bottom of the test tube?----

What is the substance at the bottom of the test tube?

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Additional Questions:

I) What is the best method for separating a mixture of insoluble solid silver iodide (Agl) and water? Explain Why. (Insoluble in water means that it does not dissolve in water)

2) What is the best method for separating a mixture of potassium iodide dissolved in water? Explain why.

3) Explain how one can separate a mixture of sand and salt from one another.

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CHEMISTRY 22Experiment 4, Quantitative Separation of a Mixture

Introduction

The purpose of this lab is to better understand mixtures. Your goal is to calculate the percentages of salt (NaCl) and sand (SiO2) in a mixture. Since the components of a mixture are not chemically bonded to one another, they can be separated by physical processes such as evaporation of water. For instance, pure salt may be obtained from salt water by heating the mixture so that the water evaporates (physical change). The salt has a higher boiling point than water so it does not evaporate at the temperature produced by heating a sample with a Bunsen burner.

In this experiment you will first add water to your mixture of salt and sand. The water will dissolvethe salt but the sand will not dissolve in water. The salt water can then be poured off leaving the sand behind. The sand is then dried to remove any water adhering to the sand. The water is also removed from the salt water solution through heating.

Since you must calculate percent of salt and sand in your mixture, it is important to understand percents. Percentages are used to compare quantities. A percent represents a part of a quantity over the whole quantity. For example, if a parking lot contains 5 red cars and 15 other cars, 25% of the cars are red:

percent = part--x 100%

=

5 red cars----x100% = 25%

Procedure

whole 20 total cars

Part 1. Obtain a sample of a mixture of salt and sand from your instructor.Weigh a clean, dry evaporating dish to the nearest 0.01 gram and then pour all of the sample into

it. Take the mass of the dish and the sample to the nearest 0.01 gram. Add about 20 mL of distilled water to the evaporating dish and let this stand for a few minutes, stirring occasionally. Be careful not to spill any of the contents. If the stirring rod is removed at any time, it should be rinsed off into the evaporating dish with a few drops of water from a wash bottle to avoid any loss of sample. After the last stirring, allow the sample to settle. Weigh a clean, dry 150-mL beaker (this mass will be needed for Part 2 so you will record it in Part 2A). Carefully pour off the water-salt solution into the pre weighed beaker. Be careful not to transfer any of the sand to the beaker. This process of pouring off the liquid from a solid is called decantation. Wash the sand residue two more times by adding 6-8 mL of water, stirring and decanting, adding each washing to the first one in the 150-mL beaker. Save the salt-water solution for Part 2.

Place the evaporating dish on a wire gauze and gently heat to evaporate the last of the water. If the heating is too rapid during this drying process the water left in the dish will spatter and throw small particles ofresidue out of the dish. After the sample appears dry, increase the heat until the dish is heated with full burner heat. Allow the dish and contents to cool to room temperature. When cool, weigh the dish and contents to the nearest 0.01 gram. It is impossible to be sure that any sample is dry from its appearance alone. Therefore, the dish and sand should be reheated for a few minutes, allowed to cool and again weighed. The only true test for dryness is if the dish and sand do not change in weight after this reheating. If the dish and sand become lighter after this second heating, the process of heating, cooling, and weighing should be repeated until a constant weight is obtained.

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Calculate the mass of sand and the percent of sand in the sample.

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Part 2.A. Heat (but do not boil) the salt-water solution in the 150-mL beaker on a wire gauze to evaporate

the water (boiling causes spattering and loss of sample). When nearly all the water has evaporated and the salt starts to separate out, slow down the heating to avoid spattering. After the salt appears dry, increase the heat gradually and finally heat with full burner heat for a few minutes. Allow the beaker and contents to cool to room temperature and weigh to the nearest 0.01 g. Heat, cool, and weigh again to be sure the salt is dry.

B. Calculate the percent of salt based on the amount of sand present in the sample. Remember that the sample contained only salt and sand. Thus, the difference between the mass of the original sample and the mass of the sand will tell you the mass of the salt.

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CHEMISTRY 22 NameExperiment 4, Quantitative Separation of a Mixture Report Sheet

PART 1. Percent of sand

Mass of evaporating dish

Mass of evaporating dish and sample

Mass of sampleCalculation:

Sample No.

Mass of dish and residue (sand) after washing and drying

Mass of dish and sand after redrying

Mass of sand Calculations:

Percent of sand in sample %Calculations:

PART 2. Percent of salt

A. Percent of salt based on amount of salt left in beaker.

Mass of beaker

Mass of beaker and salt after redrying

Mass of saltCalculations:

Percent of salt in the original sample from this data Calculations: ------------%

B. Percent of salt based on amount of sand in Part 1.

Mass of original sample (from Part 1)

Mass of salt based on amount of sand in Part 1 Calculations:

Percent of salt in sample Calculations:

True percent of salt in sample (from instructor)

------------%

------------%

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Problems:Show a logical, method of calculation for each problem.

1. If 15.00 grams of U.S. quarters contain 13.75 grams of copper, what is the percent copper in the quarters?

1.-----------

2. An alloy is made by mixing 35.0 grams of gold, 42.0 grams of nickel, and 51.5 grams of aluminum. What is the percent of each metal in the alloy?

2. -----------

3. Stainless steel used in forks and knives is an alloy containing 18 percent chromium, 10 percent nickel, and 72 percent iron. What total mass of stainless steel can be made from 50.0 grams of pure iron? (Assume excess chromium and nickel.)

3. -----------

4. Aluminum bronze contains 92.0% copper and 8.0% aluminum. What maximum mass of aluminum bronze can be prepared from 73.5 grams of copper and 42.2 grams of aluminum?

4. ----------

5. Valuable minerals (chemicals) can be obtained from ores, which are naturally occurring mixtures. One ore contains 38% magnetite, Fe3O4• Fe3O4 contains 72.4% iron. What mass of ore would contain 3.50 tons of iron (Hint: you do NOT need to convert between tons and grams)?

5. ----------

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CHEMISTRY 22Experiment 5, Paper Chromatography

Introduction

Chromatography is a general term for a family of separation techniques that are widely used in analytical chemistry. For example, gas chromatography (GC), coupled with mass spectrometry, is used to check for explosive compounds in airport security systems. High pressure liquid chromatography (HPLC) is used to analyze the purity of medicines. In this experiment you will use paper chromatography to separate the dyes that are used to color M & M candies.

In paper chromatography, the mixture to be separated is applied near the bottom of a piece of thick paper. A liquid solvent is allowed to run vertically up the paper by capillary action. As the solvent flows up the paper it carries the components of the mixture with it. But, depending on their chemical structures, different substances in the mixture will travel different distances up the paper in a given amount of time. Some substances will have a strong tendency to dissolve in the solvent and therefore will travel rapidly up the paper along with the solvent. Other substances will be more strongly attracted to the paper itself and less soluble in the solvent. These substances will cling tightly to the paper and will not travel very far with the solvent. In this way, the components of a mixture can be separated on the paper.

In this experiment you will use paper chromatography to look at the dyes used to color M & M candies. After developing your chromatogram you should be able to decide if the color of the candy is due to one single dye, or if it is colored by a mixture of dyes. You will also be able to tell which colors ofM & M candies contain yellow dye #5, a dye that can produce an allergic reaction in some individuals.

ProcedureCheck out a spot plate and a 1000 mL beaker from the stockroom. Put one piece of candy of each

available color in a separate well of the spot plate. Put one drop of yellow #5 dye solution in another one of the wells. Put one or two drops of water on each candy in the spot plate to dissolve the colored coating. When the colored coating has dissolved, but before the chocolate part is exposed, remove the candies from the spot plate.

(over)

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Obtain a piece of chromatography paper about 8 x 5 inches. Try to handle the paper by the edges to avoid getting fingerprints on it. Using a pencil, draw a line along the long dimension of the paper, about one half inch from the bottom. Lightly mark this line with 7 tick marks, about one inch apart. These marks are where you will apply the colored spots. Above each tick mark, near the very top of the paper, write in the candy color that will be applied there. One of the locations will be for yellow #5 dye. Here is an example of what your paper will look like (although your colors may be different):

Use a capillary tube to apply a small spot of the first color at the correct position along the pencil line (the origin). Keep the initial spot of color as small and compact as possible (3 mm in diameter or less). Use the other end of the capillary tube to apply the next color. Apply all the colors to their marks on the origin line, being careful to use a separate capillary tube end for each color. Dry the spots with the hair dryer provided. Once the spots are dry, you should repeat the application of the colors directly on top of the original spot. You want to do a total of four applications of color, drying the spots between each application and trying to keep the spots as small and compact as possible. For the yellow #5, only one application is needed because it is already a very concentrated solution.

Fold your chromatography paper into a cylinder (with the colors facing out) and staple the ends together without overlapping them. Your teacher will have an example of what this looks like. Put a small amount of 1 M NaCl solution into the 1000 mL beaker. The depth of the solution in the beaker should be such that the liquid will be below the level of the colored spots on your chromatography paper when you set the paper inside the beaker. Place the chromatography paper cylinder into the beaker, making sure the paper does not touch the side of the beaker. Let it sit on your bench top and watch while the solvent flows up the paper, carrying the colored dyes with it. Do not pick it up or jostle it while it is developing. Take the paper out when the solvent front is about 2 cm from the top of the paper. The solvent will continue to rise for several minutes while the paper is still wet.

When the progress of the solvent has stopped, remove the staples and set your chromatogram down flat on a clean paper towel. If your separated colored spots are very faint, you may want to circle the outline of the spots with pencil. Use your chromatogram to answer the questions on the next page.When you hand in your experiment, be sure to staple your .dry chromatogram to your lab report.

red brown green black orange blue yellow#5

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CHEMISTRY 22Experiment 5, Paper Chromatography Report Sheet

Name ------------

Questions (To be answered based on your chromatogram)

1. Is the green M & M colored by one single green dye or a mixture of blue and yellow dyes?

2. Is the orange M & M colored by one single orange dye or a mixture of red and yellow dyes?

3. Do any of the candy colors seem to be made up of the same set of dyes? Which ones?

4. Which candy colors appear to contain yellow #5 ?

5. Which color dye present in the candies seems to be the most soluble in the solvent?

6. Which color dye present in the candies seems to be the least soluble in the solvent (or the most attracted to the paper)?

Be sure to staple your dry chromatogram to the back of this report sheet before you hand it in.

6-1

CHEMISTRY 22Experiment 6, Physical and Chemical Changes

Introduction

Chemistry is the study of matter and the changes it undergoes. When matter goes through a change, we can classify the change as either a chemical change or a physical change. In a chemical change (also called a chemical reaction), the chemical composition of the matter changes and new substances are formed. For example, the burning of a match is a chemical change. The gases that you smell while it's burning and the ash that is left afterward are new substances. If new substances with a different chemical composition are produced, a chemical change has happened.

Ifthere is no change in chemical composition, then the change is just a physical change. For example, the melting of ice cubes is just a physical change. Even though there is clearly a change in state and properties, no new substance is produced. It is H2O in the solid state and still H2O in the liquid state. It is a physical change only.

In this experiment, you will do a series of changes and then classify them as chemical or physical changes. When you are trying to decide if a chemical change has occurred, remember to ask yourself the question, "do I see evidence of a new substance?" If so, the change is a chemical change.

Procedure

Note: Unless your instructor says otherwise, do this experiment in pairs to conserve chemicals.

PART 1. Get a small ribbon of magnesium metal from the chemical shelf. Examine it and describe its color, flexibility, and shininess on your data sheet.

Put an evaporating dish on your lab bench. Holding the ribbon of magnesium in your tongs, place the magnesium in the hot part of a bunsen burner flame. When it ignites, drop it into the evaporating dish.Do not look directly at the burning magnesium because it gives off ultraviolet light; look at it out of the sides of your eyes. Examine the product left in the evaporating dish and describe its properties on your data sheet.

PART 2. This part should be done as a demonstration by the teacher! The students should not be given iodine crystals. Put a few crystals of iodine in a clean, dry 150 mL beaker. Support the beaker on a wire gauze on a ring stand. Place an evaporating dish partially filled with cold tap water on top of the beaker. Gently heat the bottom of the beaker containing the iodine crystals. Stop heating when there is a noticeable deposit of solid on the bottom of the evaporating dish. Try not to breath in the iodine vapors.

Once the apparatus has cooled to room temperature, use a wooden splint to scrape some of the solid from the bottom of the evaporating dish into a large test tube. Avoid getting the solid on your hands (it is not particularly dangerous, but it will leave yellow stains). Add 1 mL of cyclohexane to the test tube and stir the solution with your wood splint. Compare this solution to a display test tube that shows what iodine looks like when dissolved in cyclohexane. Dispose of the cyclohexane solution in the waste bottle under the hood - do not put cyclohexane in the sinlc

Any iodine crystals remaining in the beaker or on the evaporating dish can be cleaned up with 0.1 M KI (potassium iodide) solution and washed down the sink.

6-2

PART 3. Add about one quarter spoonful of solid copper(II) carbonate (CuCO3) to a clean, dry, large test tube. Weigh the test tube plus the CuCO3 and record the mass on your data sheet. Note the color and appearance of the copper(II) carbonate.Heat the copper(II) carbonate powder in the test tube with a strong bunsen burner flame for about 2 minutes. Let it cool completely. Weigh the test tube plus contents and record the mass. Note the color and appearance of the substance in the test tube. You may dispose of this substance (which is cupric oxide) in the waste container labeled "CuO" or "Cupric oxide".

PART 4. Obtain a very small piece (no bigger than 3 mm across) of manganese (not magnesium) metal and put it in a small beaker. Add 3 mL of 6 M hydrochloric acid (HCl). Observe what happens. Allow the reaction to continue until all the manganese has dissolved (you may heat the beaker gently if the reaction is slow). Once the manganese has dissolved, take your beaker UNDER THE HOOD and carefully heat the beaker using your bunsen burner until all the liquid has evaporated and the residue appears dry. Examine the substance left in the dish.

PART 5. Obtain 5 mL of 1 M sodium hydroxide (NaOH) and 5 mL of 0.1 M ferric chloride (FeCh) in two separate large test tubes. Note the appearance of each solution. Now pour the solutions together and stir. Describe what you observe. To check the solubility of the brown substance, set up a piece of filter paper in a funnel like you did in experiment 3. Pour the mixture through the filter paper.

6-3

CHEMISTRY 22Experiment 6, Physical and Chemical Changes Report Sheet

Name -------------

PART 1.Describe the color, flexibility, and shininess of the magnesium metal:

Describe what you observe when the magnesium is ignited:

Is energy absorbed or given off (circle one) during this change?

Describe the appearance of the substance in the evaporating dish after the burning:

Is there evidence that a new substance was formed? (circle one) Yes No

What type of change did the magnesium undergo? (circle one) chemical physical

PART 2.Describe what you observe when the iodine is heated:

What is the color of the solution made by dissolving the solid from the bottom of your evaporating dish in cyclohexane?

Compare your solution to the display solution that shows iodine dissolved in cyclohexane. What chemical do you think the solid from your evaporating dish is?

Is there evidence that a new substance was formed during the heating? (circle one) Yes No

What type of change did the iodine undergo? (circle one) chemical physical

6-4

PART3.Mass of the test tube plus the copper(II) carbonate before heating g

Mass of the test tube plus contents after heating g

Describe the appearance of the copper(II) carbonate before heating:

Describe the appearance of the substance that is present after heating:

Is there evidence that a new substance was formed? (circle one) Yes No

What type of change happened when you heated the CuCO3? (circle one) chemical physical

Why do you think the mass was lower after heating? (hint: it has nothing to do with water or the sample drying out)

PART4.Describe the appearance of the manganese metal:

Describe what you see when the hydrochloric acid is added to the manganese:

After you evaporate the excess acid, describe the appearance of the substance remaining in the beaker:

Is there evidence that a new substance was formed? (circle one) Yes No

What type of change happened when you combined Mn and HCl? (circle one) chemical physical

PARTS.What is the color of the ferric chloride solution?

What is the color of the sodium hydroxide solution?

Describe what you see when you mix the two solutions together:

Does the brown substance dissolve in water or is it insoluble? (circle one)

soluble Is there

evidence that a new substance was formed? (circle one) Yes No

insoluble

6-5

What type of change happened when you mixed the two solutions? (circle one) chemical physical

7-1

CHEMISTRY 22Experiment 7, Chemical Reactions and Equations

Introduction

In this experiment you will be doing a series of chemical reactions. After observing the reaction, you will figure out what the products of the reaction are and write a balanced chemical equation. In order to do this, it is useful to classify chemical reactions into certain common types. Here are 5 common reaction types you should know.

1. Combination (synthesis) reactions. These are reactions in which simple substances combine to make a more complex substance. An example is when the two elements magnesium and oxygen combine to make a compound: 2Mg<sl + 0 2(gl - 2MgO<sl

2. Decomposition reactions. These are reactions in which a substance is broken down into simpler substances. Usually energy must be added to cause this to happen. An example is when mercuric oxide is heated and breaks down into its elements: 2HgO<sl 2 Hg0l + 0 <2 gl

3. Single Replacement reactions. These are reactions in which one element replaces another in a compound. An example is the reaction between copper metal and silver nitrate solution:

Cu<sl + 2 AgN0 <3 aql - Cu(N0 3) <2 aql + 2Ag<sl

4. Double Replacement reactions. These are reactions in which two compounds "switch partners". An example is the reaction between sulfuric acid and zinc carbonate:

H2S0 4<aql + ZnC0 3<ls

- H2C0 3<alq

+ ZnS0 4<aql

5. Combustion reactions. These are reactions in which a substance burns by combining with oxygen. An example is the burning of propane gas:

C3H8<lg + 50 2<gl -- 3 C0 <2 gl + 4 H20 0l

The reactions you will be doing in Part 1 of the experiment all produce a gas as one of the products of the reaction. If you can identify the gas that is produced, it is usually possible to figure out the entire reaction equation. Here is a list of some common gases and properties you can use to identify them.

Nitrogen dioxide, N0 2• This gas can be identified by its reddish brown color. It has an irritating odor.

Hydrogen sulfide, H2S. This gas is colorless, but it can be identified by its distinctive smell, the smell of rotten eggs.

Sulfur dioxide, S0 2• This gas is colorless. It has a sharp, choking odor and irritates your throat. Itis one of the smells you have smelled from a burning match or a fireworks display.

Chlorine gas, Ch. This gas has a pale yellow color. It is irritating to your eyes and throat and smells like a swimming pool.

Oxygen gas, 0 2• This gas is colorless and odorless. Oxygen accelerates combustion reactions. To test for oxygen, insert a glowing (not burning) wood splint into the container. If there is oxygen present, the wood splint should burst back into flame.

7-2

Hydrogen gas, H2• This gas is colorless and odorless. It bums rapidly with oxygen from the air. To test for hydrogen, bring a burning wood splint up to the mouth of the container. If there is H2 present, there will be a pop or a barking sound. This is due to a small explosion that happens when the hydrogen is ignited.

Carbon dioxide, CO2• This gas is colorless and odorless. The test for carbon dioxide is called the limewater test. A glass delivery tube is used to take the gas that is being tested and bubble it through limewater, which is the common name for a saturated Ca(OH)2 solution. If the gas is carbon dioxide,the following reaction will occur: C0 <2 gl +

Ca(OH)<a2

ql - CaC0<3 ls

+ H20 . You can

recognize a positive test for CO2 because the CaC0 3 precipitate will give the limewater solution a milky, cloudy, white appearance. There is a picture of a delivery tube being used below.

Figure 7-1 Delivery Tube

limewater solution

The delivery tube is used to transport gas produced by a reaction into a solution of limewater (saturated Ca(OH)2). A milky white cloud in the limewater is a positive test for carbon dioxide.

The reactions you do in Part 2 of the experiment are all double replacement reactions that lead to the formation of a precipitate. When you mix two solutions of soluble salts together, it is possible for them to switch partners and form a new salt that is not soluble in water. This insoluble product is called the precipitate. At first it will appear as a cloudy solid suspended in the solution, but it will eventually settle out to the bottom of the solution because it is more dense. A precipitation reaction happens when youmix lead(II) nitrate solution with potassium iodide solution: Pb(N03 )

a2( The solid Pbh is the precipitate in this reaction.ql + 2KI(aql - 2 KN0 3a(

q) + Pbl<2sl .

When you are trying to figure out what the insoluble product in a precipitation reaction is, it is helpfulto use a solubility table. There is one provided in the appendix section of this lab book.

7-3

Procedure

Part 1. The reactions in Part 1 all produce a gas as one of the products of the reaction. Use the tests described in the introduction to determine which gas is produced. When you are trying to decide what gas might be produced by a reaction, use your common sense to minimize the number of tests you have to do. For example, if the experiment calls for you to decompose KC1O3, the only possible gases thatcan be produced are 02 and Cb. It would make no sense to test for H2 or CO2!

Unless otherwise told by your instructor, do the following experiments with a partner to conserve chemicals.

A. Put about 10 mL of water in a large test tube. Add 2 drops of phenolphthalein solution. (Phenolphthalein is an indicator. It will change color to pink if a basic substance, such as a metal hydroxide, is present.) Bring a .QIT watch glass up to your instructor to obtain a piece of sodium metal. Do not touch the sodium metal with your fingers, use your tongs or a stirring rod to move it around. Drop the piece of sodium into the water in the test tube. Test to see what gas is being produced by the reaction. Remember, any gas produced is going to quickly escape- so think in advance about what gases might be produced by the reaction, and have the materials ready to test for those gases

B. Put 1 mL of 0.1 M Na2S into a small test tube. Under the hood, add a few drops of 6 M hydrochloric acid (HCl). Remove the test tube from the hood just long enough to smell the gas that is being produced. Dispose of the contents of the test tube in the sink under the hood, not in the classroom.

C. Put about one quarter of a spoonful of solid potassium chlorate (KC1O3) in a large test tube. Add a tiny amount of manganese(IV) oxide (MnO2) to act as a catalyst. Clamp the test tube to a ring stand at an angle. Heat the solid in the test tube with your Bunsen burner until it melts and bubbles. Test to see what gas is being produced by the reaction.

D. Put about 2 mL of 6 M hydrochloric acid (HCl) solution into a large test tube. Add one piece of mossy zmc. After it has bubbled vigorously for several seconds, test to see what gas is produced by the reaction.

E. Put 1 marble chip (CaCO3) into a large test tube. Add 2 mL of 6 M hydrochloric acid (HCl). Test to see what gas is produced by the reaction. When you clean up, be sure not to leave any marble chips in the sink.

7-4

Part 2. In this part of the experiment you will mix two salt solutions together to see if a precipitate forms. For each trial, put 1 mL of the first solution into a small test tube. Then add 1 mL of the second solution and mix thoroughly. If there is a precipitate (which may just look like a cloudy substance in the solution), describe its color on your report sheet. To figure out what the chemical formula of the precipitate is, use a solubility table. If there is no precipitate when you mix the two solutions together, then there is no chemical reaction, and you should write "no reaction" where the lab report asks for the chemical equation.

B. 0.1 MFeCb + 6 M NH4OH

C. 0.1 M NaNO3 + 0.1 M KI

E. 0.1 M Zn(NO3)2 + 0.1 M NaOH

7-5

CHEMISTRY 22 Name -------------Experiment 7, Chemical Reactions and Equations Report Sheet

Part 1.A. What gas is produced in the reaction? _

What color is the solution in the test tube after the reaction? ----

Write the equation for the chemical reaction that happened in the test tube:(hint: the color of the phenolphthalein should help you figure out the other product- go back and look at the instructions).

What type of reaction is this? (combination, decomposition, single replacement, or double replacement) -----------------

B. What gas is produced in the reaction? _

Write the equation for the chemical reaction:

What type ofreaction is this?

C. What gas is produced in the reaction? _

Write the equation for the chemical reaction: (The other product is potassium chloride. MnO2 is a catalyst, so it should not be included in the reaction equation.)

What type of reaction is this?

D. What gas is produced in the reaction? _

Write the equation for the chemical reaction:

What type of reaction is this? _

7-6

E. What gas is produced in the reaction? _

Two chemical reactions happened in this experiment. First, write the equation for the double replacement reaction that occurred when you added HCl to the CaC0 3:

One of the products of the double replacement reaction then decomposes to give the gas. Write the equation for this decomposition reaction:

Part 2.A. Is there a precipitate(PPT)? _ Color of PPT ---------

Formula and Name of the PPT: ------

Equation for the reaction: ----------------------

B. Is there a precipitate(PPT)? _ Color of PPT ---------Formula and Name of the PPT: ------

Equation for the reaction: ----------------------

C. Is there a precipitate(PPT)? _ Color of PPT ---------Formula and Name of the PPT: ------

Equation for the reaction: _

D. Is there a precipitate(PPT)? _ Color of PPT ---------Formula and Name of the PPT: ------

Equation for the reaction: _

E. Is there a precipitate(PPT)? _ Color of PPT ---------Formula and Name of the PPT: ------

7-7

Equation for the reaction: ----------------------

7-8

Additional Exercises.

1. Balance the following equations.

a) Sn(C0 3 ) 2 Sn0 2 + CO2

b) P20s + H20 - H 3P0 4

c) C4H10 + 02 - CO 2 + H20

d) FeS2 + 02 - Fe20 3 + so2

2. Write balanced chemical equations for each of the following reactions. You may need to figure out what one (or more) of the products are based on the description. Include the states (s, 1, g, or aq).

a) A piece of manganese metal is placed into a solution of hydrobromic acid. Bubbles of gas are formed, and the other product is a salt of manganese(II).

b) Potassium chromate solution is mixed with lead(II) nitrate solution. A yellow precipitate forms.

c) Water is added to some solid sodium peroxide. A gas is produced that causes a glowing wood splint to burst into flame. The other product is sodium hydroxide.

d) A piece of silver metal is placed in warm nitric acid solution. The products of the reaction are water, silver nitrate, and a reddish-brown gas.

e) Some solid aluminum carbonate is heated in a test tube. A gas is produced in the reaction and the other product is a white powder.

CHEMISTRY 22Experiment 8, Empirical Formula

IntroductionEmpirical literally means experimentally determined. The empirical formula of a

compound refers to the formula that has been experimentally determined. However, a more practical way of defining empirical formula is to say that it refers to the smallest whole number ratio of moles of elements within a compound. Therefore, in order to determine the empirical formula of a compound, one needs to know the moles of the elements within that compound, figure out their ratios, and express that ratio in smallest whole numbers.

Ionic compounds such as magnesium fluoride do not form molecules, they consists of ions. The empirical formula for ionic compounds represents the smallest whole number ratio between the ions. Ionic compounds are only expressed using empirical formula. On the other hand, covalent compounds such as glucose (a sugar present in blood and used by cells as fuel) can be represented by either a molecular formula or an empirical formula. The molecularformula of glucose is C6H12O 6, which tells you the number of atoms present in one molecule of glucose. The empirical formula of glucose is CH2O because that is the smallest whole number ratio of moles of elements within glucose.

The purpose of this experiment is to determine the empirical formula of a sulfide of copper. A sample of copper is heated in the presence of sulfur in a crucible inside the hood. After the sample cools to room temperature its mass is determined. The mole ratio and empirical formula will be calculated.

Procedure

Heat a crucible until the bottom turns red using a Bunsen burner inside the hood, and wait for it to cool to room temperature and measure its mass to ±0.0lg.

While waiting for your crucible to cool off you can remove the plastic coating on the copper wire. Take a piece of copper wire about 1 foot long (roughly 2 grams). Using your tongs to hold it, heat the copper wire directly in the bunsen burner at the hottest part of the flame.Keep the wire moving (so it does not melt) and be sure to heat the entire length of the wire. This process should bum off the plastic coating. Now sand your wire with a small piece of sandpaper.

Shape your wire into a small coil by winding it around a pencil so that it fits into a crucible. Measure the mass of the crucible (without the lid) plus wire to ±0.01 g. Add an amount of sulfur that weighs about half as much as the copper to the same crucible (about 1 g). This will make the sulfur the excess reactant. It is not necessary to measure the mass of sulfur precisely because any sulfur that does not chemically combine with copper will burn, forming sulfur dioxide gas which escapes. That is why one should heat the copper and sulfur mixture inside the hood.

Cover the crucible and place it on a clay triangle that is supported on a ring stand inside a hood. Heat slowly at first and then gradually increase the rate of heating until the hottest part of the flame is just underneath the crucible. After the sulfur has stopped burning (when you pick up the lid using a crucible tong, there is no blue flame inside the crucible), keep the crucible red hot for about 5 more minutes. Then pick up the burner and heat all around the crucible using the hottest part of the flame so that it is also red hot (so you burn all excess sulfur off as SO2). When no more sulfur is present, let the crucible cool to room temperature. When cool, remove the lid and measure the mass of the crucible and its contents to ±0.0lg. Reheat the crucible, cool and

8-1

8-2

weigh. If all of excess sulfur is gone, the mass should not change. Record this mass. Remove the piece of copper sulfide and break it. Note that its properties are different from both copper and sulfur.

8-3

CHEMISTRY 22Experiment 8 Name _ Empirical Formula of a Sulfide of Copper Report Sheet

Mass of empty crucible g

Mass of crucible and copper wire g

Mass of copperShow work for calculating mass of copper:

Mass of crucible and copper sulfide afterheating to a constant mass g

Mass of sulfur in copper sulfide Show work for calculating mass of sulfur in copper sulfide:

Moles of copperShow work for calculating moles of copper:

-------

moles

Moles of SulfurShow work for calculating moles of sulfur:

Empirical Formula:Show work for calculating empirical formula:

--------

moles

Write a balanced chemical equation for the reaction that took place between copper and sulfur.

8-4

Problems:

1) In a similar experiment, a student used 1.50 g of manganese to react with excess fluorine to form manganese fluoride. If 4.09 g of manganese fluoride is formed, what is the empirical formula of that compound?

1.--------

2) The decomposition of 5.00 g of vanadium sulfide produced 1.94 g of vanadium. What is the empirical formula of that compound?

2.--------

3) A compound is found by analysis to contain 81.82% carbon and the rest hydrogen. What is the empirical formula?

3.--------

4) If the molar mass of the compound in problem 3 is 132 g/mol, what is the molecular formula of the compound?

4.--------

5) Calculate the simplest formula of a compound containing 4.578 g of potassium, 2.924 g of arsenic, and 2.498 g of oxygen.

5.--------

9-1

CHEMISTRY 22Experiment 9, Empirical Formula of a Hydrated Salt

Introduction

The crystals of many ionic compounds (salts) contain water molecules in their crystal structures.These salts are called hydrated salts. An example of a hydrated salt is sodium acetate trihydrate, which has the formula NaC2H3O2 • 3H2O. The formula shows that for every one mole of sodium acetate in the crystal, there are three moles of water. When writing the formula of a hydrated salt, a dot is used to separate the water of hydration from the formula of the salt.

It is often possible to remove the water of hydration by heating the hydrated salt. For example, when copper(II) sulfate pentahydrate is heated, it loses all 5 molecules of water, which evaporate away: CuSO4• 5 H2O<sl CuSO4<s l + 5 H2O(gl. The salt that is left behind after removing the water of hydration is called the anhydrous salt.

In this experiment you will be given a sample of a hydrated salt. You will heat the salt to remove the water of hydration. By measuring the mass of water that is lost during heating, you will be able to calculate the percent water in your salt and determine the empirical formula of the hydrated salt.

Procedure

Obtain a sample of a hydrated salt from your instructor. It probably doesn't look wet, but there 1.§. water in the crystal of the salt you will be given. Weigh a clean, dry evaporating dish to the nearest 0.01 gram and record the mass on your data sheet. Put your entire hydrated salt sample into the evaporating dish and weigh it again. (Don't worry iflittle bits of your salt stay stuck in the sample tube.)

Place the evaporating dish on a wire gauze on a ring stand. Heat it with your Bunsen burner, starting with a low flame. If you heat it too rapidly at the beginning, the salt may pop or sputter and you will lose sample. Gradually increase the temperature of the flame and heat the salt for several minutes.However, if you see the salt turning black or brown around the edges, you are heating it too strongly and you should decrease the flame. Turn off the burner and let the evaporating dish cool for 15 minutes.Weigh the dish. The only way to know if the salt has been completely dried is to heat it to constant mass. Heat the dish again with a strong flame for 5 minutes. Cool it off completely and weigh it again. Repeat as necessary until the mass reaches a constant value.

Calculate the percent of water in your hydrated salt. Take your lab book up to your instructor to get the true percent water and the formula of the anhydrous salt. Calculate the empirical formula of the hydrated salt. That is, calculate how many moles of water there are per mole of anhydrous salt. There is a sample calculation shown on the next page.

(

l

3

l

9-2

Sample Calculation

A student does this experiment and gets the following data: Mass of evaporating dish 29.36 g Mass of evaporating dish + sample 34. 88 g Mass of evaporating dish + salt(after heating to constant mass) 31.76 g

The mass of the hydrated salt sample is 34.88 g- 29.36 g = 5.52 g The mass of water lost during heating is 34.88 g- 31.76 g = 3.12 g The mass of the anhydrous salt is 5.52 g-3.12 g = 2.40 g

3 12So, the experimental percent water in the hydrated salt is: ( · g ) (100% )= 56.5%water

5.52g

The student now finds out from the instructor that the true percent water in the salt is 56.9% and the formula of the anhydrous salt is Na3PO4• The next step is to determine what the empirical formula of the hydrated salt is. That is, she needs to figure out what Xis in the formula: Na3PO4 · XH 2O.

3. 1 2g H20 ) ( 1 mol e H2 0 = 0.173moles H2O18.02 g H2O

2. 40g N a P 3 0 4)( lmoleNa3P0 4 =0.0146molesNa3PO4( 163.94 g Na3PO4

0.173moles H 2O=

l l .8moles H2O So the formula of the salt is Na3P O4 ·12 H2O0.0146moles Na3PO4 l .0Omoles Na3PO4

If your experimental percent water value has more than a 5% error when compared to the true value, then you should calculate the empirical formula using the given percent water, rather than your experimental mass data. Assume that you have a 100 g sample for this calculation. For example, the student above is told that her true percent water is 56.9%. This means that a 100 g sample of her salt would contain 56.9 g of water and 43.1 g (100.- 56.9g) of anhydrous salt. Using these numbers, the calculation would look like this:

56. 9g H20 ) ( 1 mol e H 2 l 0 = 3.16moles H2O18.02 g H2O

43. l g N a P 3 0 4 )( l m o l e N a 3 P 0 4

( 163.94 g Na3PO

4

= 0 .263mo les Na P O 4

3.16moles H2O=-12-.0-m-o-les -H-2O--

So the formula of the salt is Na PO4 ·1 2H O

(

9-3

0.263moles Na PO 1.00 moles Na PO3 2

3 4 3 4

I

9-4

CHEMISTRY 22Experiment 9, Hydrated Salt Report Sheet

Name ------------

Sample Number _

ExpenmentaID ataMass of evaporating dish g

Mass of evaporating dish + sample g

Mass of evaporating dish + sample after first heating g

Mass of evaporating dish + sample after heating to constant mass g

Calculated Values (show work)Mass of hydrated salt sample

gMass of water in the sample

gPercent water in the hydrated salt

%

Values from Your Instructor Correct % water Formula of Anhydrous Salt

Calculated Values (show work)% error for percent water in salt

%

Moles of water in the hydrated salt

mol

Moles of anhydrous salt

mol

Empirical formula of hydrated salt

9-4

Additional Problems

1. What is the percent water in PtC14 • 4 H 2O ?

2. A 8.00 g sample of hydrated CoC12 is heated and loses 3.63 g of water. How many moles of water are combined with each mole of CoC12 in the hydrated salt?

3. A hydrated salt sample is analyzed and found to contain 4.07 g of potassium, 1.61 g of phosphorus,2.91 g of oxygen, and 1.41 g of water. What is the empirical formula of this hydrated salt?

10-1

CHEMISTRY 22Experiment 10, Stoichiometry

Name -------------

Stoichiometry is the study of the quantitative relationships between the amounts of reactants and products formed during a chemical reaction. A balanced equation gives the quantitative proportions of reactants and products in a chemical reaction. From the masses of reactants (or products) you can calculate the amount of products (or reactants) that would be formed in the chemical reaction by using the following solution map.

A ► B

Moles of Moles of Mass of

i ► . nceA i ►.nceB i►.nceB

molar mass coefficients from balanced equation molar mass

When you have two or more reactants, the amount of product that is formed in the reaction is based on the limiting reactant - the reactant that limits the amount of product that is formed. You can find this out by calculating the amount of product formed from the quantities of each of the reactants. The limiting reactant is the reactant that makes the least amount of product. The amount of product that can be made in the reaction is called the theoretical yield and is based on the limiting reactant. The other reactant(s) is called the excess reactant unless it also gives exactly the same amount of product as the limiting reactant. However, the actual amount of the product formed in the chemical reaction (experimental yield) is always less than the theoretically calculated amount of product due various reasons.

The percent yield of the reaction is:

% yield = experimental yield x I 00theoretical yield

In this experiment you will mix a solution containing a weighed amount of calcium chloride with a solution containing a weighed amount of sodium carbonate. One of the products of this double displacement reaction is calcium carbonate, which precipitates out. This is separated from the other soluble product, sodium chloride, and excess reactant by filtration. You will calculate which is the limiting reactant, the theoretical yield and the percent yield of the reaction.

Mass of Substance A

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Procedure

Use only distilled water for the experiment. Keep all reagent bottles tightly closed when reagents are not being used.

Take two clean and dry beakers (100 mL,150 mL or 250 mL). Number the beakers 1 and2. Place beaker 1 on the balance pan and press "TARE" or "RE-ZERO". The mass of the beaker should show as 0.0 g. Add calcium chloride to the beaker till the mass reading is close to 2 g. Note down the precise reading for the mass of calcium chloride in the report sheet. Similarly, place beaker 2 on the balance pan. Press "TARE" or "RE-ZERO" and then add an amount of sodium carbonate that is equal to the mass of calcium chloride. Try to stay within 0.05 g of the mass of calcium chloride and record the mass. Add 10 mL of water to beaker 1 and stir till all the solid is dissolved. Add 20 mL of water to beaker 2 and stir until all the solid is dissolved. Add the solution in beaker 2 (sodium carbonate solution) to beaker 1 (calcium chloride solution). Stir the mixture for about 5 minutes. A white precipitate should form soon after the addition.

Take a filter paper and write your name on it with pencil. Place the filter paper on a watch glass and record the mass of the filter paper and the watch glass. Fill your wash bottle with distilled water. Fold the filter paper, fit it into a funnel and moisten it with distilled water from a wash bottle. Place the funnel with filter paper on the ring stand with clay triangle and set it up for filtration. Stir the contents of the beaker 1 and transfer it to the filter paper with the aid of a rubber policeman. Do not fill the filter paper more than two-thirds full or else the product may seep through the space between the filter paper and funnel and there will be a loss of product.Scrape the product from the walls of the beaker with the rubber policeman and rinse the beaker with distilled water from the wash bottle. Transfer all the precipitate into the filter paper. You may rinse the beaker with several small portions of distilled water and transfer it to the filter paper. This will dissolve all the soluble compounds from the precipitate. Always let it drain between the rinses and make sure the funnel is never more than two thirds full. A maximum of 100 mL water may be used.

After all the liquid drains down, carefully open the filter paper with the precipitate and place it on the watch glass. Keep it in the oven (temperature should be around 110 °C) for 30 minutes for drying. Take it out of the oven allow the watch glass with the filter paper to cool to room temperature. Record the mass of the watch glass and filter paper and calculate the amount of calcium carbonate formed experimentally. From the masses of calcium chloride and sodium carbonate you used for the reaction find out which is the limiting reactant and calculate the theoretical yield of calcium carbonate that would be formed in the reaction. Calculate the percent yield of the reaction using the formula given above.

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CHEMISTRY 22Experiment 10, Stoichiometry Report Sheet

Name -------------

Balanced equation for the reaction of calcium chloride with sodium carbonate:

Mass ofCaClz g

Mass of Na2C03 g

Mass of watch glass + empty filter paper

(Write your name on filter paper)

g

Mass of watch glass+ filter paper+ CaC0 3 (after drying) g

Mass of CaC0 3 (This is experimental yield) g

Theoretical yield of CaC0 3 (show calculations below) g

Calculations for theoretical yield of CaC03:

Which is the limiting reactant?

Which is the excess reactant?

Amount of excess reactant left over after the reaction (show calculations below)

g

Calculations for amount of excess reactant left over:

Percent yield of reaction (show calculations below)

%

Calculations for percent yield:

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Experiment 10, Stoichiometry Additional Problems

Show all calculations using dimensional analysis. Your answer should have the correct number of significant figures.

1. What is the mass of mercury that can be prepared from 1.25 g of cobalt metal? The reaction is:CorsJ + HgCh (aqJ - CoCh (aqJ + Hg (IJ (not balanced)

2. The US space shuttle employs a mixture of aluminum and ammonium perchlorate for fuel. The reaction is:

Al (sJ AlzO3 ( s) + AlCh (s) + NO (g) + H2O (g)

(a) Balance the equation.

(b) What mass of ammonium perchlorate (in kg) should be used in the fuel mixture with5.00 kilograms of aluminum?

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3. (a) Write the balanced equation for the combustion of octane(CsH1s).

(b) Assuming gasoline is 100% octane, with a density of 0.692 g/mL, what mass of carbon dioxide will be produced by the combustion of 15.0 gallons of gasoline?

4. 50.0 g of tin(II) fluoride reacts with 50.0 g of sodium phosphate

(a) Write the balanced equation for the double displacement reaction.

(b) Calculate the grams of tin(II) phosphate that would be formed in the reaction.

(c) How many grams of tin(II) fluoride and sodium phosphate remain after the reaction is complete.

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5. In a lab experiment, the reaction of 1.55 g oflead(II) nitrate with excess potassium iodide gives 2.00 g oflead (II) iodide. What is the percent yield of this reaction?

6. Urea (CH4N2O) is a common fertilizer that can be synthesized by the reaction of ammonia with carbon dioxide according to the following equation:

NH3 + + H20 (equation not balanced)

In an industrial synthesis, 43.0 kg of urea was obtained when 34.0 kg of ammonia was reacted with 50.0 kg of carbon dioxide. Determine the limiting reactant and calculate the percent yield of the reaction.

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CHEMISTRY 22Experiment 11, Preparation and Properties of Hydrogen Gas

IntroductionHydrogen is the most abundant element in the universe, making up more than 75% of the mass of the

universe. Hydrogen is also very abundant on Earth, but most of it is found in water and other compounds. Hydrogen gas, H2, is actually quite rare on Earth. Only one out of every one million gas molecules in our atmosphere is Hz. Hydrogen gas is colorless and odorless. It has the lowest density of all gases, about one fifteenth of the density of air. At standard pressure, hydrogen must be cooled down to minus 253°C to turn it to liquid.

Hydrogen gas can be prepared in many ways. It can be obtained directly by the electrolysis of water:2H20 electricity )0 g2( l + 2

H2g(

) ' Hydrogen is produced in the reaction of very active metals (from

periodic table group IA and IIA) with water: 2Li<sl + 2H02 H2g( l + 2LiOH(aq)· It is also producedwhen metals react with acids: Mn<sl + 2HCl<aql H<2g l + MnCl2( aql Industrially, hydrogen is often produced from methane gas at high temperatures: CH4(gl + H20< gl CO(gl + 3 HC2 gl

One of the chemical properties of hydrogen gas is that it burns rapidly with oxygen to make water, releasing a large amount of energy. This property is used in the standard laboratory test for hydrogen gas. If you bring a burning wood splint near the mouth of a container where hydrogen gas is being produced, you will hear a pop or a barking sound as a small explosion happens between the hydrogen gas and oxygen from the air. This small explosion is not dangerous, but it is dangerous to ignite a large amount of hydrogen. For this reason, you should be careful not to bring any flames near your zinc/sulfuric acid hydrogen generator.

In this experiment you will prepare H2 gas, and investigate some of its chemical and physical properties.

ProcedurePart 1.A. Bring a 4ry_ watch glass up to your instructor to obtain a piece of sodium metal. Do not touch the

sodium metal with your fingers, use your tongs or a stirring rod to move it around. Drop the piece of sodium into a small test tube that contains about 3 mL of water. While the reaction is taking place, test for hydrogen by bringing a burning wood splint near the mouth of the test tube.

B. Place a piece of calcium metal into a large test tube that contains about 5 mL of water. If the evolution of gas is slow, loosely stopper the tube and let the reaction run for a couple of minutes. Test for hydrogen using a burning wood splint.

C. Put 3 mL portions of 6M HCl into each of three large test tubes. Put a piece of magnesium metal into one test tube, a piece of aluminum metal into the next, and a piece of copper metal into the last. Watch carefully to see if any reaction happens. When vigorous bubbling occurs, test for hydrogen gas. Note: It may take five minutes for the reaction to begin because the pieces of metal may be covered in an oxide coating. If any of the metal pieces do not react, please rinse them off and put them in the appropriate recycling container. Do not put metal pieces in the sink!

Part 2.A. Assemble a hydrogen generator according to the following directions (see picture on the next

page). From the cabinets under the hoods obtain a gas-collecting trough (pneumatic trough) and two

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wide mouth bottles. Check out a funnel tube from the stockroom, and get two glass squares from the

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reagent shelf. Put 5-8 pieces of mossy zinc into one of the wide mouth bottles. Stopper it with the funnel tube (also called a thistle tube). Clamp this bottle to a ring stand. Hook up the rubber hose of the funnel tube to the gas inlet on the pneumatic trough. Using a beaker, pour tap water into the pneumatic trough until the water is to a depth of about one inch.

Fill a wide mouth bottle completely with water. Cover it with a glass square and turn it upside down.Slide the inverted bottle into the water in the pneumatic trough, making sure no air bubbles get inside. You can leave the bottle standing in the trough until you are ready to collect hydrogen gas. This is how you will prepare all the bottles and test tubes in which you collect hydrogen gas.

To begin making hydrogen gas, pour 30 mL of 6M hydrochloric acid into the generator bottle through the funnel tube. You should see a reaction begin and gas should be bubbling out through the hole in the bottom of the pneumatic trough. If the reaction is too slow, you can speed it up by adding a few drops of 0.1M cupric sulfate solution, a catalyst for this reaction.

Collect one bottle of gas by sliding the inverted bottle over the hole in the pneumatic trough and allowing the hydrogen gas to displace the water. Discard this first bottle. Collect one large test tube of gas. Bring a burning wood splint near the mouth of this tube. If it “pops” too loudly, that means there is still too much oxygen in the system and you need to let the generator run for a couple more minutes.Collect another test tube of hydrogen and test with the burning wood splint. If it gives a quiet pop, then you are ready to begin collecting hydrogen for the rest of the experiment. Collect three large test tubes of hydrogen and one wide mouth bottle. Store the hydrogen gas by standing the containers upside down on a glass square on your bench. When you clean up your hydrogen generator, do not leave pieces of zinc in the sink. Rinse off the zinc and put the pieces in the “used zinc” container.

B. Let one of the test tubes of hydrogen stand un-stoppered in your rack, mouth upward, for two minutes. Then test it for hydrogen using a burning wood splint.

C. In one hand hold a test tube of hydrogen, mouth downward. In your other hand hold a empty test tube (which is actually full of air) mouth upward. Bring the mouths of the two test tubes together. Keeping the test tubes together, invert the tubes so that the hydrogen-containing tube is directly below the tube with air in it. After 5 seconds, separate the tubes and test each one for hydrogen.

D. Light a wood splint. Take the bottle of hydrogen you collected, hold it with the mouth downward, and immediately insert the burning wood splint up into the bottle. Does the splint burn inside the bottle? Does the splint re-ignite as you withdraw it from the bottle?

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CHEMISTRY 22 Name Experiment 11, Hydrogen Gas Report Sheet

Part 1.A. What did you observe when you put the sodium metal into water?

Write the balanced equation for the reaction between sodium and water:

B. What did you observe when you put the calcium metal into water?

Write the balanced equation for the reaction between calcium and water:

C. Which of the three metals did react with HCl?

Write balanced equations for the reactions. If there was no reaction, write “no reaction” instead.

Mg HCl

Al HCl

Cu HCl

Part 2.A. Write the balanced chemical equation for the reaction between zinc metal and hydrochloric acid

in your hydrogen generator:

Write the balanced chemical equation for the reaction that makes the “pop” sound (the burning of hydrogen with oxygen):

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-

Why was the first bottle of hydrogen you collected discarded?

B. Did this test tube give a positive test for hydrogen gas?

What happened to the hydrogen gas originally present in the tube?

C. What were the results of your experiment?

Explain the results and what it tells you about the properties of hydrogen gas.

D. Does the wood splint bum inside the bottle of hydrogen?

Does the wood splint re-ignite as you withdraw it from the bottle of hydrogen?

Explain these results:

Additional Questions.Write balanced chemical equations for each of the following reactions. If there would be no reaction,

write "no reaction" instead. (Hint: each of these reactions is similar to one you did in the experiment).

L sl + H20

Mgcsl + HzS04(aq) - Cues) + HBfcaql - Zn<sl + HClcaql - A s> + HzS04(aq) -

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Bacs> + H 20 -

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CHEMISTRY 22Experiment 12, Preparation and Properties of Oxygen Gas

Introduction

Oxygen is the most abundant element of the Earth's crust. It is widespread in familiar compounds, including not only water, but also many rocks and minerals and most biological molecules. As an element, oxygen is usually found in the form of a diatomic gas, 0 2. Diatomic oxygen makes up about 21% of the Earth's atmosphere by volume. 0 2 is colorless and odorless. It must be cooled down to minus 183°C (at standard pressure) to turn it into a liquid.

Oxygen gas can be prepared in many ways. Plants, of course, produce oxygen continually as a product of photosynthesis. Industrially, oxygen gas is usually produced by liquefying air, then purifying the oxygen through distillation. In the lab, oxygen can be produced by the decomposition of oxygen containing compounds. It can be obtained by the electrolysis of water: 2 H2O electncity ) O<2gl + 2H<2gJ •

Potassium chlorate can be decomposed thermally: 2KC1O<3sl 2KCl<sJ + 3O<2 Jg · In this experiment you will prepare 0 2 gas by the decomposition of the unstable compound hydrogen peroxide:2 H2O20l - O <2 gl + 2H2Oc iJ · The decomposition of hydrogen peroxide is spontaneous (that is, youdon't have to add energy to get it to happen), but it is slow at room temperature. You can speed up the reaction by adding a catalyst, a substance that increases the rate of a chemical reaction without itself being consumed in the reaction. The decomposition of H2O2 is catalyzed by several things, including ultraviolet light, the enzyme catalase (which is found in blood), and manganese(IV) oxide.

In this experiment you will prepare oxygen gas by the decomposition of hydrogen peroxide using MnO2 as a catalyst. You will collect the gas by water displacement using a pneumatic trough, and then investigate some of the reactions and properties of oxygen gas.

Procedure

Preparation of Oxygen.From the cabinets under the hoods obtain a gas-collecting trough (pneumatic trough) and five wide

mouth bottles. Check out a funnel tube from the stockroom, and get four glass squares from the reagent shelf. Put a very small amount (just a pinch) of manganese(IV) oxide powder into one of the bottles and stopper it with the funnel tube (also called a thistle tube). Make sure the end of the thistle tube is close to the bottom of the glass bottle. Clamp the bottle to a ring stand. Hook up the rubber hose of the funnel tube to the gas inlet on the pneumatic trough. Using a beaker, pour tap water into the pneumatic trough until the water is to a depth of about one inch. There is a picture of this assembly shown in experiment #11.

Fill a wide mouth bottle to the very top with water. Cover it with a glass square and tum it upside down. Slide the inverted bottle into the water in the pneumatic trough, trying not to let air get inside. You can leave the bottle standing in the trough until you are ready to collect oxygen gas. Prepare a total of four bottles this way.

In a graduated cylinder, obtain 50 mL of 9% hydrogen peroxide solution. To begin generating oxygen gas, pour about 10 mL of the hydrogen peroxide solution into the generator bottle through the funnel tube. You should see a reaction begin and gas should be bubbling out through the hole in the bottom of the pneumatic trough. If there is no gas bubbling out, check all your connections to make

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sure

* You can use very small quantity of yeast instead of manganese(IV) oxide

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they are tight. Whenever the rate of oxygen gas evolution slows down, add more of the hydrogen peroxide solution through the funnel tube.

Collect one bottle of gas by sliding the inverted bottle over the hole in the pneumatic trough and allowing the oxygen gas to displace the water. Remember which bottle is the first bottle collected. Collect three more bottles of oxygen gas. To remove a bottle of gas from the pneumatic trough, cover the mouth of the bottle with a glass square before removing it from the water. Store it upright on your bench, with the glass square over the bottle opening.

Properties of OxygenA. Ignite a wood splint in your bunsen burner and let it bum for about 5 seconds. Blow out the

flame and immediately insert the glowing splint into the first bottle of oxygen collected. What do you observe?

B. Put a small piece of charcoal (mostly carbon) in a deflagrating spoon. Hold the spoon in a burner flame until the charcoal glows orange. Quickly lower the spoon into a bottle of oxygen, being careful not to touch the sides of the bottle with the hot spoon or the bottle may crack. What do you observe? Hold the spoon suspended in the bottle while the charcoal burns. Once it has finished burning, remove the deflagrating spoon and pour about 10 mL oflimewater (saturated Ca(OH)2 solution) into the bottle. Cover with a glass square and shake the bottle. What happens to the limewater? If you have forgotten about the limewater test and what gas it is used to identify, look back to experiment 7.

C. Fill the deflagrating spoon about one eighth full with sulfur powder. Take the spoon, one of your oxygen bottles and a bunsen burner to the fume hood. Under the hood, ignite the sulfur using a burner flame and then lower the spoon into the bottle of oxygen. What do you observe? When the sulfur is done burning, pull the bottle out of the hood just long enough so that you can cautiously smell the odor of the sulfur dioxide that is produced in the reaction. To clean up your deflagrating spoon, heat it in a burner under the hood until all the sulfur bums off.

D. Pour 25 mL of water into a bottle of oxygen and replace the cover. Set this bottle near your bunsen burner. Obtain a loose clump of steel wool, about 4 cm in length. Hold the steel wool in your tongs and heat it in the hottest part of the burner flame until you see some of the steel glowing orange. Quickly lower the still-glowing steel wool into the oxygen bottle. What do you observe?

Notes: The water is placed in the bottle to keep it from cracking if the hot steel wool is accidentally dropped. Steel wool is mostly iron. When iron bums in an oxygen rich environment, it becomes fully oxidized. In other words, the iron should form an ion with the highest charge possible for iron.

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CHEMISTRY 22Experiment 12, Oxygen Gas Report Sheet

Preparation of Oxygen.What is the purpose of the Mn0 2 in the oxygen generator?

Name -------------

Write the equation for the decomposition of hydrogen peroxide. (A catalyst is not changed in the reaction, so it should not be included in the overall equation).

When you collected hydrogen gas in the previous experiment, you stored the bottles of gas with the mouth facing downward. When you collected oxygen gas, it was OK to store them with the mouth upward. Why are the two gases treated differently?

Properties of OxygenA. What did you observe when you put the glowing wood splint into the bottle of oxygen?

B. What did you observe when you put the glowing charcoal into the bottle of oxygen?

What did you observe when you put the limewater into the bottle?

What gas was present in the bottle (as demonstrated by the limewater test)?

Write the balanced equation for the reaction between the limewater (Ca(OH)2caq)) and that gas:

Write the balanced equation for the reaction that happened as the charcoal (carbon) was burning:

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C. What did you observe when you put the burning sulfur into the bottle of oxygen?

Write the balanced equation for the reaction that happened as sulfur was burning:

D. What did you observe when you put the glowing steel wool into the bottle of oxygen?

Write the balanced equation for the reaction that happened as the steel wool (iron) was burning:

Oxygen gas itself is considered a non-flammable gas. However, how would you describe the effect that oxygen has on the combustion reactions of other substances?

Additional Questions.Write a balanced chemical equation for the combustion of each of the following substances with

oxygen gas.

co

CsH1s(octane)

13-1

CHEMISTRY 22Experiment 13, Charles' Law Experiment

IntroductionA law is a summary of a set of observations that happen over and over again

under the same set of circumstances. J.A.C. Charles, a French scientist (1746-1823), observed that the volume of a gas is directly proportional to its absolute temperature under constant pressure and number of moles. This is called Charles's law and can be used to predict change in volume of a gas as a result of change in its temperature. This law can also be expressed mathematically as:

where V1 and V2 are initial and final volumes respectively, T1 and T2 are initial and final temperatures respectively. The temperatures must be expressed in Kelvin.

This behavior of gases can be explained using the kinetic molecular theory of gases. The kinetic molecular theory of gases assumes that a gas is made up of tiny particles called gas molecules. These molecules are in constant random motion colliding with each other and sides of the container. Pressure is a result of these collisions. When temperature is increased, the average speed of these molecules increases and therefore they collide with the sides of the container more frequently and with greater force. This increases the pressure, and in order for the pressure to stay the same the volume has to increase. Therefore, the volume increases when temperature is increased if pressure is to be kept constant.

In this experiment you will measure the decrease that takes place in volume of air as it is cooled from boiling point temperature of water to room temperature. Then you compare your measured volume to the calculated volume from Charles's Law.

ProcedureCheck out a 600-mL beaker, a thermometer, and a rubber stopper with glass tubing from the stockroom (A 400ml beaker also works ifno more 600 mL beakers are available).Use a 250mL flask (a 125 mL flask if you are using a 400 mL beaker) along with the rubber stopper with glass tubing, 600 mL beaker (or 400 mL beaker), a ring stand, wire gauze, and a utility clamp to set up the apparatus as shown in figure 13-1 on the next page. Do NOT wash the flask. The flask, stopper and the tube must be absolutely dry inside. Any trace of water in the flask will ruin the experiment.

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Figure 13-1Add water to the beaker up to the neck of the flask, and heat the flask by boiling the water inside the beaker. While the beaker is being heated fill a large pan (the pans are in cabinets under the hood) with sufficient cool tap water to cover the flask. Measure the temperature of the boiling water. When the air in the flask has been heated for fifteen minutes, wet a finger and tightly close the glass tube by placing your finger over it. Invert the flask in the pan of cool water, and after the end of the tube is under water remove your finger. Some water will go into the flask as the flask cools. Keep the flask under water for about five minutes while the temperature of the flask and cool water are equalizing. Then make sure that the pressure inside the flask is the same as the atmospheric pressure by making sure the level of the water inside and outside the flask are equal. You can do this by raising or lowering the flask until the water levels (inside and

outside of the flask) are equal. Then close the tube again with your finger. Take the flask out of the pan and measure the volume of the water that went into the flask. The volume of the water inside the flask represents the shrinkage in volume of the gas as it changed from the temperature of the boiling water to the temperature of cool water. Measure the volume of water with a graduated cylinder as accurately as possible. Then determine the total volume of the flask and glass tubing by filling the flask with water up to the brim, inserting the stopper and tube, then removing the stopper and tube and measuring the volume of the water in the flask using a large (250 mL or 500 mL) graduated cylinder.You can find the 250 mL or 500 mL graduated cylinders between the bench tops. This is the volume of the air V1 at boiling water temperature. The volume V1 minus the volume of water drawn into the flask is the volume of air at the lower temperature V2•

Calculate what the volume of air should be at the lower temperature using Charles's Law and compare the calculated and measured values.

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CHEMISTRY 22Experiment 13, Charles' Law Report Sheet

Name----------

Temperature of boiling water (T1)

Temperature of air after cooling (temp. of water bath) (T2)

OC= K

OC= K

Volume of water drawn into flask at the lower temperature mL

Volume of air at the temperature of boiling water (V1) mL

Volume of air at the lower temperature (measured)(V2) mL

Volume of air at the lower temperature (calculated) (V2) mL

Show work for calculating V2 (calculated):

Percent error in the measured value of

(V2) Show work for calculating %

error:

-------%

Problems:

1) The volume of a gas is 1.00 Lat 12 °c. If the pressure remains constant, to what temperature (in °C) must the gas be heated to triple the volume?

1.-------

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2) A certain quantity of a gas occupies 25 cubic meters at a pressure of 1.5 xl 04 Pa. What will be the volume of that gas if its pressure is increased to 3.00 atm under constant temperature?

2.------

3)gas has a volume of 45.0 mL at STP. Calculate the volume of that gas at 2.00 atm and 85 °c.

3.-------

4) A given mass of helium gas has a certain volume at 35 °c and 565 mmHg. Calculate the temperature (in °c) necessary to double the original volume of that gas when the newpressure is 720. torr.

4.-------

14 - 1

CHEMISTRY 22Experiment 14, Identifying Ions in a Solution

Introduction

Have you ever wondered what makes the different colors in fireworks? The presence of certain metals in the fireworks produces the brilliant color displays. When the firework is lit, energy is supplied to the sample. The atoms within the fireworks absorb the energy causing the electrons to become "excited," which means that the electrons jump from a lower to a higher energy level within each atom.

The electrons cannot remain excited. They must "relax" or lose the excess energy. To lose the energy, the electrons then jump from a higher to a lower energy level releasing light in the process. The wavelength of light released depends on the spacing between the higher and lower energy levels. Since elements have a specific spacing between energy levels, they release specific wavelengths of light.

Each specific wavelength of light released when a metal atom relaxes is related to a specific color.Thus when you excite atoms in a metal either by lighting fireworks or simply heating a solution containing metal ions, the sample will emit certain colors. The colors will depend on which metal is present in your solution. Thus you can identify which metal ion is present in your solution based on what color of light is emitted (released) when you heat a sample containing the ion.

In this lab, your ultimate goal is to identify the cation (positively charged ion) and anion (negatively charge ion) present in an unknown solution. Your unknown will only contain one cation and one anion. To identify the cation, you will first have to determine the different color flames that each metal ion emits when heated with a Bunsen burner. To identify the anion, you will add silver nitrate which may cause a distinctive colored solid to form. The color of the solid formed or lack of a solid being formed will allow you to identify the anion present in your unknown.

Procedure

Flame Test for Metal CationsAdd about 10 drops of each test solution into separate wells in the spot plate. Each test solution

contains one of the following cations: sodium, potassium, calcium, lithium, or barium ions. Mark or label wells accordingly. Clean the nichrome wire by dipping the loop in about 2 mL of 6 M HCl (CAUTION!!) and heating it in the hottest part of the flame until no color is imparted to the flame. Dip the clean wire into one of the test solutions and record the observed flame color that results. Some of the flame colors are very similar so it is also helpful to note the duration of the colored flame. Some flame colors last a long time while others disappear quickly, lasting only a few seconds. In the latter cases, observations must be recorded immediately. Repeat the flame test for the other solutions, cleaning the wire between each test. Even the slightest bit of sodium will impart a yellow-orange flame that can easily mask some of the other colors particularly if the color does not persist. Since this is particularly a problem when observing the potassium flame, you should view the flame for potassium through two pieces of cobalt glass.

Precipitation Test for AnionsAdd 1 mL of chloride ion, er, to a small test tube. Add 1 drop of 0.1 M AgNO3 to the test tube.

Note whether any solid forms. A solid that forms when two solutions are mixed together is called a precipitate. If a precipitate forms, indicate the color of the precipitate on your lab report. Repeat the same procedure for each of the anions. Clean up any spills of silver nitrate solution thoroughly- it will leave purple stains on the lab bench.

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Unknown AnalysisBring a small test tube to your instructor to obtain an unknown solution. The solution contains a salt,

with only one cation and one anion. Determine which ions are present in your unknown by using the flame and precipitation tests that you practiced in the previous parts of the experiment. Once you have figured out which ions are present, write the formula of your salt, writing the cation first and the anion second. Do not include charges when writing the formula of your salt.

14 - 3

CHEMISTRY 22Experiment 14, Identifying Ions in a Solution Report Sheet

Flame Test Results:

Name ------------

Cation Tested Observed Flame Color

Na+

K+

Ca2+

Lt

Ba2+

Precipitation Test Results:

Anion Tested Color of precipitate (if no precipitate forms, write "no ppt")

Chloride, erIodide, :r

Phosphate, Po/-

Sulfate, so/-

Unknown Results: Unknown Letter -----Observation from Flame or

Precipitation TestIon Present in Unknown

Cation - Flame Test

Anion - Precipitation Test

Formula of Unknown Salt (do not include charges) ------------------

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Questions:

1. Why do different elements have different colored flames when they are heated?

2. Why is it important to dip the wire in 6 M HCl before performing a flame test for each cation?

3. Why is it easy to get a false positive flame test for sodium?

4. Why is cobalt glass used when testing for potassium?

5. What two chemicals might be used to produce the red and green flame in a "Christmas log"?

6. How did you determine which ions were present in your unknown based on your flame and precipitation tests?

15-1

CHEMISTRY 22Experiment 15, Solutions

Introduction

Solutions are homogenous mixtures of two or more substances. Solutions can be prepared from all possible combinations of physical states. That is, solids can dissolve in liquids, liquids can dissolve in other liquids, gases can dissolve in solids, gases can dissolve in gases, etc. For example, air is a solution; it is a homogeneous mixture of nitrogen gas, oxygen gas, and several other gases. In this experiment, you will look primarily at solutions made with solid or liquid solutes dissolved in liquid solvents. You will investigate how the polarity of molecules is related to solubility, what factors affect the rate of dissolving, the difference between true solutions and suspensions, and the properties of supersaturated solutions.

Before you do this experiment you should read the chapter in your textbook about solutions. It will provide definitions for the many terms that are used to describe solutions and help you understand the factors that affect solubility.

Procedure

PART 1. Supersaturated Solution. Choose a large test tube from your locker that looks smooth and not scratched. Clean the test tube with soap and a test tube brush. Rinse the test tube several times with tap water and then twice with distilled water. Shake out the excess water, but do not dry with a paper towel.

Fill the test tube about half full with solid sodium acetate, then add about 2 mL of distilled water, rinsing down the sides of the test tube as you add it. Clamp the test tube to a ring stand and warm gently to get the crystals of sodium acetate to dissolve completely. Do not let the solution boil and avoid tipping or jostling the test tube. After the crystals have completely dissolved, let the test tube cool for a few minutes and then carefully set it in a beaker of cold water. While it is cooling in the water bath, go on to do the other parts of the experiment. Check your test tube every now and then - it should remain clear like water. If you see any crystals forming or any cloudiness while it is cooling down, reheat the test tube to get the crystals to dissolve.

After you finish parts 2-4 and your sodium acetate solution has cooled completely, go get one small crystal of solid sodium acetate from the reagent shelf. Remove the test tube from the water bath and drop in the single crystal of sodium acetate (the seed crystal). Observe what happens. Be sure to feel the outside of the test tube to note any temperature change.

PART 2. Polarity and Solubility. Place 0.5 mL of water in a small test tube and 0.5 mL of cyclohexane in another small test tube. Add 10 drops of oil (a nonpolar substance) to each test tube, stopper and shake. Observe to see if the oil dissolves.

Once again, place 0.5 mL of water in a small test tube and 0.5 mL of cyclohexane in another small test tube. This time add 10 drops of glycerol to each test tube, stopper and shake. Observe to see if the glycerol dissolves.

Cyclohexane or oil should not be put in the sink. Dispose of the mixtures containing cyclohexane or oil in the appropriate waste bottles in the hood.

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To help you decide if glycerol and cyclohexane are polar molecules, their structures are shownbelow. cH 2

OH OH OHI I I

CH/ 'CH 2f ICHi-CH-CH2 CH2..... /CH2

CH2

glycerol cyclohexane

PART 3. Factors that Affect the Rate of Dissolving.A. Obtain two crystals of solid sodium chloride of about the same size from the bottle labeled

large crystals. Grind one of the crystals with a mortar and pestle until it is a fine powder. Place this powder into a large test tube, and place the whole (unground) crystal into another large test tube. Add 15 mL of distilled water to each test tube. Stopper each test tube and shake to completely dissolve the sodium chloride. Note how long it takes for the solid to dissolve in each case. Save one of these solutions to use in Part 4.

B. Obtain two crystals of solid sodium chloride of about the same size from the bottle labeled large crystals. Place each crystal into a separate large test tube and then add 15 mL of distilled water to each test tube. Leave one test tube undisturbed in the rack, but stopper and shake the other one. Note how agitation affects the rate of dissolving.

C. Place 15 mL of cold tap water into one large test tube and 15 mL of hot tap water into another large test tube. Add ¼ spoonful of sugar (sucrose) into each test tube. Stopper and shake the test tubes to see how long it takes for the sugar to dissolve.

PART 4. The Differences between a Suspension and a Solution.A. Examine the sodium chloride solution saved from Part 3A. Hold the stoppered test tube

sideways over this page and see if you can read the print. Is the solution clear or cloudy? Make another mixture by combining ¼ spoonful of starch with 15 mL of distilled water in another large test tube. Stopper the test tube and shake vigorously. Can you read print through this starch suspension? Is the suspension clear or cloudy?

B. Set the test tubes containing the sodium chloride solution and the starch suspension in your rack for five minutes. Do the sodium chloride particles settle out of the solution? Do the starch particles settle out of the suspension?

C. Set up a piece of filter paper in a funnel like you did in Experiment 3. Pour the sodium chloride solution through the filter, collecting the filtrate in a small beaker. Do any of the sodium chloride particles get trapped on the filter paper or do they stay mixed in the water? Shake the starch suspension and then pour that through the filter paper. Do any of the starch particles get trapped on the filter paper?

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CHEMISTRY 22Experiment 15, Solutions Report Sheet

PART 1.

Name

Describe what you observe when the seed crystal is added to the supersaturated sodium acetate solution:

Is energy absorbed or given off (circle one) during this change?

Explain why the crystals form in the supersaturated solution when you add the seed crystal.

What would happen if you dropped a crystal of sodium acetate into an unsaturated solution of that salt?

PART2.For each of the four substances you used, circle if they are polar or nonpolar molecules. Then, in the boxes write down if the solute was soluble in the solvent or was not soluble.

Solutes

SolventsOil

polar I nonpolarGlycerol

polar I nonpolar

Waterpolar I nonpolar

Cyclohexanepolar I nonpolar

Based on your observations above, write a general rule about how the polarity of molecules is related to solublity:

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PART3.A. Which dissolved faster, the sodium chloride powder or the single large crystal (circle one)?

B. Which dissolved faster, the sodium chloride in the shaken or the undisturbed test tube (circle one)?

C. Which dissolved faster, the sugar in the hot water or the sugar in the cold water (circle one)?

In summary, write three ways you can speed up the rate of dissolving of a solid solute in a liquid solvent:

1.

2.

3.

For one of the three things you listed above, explain why it will increase the rate of dissolving.

PART4.Fill in the following table comparing the properties of the sodium chloride solution (a true solution) and the starch/water mixture (a suspension).

sodium chloride solution starch suspensionA. Is it clear or cloudy?

B. Do the solute particles settle out over time?C. Are the solute and solvent separated by filtration?

Based on what you observed above, how do you think the size of particles of sodium chloride in the solution compare to the size of the starch particles in the suspension? Explain.

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CHEMISTRY 22Experiment 16, Acid-Base Titrations

Introduction

Titration is an important technique in analytical chemistry. It is a volumetric technique in which two solutions are reacted with each other, and the volumes of the solutions needed to complete the reaction are carefully measured. The volumes are usually measured with an instrument called a buret.

There are many types of titrations. In this experiment, we will do an acid-base titration. A buret will be used to carefully dispense a known volume of an acid solution into a flask. A base solution will be added gradually from another buret until we have added just enough base to neutralize the acid completely. This point at which the acid and base have just neutralized each other is called the equivalence point or endpoint. We will use a pH indicator, phenolphthalein, to help us recognize when we have reached the endpoint.

There are two parts to this experiment. In part 1 you will make up a solution of sodium hydroxide that is approximately 0.2 M. In order to determine the precise molarity of your sodium hydroxide you will do a titration with a standard hydrochloric acid solution. It is called a standard solution because its precise molarity is known. After measuring how much of your base solution is needed to neutralize the standard acid, you will be able to calculate the precise molarity of your base solution, thereby standardizing your base solution.

In part 2 you will use your standardized sodium hydroxide solution to titrate an acetic acid sample of unknown molarity. After measuring how much of your base solution is needed to neutralize the acetic acid solution, you will be able to calculate the precise molarity of the unknown acetic acid solution.

Your instructor will demonstrate the technique of titration in class and also show you examples of titration calculations.

Procedure

PART 1. Preparation and Standardization of the Base SolutionPlace about 16 mL of 6M NaOH into your 500 mL bottle. Add distilled water to fill the bottle to the

shoulder. Obtain a rubber stopper that fits your bottle and then wrap it with parafilm. Stopper the bottle and mix the contents thoroughly by repeatedly inverting the bottle. Keep your bottle ofNaOH solution stoppered as much as possible.

Obtain two burets and a buret clamp from the stockroom. Wash each buret thoroughly with tap water. Make sure you rinse the sides of the buret and also rinse out the stopcock by letting liquid run through. Do one more rinse using distilled water. Let excess water drain from the burets by setting them in the clamp upside-down, with the stopcock open.

The final step in the preparation ofburets is always to rinse them with the solution that you will be placing in them. Use about 5-10 mL of your base solution to rinse the base buret. Make sure you rinse the sides of the buret and also rinse out the stopcock by letting liquid run through. Discard the rinse solution. Repeat this rinse procedure with another 5-10 mL sample of your base solution. Now your buret is ready to fill up with base solution. Go ahead and fill it a little above the zero mark at the top of the buret. Open the stopcock quickly to let a few mL of the base solution drain out into the sink or a waste beaker. This step will push out any air bubbles caught in the tip of the buret. Make sure the level of the solution in the buret is at or below the zero mark on the scale (drain out some base solution if necessary). Read the initial buret volume to the nearest 0.01 mL and record this initial reading on your

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data sheet. If a drop is dangling from the tip of the buret, remove it by gently touching the tip with a beaker and letting the drop run off into the beaker.

Obtain some of the standard hydrochloric acid solution in a beaker. Make sure you write down what the precise molarity of the acid is (it should be written on the jug containing the acid). Clean and prepare your acid buret in the same way you prepared your base buret, except do the final rinses with the acid solution, not the base solution. Make sure to record the initial volume reading for the acid buret on your data sheet.

Using the buret, dispense about 25 mL of the acid solution into a clean 250 mL flask. (The flask does not have to be dry. Why not?) Add three drops of phenolphthalein indicator to this solution. Now begin to add base solution from the other buret, swirling the flask to mix in the base. At first you can add the base rapidly. As you get near the endpoint, which is indicated by the fact that the pink color of the phenolphthalein takes longer to disappear, you should use your wash bottle to rinse down the sides of the flask with a little distilled water. When you are close to the endpoint, you need to add the base solution drop by drop to the flask. You must stop the titration on the first drop of base that turns the solution pink, and the pink color persists, even after swirling the solution. The shade of pink is not important; what is important is that you stop on the first drop that causes the solution to tum pink.When you have reached the endpoint, record the final buret readings (to the nearest 0.01 mL) for both the acid and base on your data sheet. If you accidentally go past the endpoint, you can add a little more acid from the acid buret until the solution turns colorless, and approach the endpoint again.

Do two more titrations, starting with a clean 250 mL flask each time. Calculate the molarity of your NaOH solution based on the data from each of three titrations. The three values you obtain for the molarity should all be within 0.002 M of each other. If they are not, repeat the titration as necessary until you obtain three consistent results. Take the average of these three molarity values to get the final molarity of your solution. Save your NaOH solution in your locker to use in part 2.

PART 2. Determining the Molarity of an Unknown Acid SolutionObtain an acid sample from your instructor. The acid is acetic acid, but its molarity is unknown.

Following the same general procedure you used in Part 1, titrate the unknown acid with your NaOH solution. Do the titration at least three times. Calculate the molarity of your unknown acid based on the data from each of your three titrations. The average of these three molarity values is the final answer you should report as the molarity of your unknown sample.

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CHEMISTRY 22Experiment 16, Titrations Report Sheet

Name -------------

PART 1. Standardization of the NaOH Solution

Molarity of the Standard HCl Solution _

Write the chemical equation for the reaction ofHCl and

NaOH:

Titration DataTitration #1 Titration #2 Titration #3

Acid Base Acid Base Acid BaseInitial buret reading (mL)

Final buret reading (mL)

Volume used (mL)

Molarity of base solution

(show calculations below)

Average Molarity ofNaOH Solution _

Show molarity calculations here:

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PART 2. Determining the Molarity of an Unknown Acid Solution

Molarity of your NaOH Solution (your average value from Part 1)

Write the chemical equation for the reaction of acetic acid and NaOH:

Titration DataTitration #1 Titration #2 Titration #3

Acid Base Acid Base Acid BaseInitial buret reading (mL)

Final buret reading (mL)

Volume used (mL)

Molarity of acid solution

(show calculations below)

Average Molarity of Acetic Acid Unknown Solution _

Show molarity calculations here:

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Experiment 16, Titrations Additional ProblemsFor all of the problems you must show a neat, complete and logical method of calculation. Use the conversion factor (dimensional analysis) method and make sure all numbers are labeled with the appropriate units. Pay attention to significant figures and put your answer in the box provided.

1. This question refers to your calculated molarity value for your unknown acid solution in Part 2.(a) How many grams of acetic acid would be present in 1.00 L of your unknown solution?

(b) If the density of your unknown acid solution is 1.01 g/mL, what is the concentration of your unknown expressed as percent acetic acid (by weight)?

2. A 30.0 mL sample of 0.220 M sulfuric acid is titrated with a potassium hydroxide solution. If it takes26.4 mL of base solution to reach the endpoint, what is the molarity of the base solution?

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3. How many mL of 0.200 M hydrochloric acid are needed to completely neutralize 3.50 g of Al(OH)3?

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CHEMISTRY 22Experiment 17, Properties of Acids and Bases

Introduction

Acids and bases have been recognized as important chemical families for hundreds of years. At first, acids and bases were identified by their observable properties. For example, acids taste sour and they react with metals to make hydrogen gas; bases feel slippery and they neutralize acids.

As chemists in the 19th century began to understand more about chemical bonding, it became clear that acids were compounds that dissociated in water to increase the concentration of hydrogen ion, H+, and bases were compounds that increased the concentration of hydroxide ion, 011. HCl is an acid because in water it ionizes to give H+ ions and c1- ions. Sodium hydroxide is a base because it dissociates in water to give Na+ ions and OI-r ions.

Acids can be described as strong acids or weak acids. Strong acids, like HCl, ionize 100% in water to give H+ ions and er ions: HCI<aql woo/o H\aql + cr<aq). Weak acids, like acetic acid, only ionizepartially in water: HC2H3Oz(aq) H\aq) + C2H3O- 2a( q)· Most of the acid molecules stay "stucktogether", only about 1% of them ionize. Similarly, bases can be described as strong or weak depending on the degree to which they ionize to give OI-r ions.

Some compounds don't contain H+ or OI-r themselves, but they react with water to make either acids or bases. The oxides of nonmetals are called acid anhydrides because they react with water to make an acid. For example, sulfur trioxide is an acid anhydride. When mixed with water, the following reactionto make sulfuric acid occurs: SO3g(

l + H2O - H2SO4a( q ) . The oxides of most metals are basic

anhydrides because they react with water to make a hydroxide compound. For example, potassium oxide is a basic anhydride. When mixed with water, the following reaction occurs:K2O(s) + H2O - 2KOH(aq) .

Certain compounds change color as the acidity of the solution changes. These compounds are called pH indicators. Indicators have been used by chemists for at least 700 years to measure the acidity of solutions. In this experiment you will investigate two of these indicators, litmus (a dye that naturally occurs in several species of lichens) and phenolphthalein.

Procedure

PART 1. Behavior of IndicatorsThe purpose of this section is to determine the colors of two common indicators, litmus and

phenolphthalein, in both acidic and basic solutions.Place 5 mL of distilled water in a large test tube. Do each of the following additions, in order, to the

test tube. After each substance is added, mix it in thoroughly and then record the color you see on your data sheet. 1) Add 10 drops of litmus solution. 2) Add one drop of 6M HCl. 3) Add two drops of 6M NaOH. 4) Add two drops of 3M H2S O4. 5) Add two drops of 6M NH4OH.

Place 5 mL of distilled water in another large test tube. Repeat the sequence of additions you didabove, but use 3 drops of phenolphthalein solution instead of the litmus solution.

17-2

PART 2. Preparation and Properties of Acids.A. Check out a delivery tube from the stockroom. Place 2 marble chips (calcium carbonate) into a

large test tube. Into a second large test tube, put 10 mL of water and 10 drops of litmus solution. Clamp the test tube containing the marble chips onto a ring stand using a utility clamp. Add 2 mL of 6M HCl to the marble chips, then immediately stopper with the delivery tube. Now move the test tube containing the litmus solution into position so that the carbon dioxide gas coming out from the delivery tube will bubble through the litmus solution. Let the gas bubble through the litmus solution until a color change is observed.

When you clean up, don't leave the marble chips in the sink; rinse them off and put them in the "usedmarble chips" container.

B. Put 2 rnL of 6M HCl into a large test tube and 2 mL of 6M acetic acid into another large test tube. Obtain two pieces of mossy zinc of about the same size. Drop one piece of zinc into each of the two acids. Compare the rate of the reactions by observing the rate of evolution of hydrogen gas.

When you clean up, don't leave pieces of zinc in the sink; rinse them off and put them in the "used zinc" container.

C. Put 2 rnL portions of 6M HCl into two large test tubes. Put about 150 mL of room temperature water into a 250 rnL beaker. This is your cool water bath. Put about 200 mL of water into a 400 mL beaker and heat until it boils gently. This is your hot water bath. Put one test tube of hydrochloric acid into the hot water bath and one test tube of acid into the cool water bath. Wait about 5 minutes for the temperature of the acid to equal the temperature of the water bath. Obtain two pieces of zinc foil of about the same size. Drop one zinc square into each of the test tubes of acid. Compare the rates of the reactions.

When you clean up, don't leave pieces of zinc in the sink; rinse them off and put them in the "used zinc" container.

PART 3. Preparation and Properties of Bases.A. Bring a 4ry watch glass up to your instructor to obtain a piece of sodium metal. Do not touch the

sodium metal with your fingers- use your tongs or a stirring rod to move it around. Fill a 150 mL beaker about half full of water and add a few drops of phenolphthalein. Drop the piece of sodium into the beaker and observe the reaction. When the reaction is finished, take a few drops of the solution out of the beaker and rub it between your fingers. What does it feel like? Wash your fingers with water right after doing this.

B. Obtain a small piece of calcium metal from the reagent shelf and drop it into a large test tube containing about 10 mL of water and a few drops of phenolphthalein. Note any signs of a chemical reaction.

C. Add 1/8 of a spoonful of calcium oxide to 10 mL of water in a 50 mL beaker. Stir the mixture with your stirring rod for about 30 seconds. Feel the outside of the beaker to see if heat is being given off or absorbed. Add a few drops of phenolphthalein to the solution.

D. Put 2 mL of O. lM ferric chloride solution in a large test tube. Add 1 mL of 6M NaOH to the test tube and mix thoroughly. Record your observations on your data sheet. Now add 2 mL of 6M HCl to the test tube and mix thoroughly.

17-3

CHEMISTRY 22Experiment 17, Acids and Bases Report Sheet

Name -------------

PART 1. Behavior of IndicatorsFill in the co1ors you observe.d m the fio11owmgtable:

Litmus PhenolphthaleinIn water

After adding HCl

After adding NaOH

After adding H2S04

After adding NH40H

Summarize the colors of the indicators:

Color of litmus in acid ------

Color of litmus in base

------

Color of phenolphthalein in acid _

Color of phenolphthalein in base _

PART 2. Preparation and Properties of Acids.A. Write the chemical equation for the reaction between the marble chips (calcium carbonate) and HCl to generate carbon dioxide gas:

What color did the litmus turn when the carbon dioxide gas was bubbled through it? _

Did the carbon dioxide make the solution acidic or basic? (circle one)

Write the chemical equation for the reaction between carbon dioxide and water that happened when the CO2 was bubbled through the litmus solution:

B. With which acid did the zinc metal react faster, hydrochloric acid or acetic acid? (circle one)

Explain why one of the acids reacted faster. (Hint: it has to do with strong versus weak acids).

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C. Did the zinc metal react faster with the acid at the higher or lower temperature? (circle one)Explain why.

PART 3. Preparation and Properties of Bases.A. What did you observe when the sodium metal was placed in the water?

What was the color of the phenolphthalein in the solution after the reaction? _

How did the solution feel when rubbed between your fingers?

Write the chemical equation for the reaction of sodium metal and water:

B. What did you observe when the calcium metal was placed in the water?

What was the color of the phenolphthalein in the solution after the reaction? _

Which metal reacted more vigorously with the water, calcium or sodium? (circle one)

Write the chemical equation for the reaction of calcium metal and water:

C. What was the color of the phenolphthalein in the solution after the reaction? _

During the reaction, was heat given off or absorbed? (circle one)

Write the chemical equation for the reaction of calcium oxide and water:

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D. What did you observe when the ferric chloride solution was mixed with the NaOH solution?

Write the chemical equation for the reaction between FeCh(aq) and NaOH(aq):

What did you observe when hydrochloric acid was added to the mixture containing the precipitate?

Write the chemical equation for what happened when you added the HCl to the precipitate:

Additional Questions.

1. Each of the following substances reacts with water to give an acidic or a basic solution. For each substance, indicate whether it would give an acidic or basic solution. Then write the chemical equation for the reaction of the substance with water. (Hint: all of these reactions are similar to ones you did in the experiment. Also, there is helpful information in the introduction of the experiment.)

a) sodium oxide acidic basic (circle one)

b) sulfur dioxide acidic basic

c) potassium metal

d) diphosphorus pentoxide

acidic basic

acidic basic

2. What are three properties that acids have in common?

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3. What are three properties that bases have in common?

4. The [H+] (H+ concentration) in a solution is 1 x 10-8 M. What is the pH of the solution?

5. The [H+] in a solution is 0.01 M. What is the pH of the solution?

6. If the pH of a solution is 11.0, what is the [H+]?

For the following problems, show a method of calculation.7. If the pH of a solution is 4.0, what is the [Oir]?

8. What is the pH of a solution with a [H+] of 4.9 x 1o-6 M?

9. What is the pOH in a solution with a pH of 12.18?

10. What is [Oir] in a solution with a pH of2.85?

"18-1

CHEMISTRY 22Experiment 18, Ethanol Distillation

Introduction

The fermentation of sugar to make ethanol (ethyl alcohol) is one of the oldest chemical reactions known to man. There are beer making recipes in ancient Sumerian writings dating back to 3000 B.C. In 1857 French chemist Louis Pasteur concluded that yeast was involved in the fermentation process.German chemist Eduard Buchner, who won the 1907 Nobel prize, isolated the enzyme from yeast, zymase, that is responsible for fermentation.

The overall chemical equation for alcoholic fermentation is:C6HJ206 yeast ) 2CzH50H + 2C02

The yeast takes a six carbon sugar, like glucose, and turns into two molecules of ethanol and two molecules of carbon dioxide. In the process, the yeast obtains energy that it needs to live and reproduce. Fermentation to produce ethanol only happens in the absence of oxygen (anaerobic conditions). If oxygen is present, the yeast will metabolize the sugar completely to carbon dioxide. In our experiment, we will start with the sugar sucrose, a twelve carbon sugar. Before fermentation happens, the yeast will break down the sucrose into the two six carbon sugars that is composed of, glucose and :fructose.

The ethanol fermentation reaction is used by humans for two main purposes. It is used to make alcoholic beverages to drink, and it is also used to make bread rise (in this case it is the carbon dioxide that is the desired product). Over one thousand years ago, chemists in the Middle East discovered that the ethanol could be concentrated and purified by the process of distillation. In a natural fermentation mixture, like beer or wine, the maximum alcohol content is about 15%. When the alcohol level gets higher than this, the yeast go dormant. Using distillation a solution that contains a much higher alcohol content can be obtained. In fact, with repeated distillation, a solution of 95% ethanol is possible.The concentration of alcohol in distilled spirits is often expressed as the proof of the solution. The proof is equal to twice the percent alcohol by volume. For example, vodka that is 140 proof is 70% alcohol.

In this experiment ethanol will be made by the fermentation of sucrose. After the fermentation is complete, which will take at least a week, the mixture will be distilled. You will collect several fractions of the distillate and determine the percent alcohol in these fractions by measuring the densities of the solutions. It is expected that the first :fraction you collect should have the highest percent alcohol, and then the percent alcohol in the distillate will decrease as you continue the distillation.

Procedure

PART 1. Preparation of the Fermentation MixtureFill a 250 mL beaker about half full with sucrose(sugar). Transfer the sugar to a 400 mL beaker and

add water until the beaker is almost full. Stir to dissolve all the sugar. Pour the sugar water into a 500 mL bottle. Add one scoop of dry yeast and 35 mL of yeast nutrient solution. Add more water as necessary to bring the level up to the shoulder of the bottle. Mix the contents of the bottle gently.Obtain a one-holed rubber stopper that fits your bottle (you can trade stoppers at the stockroom). Put a piece of tape over the hole in the stopper and then use a pin or a sharp pencil to poke a tiny hole in the tape. The purpose of this is to keep oxygen out of the bottle, but let the carbon dioxide gas escape. Put the stopper in the bottle and store the fermentation mixture in your locker.

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PART 2. DistillationCheck out the distillation equipment from the stockroom and set it up like you did in experiment 3.

You will be using a thermometer to measure the temperature throughout the distillation. When you put the stopper in the distilling flask, the bulb of the thermometer should be at the same height as the side arm of the distilling flask.

Do not stir up your fermentation mixture, let the yeast stay settled at the bottom of the bottle. If you have mold floating on the top of the mixture, you can try to remove it. Pour about half of the liquid into the distilling flask. Add two boiling chips (marble chips). Before you begin the distillation, weigh one clean, dry, empty 10 mL graduated cylinder and two 50 mL graduated cylinders. Record the masses on your data sheet and be sure to remember which cylinder is which.

Begin heating the liquid in the distilling flask and tum on the water flow through the condenser. Put the weighed 10 mL graduated cylinder in position to collect the distillate. When the mixture begins to boil, lower the heat. It is important that the mixture is boiled gently. If you boil it too vigorously, you will get poor results. When the first drop of distillate falls into the graduated cylinder, record the temperature shown by the thermometer in the distilling flask. Continue to collect distillate in the cylinder until the volume is 10.0 mL, or slightly less. This is distillate fraction #1. Remove the 10 mL cylinder and replace it with a 50 mL cylinder to collect distillate fraction #2. Each time you switch cylinders, record the temperature on your data sheet. You should collect about 20 mL of distillate in fraction #2, then switch to the last cylinder to collect fraction #3, which should also be about 20 mL. After collecting fraction #3, you can stop the distillation.

Once the distillate fractions in the graduated cylinders have cooled completely, weigh them and record the masses on your data sheet. Compute the densities of each of the three fractions. Use the table provided to determine what the percent alcohol is in each of your fractions.

Pour a few mL of fraction #1 into your large watch glass. Light it on fire using a burning wood splint. If you can't see the ethanol burning (ethanol burns with an almost colorless flame), hold your hand over the watch glass to see if you can feel the heat coming off.

Densities of Ethanol and Water Solutions at 20°C% ethanol density

(ct/ml)% ethanol density

(ct/ml)% ethanol density

(g/ml)2.0 0.995

0.9910.9880.985

0.9820.9790.9770.974

0.9710.9690.9660.963

0.9600.9570.9540.9500.947

36.038.040.042.0

44.046.048.050.0

52.054.056.058.0

60.062.064.066.068.0

0.9430.9390.9350.931

0.9270.9230.9180.914

0.9100.9050.9000.896

0.8910.8870.8820.8770.872

70.072.074.076.0

78.080.082.084.0

86.088.090.092.0

94.096.098.0

100.0

0.8680.8630.8580.853

0.8490.8440.8390.834

0.8280.8230.8180.813

0.8070.8010.7950.789

4.06.08.0

10.012.014.016.018.020.022.024.026.028.030.032.034.0

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CHEMISTRY 22Experiment 18, Ethanol Distillation Report Sheet

Name -------------

Data from DistillationMass of empty graduated cylinder

Temp. at the beginning of collecting the fraction

Volume of distillate in the fraction

Mass of graduated cylinder plusdistillate

Mass of distillate in the fraction

Fraction #1(collect about 10 mL)Fraction #2 (collect about 20mL)Fraction #3(collect about 20mL)

Density Calculations. Calculate the density of the distillate in each of the three fractions. Show your calculations below and put your answers in the table. Use the table on page 2 of this experiment to detennine what the percent alcohol in each fractions is.

Fraction #1: Fraction #2:

Fraction #3:

Calculated Density (g/mL) % Ethanol (by volume)Fraction #1

Fraction #2

Fraction #3

(over)

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Additional Questions:

1. Did the liquid in fraction #1 bum? Yes No

2. Ethanol has a boiling point of 80°C. The rest of the fermentation mixture is mostly water, which has a boiling point of 100°C.

a) As you progress through the distillation of a fermentation mixture, what happens to the percent ethanol in the fractions collected? (check one)

The percent ethanol in the distillate fractions gets lower and lower as you continue withthe distillation

The percent ethanol in the distillate fractions gets higher and higher as you continue withthe distillation

b) Explain why the first fraction collected will have the highest concentration of ethanol.

SA 1 - 1

CHEMISTRY 22Study Assignment 1, Significant Figures and Scientific Notation

Name -----------

1. Determine the number of significant figures in the following values:

a) 190.82 mL f) 8 lions

b) 0.0065 g g) 5.20 X 1014 pg

c) 95.00 mm h) 3.07 x 1011 gal

d) 5300 nm i) 6.0300 X 105 m

e) 732.010 lb j) 3215 miles

2. Round the following values to 3 significant figures:

a) 3.68801 f) 0.0421975

b) 3.68493 g) 150.7256

c) 3.68529 h) 0.013877

d) 17.048192 i) 4.86250 X 109

e) 8.75430 j) 4.0098 X 10-12

3. Perform the following unitless calculations and round the final answer to the proper number of significant figures. Assume that all of the numbers came from measurements.

a) 3.41 - 0.086652 = -------

b) 17.441 +3 = _

c)21.01 x2.0= _

f) 0.87 + 4.061 + 10.4 = ------

g) 16 X 841.4 + 16.300 = _

h) 14.044 + 8.11 + 3.4 = _

d) 18.7644- 3.472 + 0.4101 = _

4. Convert the following values into scientific notation, or if given in scientific notation, convert back to a regular number:

-a) 52,000

b) 0.0006 _-

c) 789,000,000 _

d) 7.66 x 10-2 -------

e) 4.75 X 10-4 -------f) 9 X 108

-------

-g) 250,000,000,000,000,000,000 ------

h) 0.0000006509 _

i) 0.0000009548 -------

j) 7.8 X 105 _

k) 6X 10-3 -------

1) 6.022 X 104 _

8

SA 1 - 2

5. Calculate the average of the following set of values: 18.4, 12.99, 13.772 and 9.704. Round the final answer to the proper number of significant figures.

6. Perform the following mathematical operations

a) 103 x 104 = _d) 10-

13X 106 = --------

b) 10-7X 10-

12 = --------103

e) -10-5 = _

10-2c) 10-6= -------

104f) 10-13= --------

7. Perform the following operations and express each answer in scientific notation:

a) (2.0 X 107) (9.0 X 106

) = --- --

b) (3.5 x 1013) (7.1 x 10-6

) = _

c) 6. 7x l0 =4.8 x 10 -9 --------

d) 7.9X 10-15=

5. 5x l0 -7 --------

Express the sum of 1.380 kg + 7.1 g + 85 mg using the correct number of significant figures. Show a complete method. (Hint 1: Convert all to the same units. Hint 2: Arrange the numbers vertically before adding them together.)

g

SA 2-1

CHEMISTRY 22Study Assignment 2, Unit Conversions

Name ------------

Show all work with correct conversion factors and in one set-up using dimensional analysis. Make sure you include the appropriate units and significant figures.

1. Convert 0.50 m to mm.

2. Convert 22.4 L to rnL.

3. Convert 2.00 ks to s.

4. Convert 10.0 nm tom.

5. Convert 1.2 g to µg.

6. Convert 8.00 mm to cm.

7. Convert 3.22 Gg to kg.

8. Convert 3.001 cg to mg.

SA 2-2

9. Convert 5.00 m3 to cm3

10. Convert 3.5 mm3 to dm3•

11. Convert 10.0 cm to in.

12. Convert 1.50 gal to L.

13. Convert 0.800 mi to cm.

14. Convert 15.0 lb to kg.

15. Convert 25.00 mL to qt.

16. Convert 66.3 m3 to mL.

17. Convert 32.5 mi/h to cm/ms.

18. Convert 71.63 g/cm3 to mg/dm3.

.

SA 3 -

CHEMISTRY 22 Name -------------Study Assignment 3, Dimensional Analysis

Show all work with correct conversion factors and in one set-up using dimensional analysis. Make sure you include the appropriate units and significant figures.

1. You have been working at a fast food restaurant for the past 35 years wrapping hamburgers. Each hour you wrap 184 hamburgers. You work 8 hours per day. You work 5 days a week. You get paid every 2 weeks with a salary of $840.34. How many hamburgers will you have to wrap to make your first one million dollars assuming you can continue to wrap hamburgers as long as it takes (Hint: You do not need to use 35 years) (keep 3 significant figures in your final answer).

2. How many miles in 1.2 x 104 yards?

3. The speed oflight is 3.0 x 1010 cm/s. Express this speed in km/hr.

4. What is the cost in dollars of 16 onions if 3 onions weigh 1.5 lb and the price of onions is 33 cents per kilogram?

5. How many minutes will it take to drive to Los Angeles from San Francisco if an average speed of 72 mi/hr is maintained? The distance between the two cities is 405 miles.

6. What is the cost to drive from San Francisco to Los Angeles (405 mi) if the cost of gasoline is$2.34/gal and the automobile gets 8.15 mi/L?

SA 3-2

7. How many oranges are in a crate if the price of a crate is $1.60 and the price of oranges is$0.20/lb? On the average, there are three oranges per pound.

8. You're throwing a pizza party for 15 and figure each person might eat 4 slices. You call up the pizza place and learn that each pizza will cost you $14.78 and will be cut into 12 slices. You tell them you'll call back. How many dollars will the pizza party cost?

9. A sample of seawater contains 6.277 g of sodium chloride per liter of solution. How many mg of sodium chloride would be contained in 15.0 mL of this solution?

10. At a local university the students have been overdosing on caffeine to help them study for exams. However, many students have been getting quite sick from taking too much coffee and cola. It is known that 3.00 g of caffeine taken in one day is quite dangerous.A. How many cups of coffee would be too much and at the dangerous level (of 3.00 g)?

You know that coffee contains 21.5 mg caffeine per ounce and cola contains 4.20 mg per ounce and a cup is 8 oz. and a can is 12.0 oz.

B. How many cans of cola would be too much and at the dangerous level?

11. What is the area of a piece of aluminum in cm2 that has a length of 7.69 ft and a width of 1.23 ft?

SA4-1

CHEMISTRY 22Study Assignment 4, Electron Configurations Name _ and Periodic Properties

1) Arrange the following electromagnetic waves according to each description:

X-rays, radio waves, blue light, red light, microwaves, gamma rays

a) In terms of increasing wavelength.

b) In terms of increasing energy.

c) In terms of increasing frequency.

2) Draw the three different orientations of the p-orbital.

3) Determine the maximum number of electrons that each of the following sublevels can hold.a) 6sb) 5pc) 4dd) Sf

4) When an electron travels from 5d orbital to 2s orbital in an excited hydrogen atom (Circle the correct answer below)a) a photon of light is absorbed by the atomb) a photon of light is emitted by the atomc) the atom does not lose or gain energy

5) Write the complete electron configuration and orbital diagram of the following.

a) C

b) P

c) p-

SA4-2

d) Fe2+

e) Pb (you may use noble gas abbreviation)

f) U (you may use noble gas abbreviation)

6) How many valence electrons and unpaired electrons does each atom/ion in question # 5 have?

Species Number of valence electrons Number of unpaired electronsCp

FFe2+Pbu

7) Arrange the following elements according to each

description: F,Li,K,N,

a) In terms of increasing size.

b) In terms of increasing ionization energy.

8) Which element(s) belong to period 3 and have/has 2 unpaired electrons?

9) Which element has greater ionization energy than bismuth (Bi), has a larger atomic size than iodine, and has one unpaired electron?

11) Which of the following is a possible electron configuration for an excited state of oxygen atom? a) l s22s22p4 b) ls22s22p33s2 c) l s22s22p33s1

SA 5-1

CHEMISTRY 22 Name-------------Study Assignment 5, Ionic Formulas

Write the formula of each ion on the line below the name of the ion and fill in the boxes with the correct fonnula of the compound fanned by the combination of the ions. Write the names of the compounds below the fonnula.

Potassium Magnesium Copper(!) (Cuprous)

Iron(II) (Ferrous)

Aluminum

Bromide

Nitrate

Carbonate

Oxide

Sulfite

Phosphate

SA6-1

CHEMISTRY 22Study Assignment 6, VSEPR

Name------------

Molecule Lewis Structure Electronic Geometry

Molecular Geometry

Bond Angle

Example:

PBr3

I I

: .B.r /1 -·B·r:. . • r• ..

••• •

Tetrahedral Trigonal pyramidal 109°

CS2

CH2h

OF2

SiH4

ChCO

BF3

H2Se

SA6-2

Ion Lewis Structure Electronic Geometry

Molecular Geometry

Bond Angle

H30+

-Cl02

-Br03

sol-

-NH2

NH4+

C0 3-2

N02+

SA 7-1

CHEMISTRY 22Study Assignment 7, Polarity

Name -------------

1. Electronegativity.Electronegativity is a measure of the attraction an atom has for the electrons in a bond. If an atom has a high electronegativity, then it can pull the bonding electrons toward itself, shifting them away from other atoms.Electronegativity is a periodic property. That is, electronegativity changes in a predictable way on the periodic table. Use a table of electronegativities in your book to answer the following questions.

Which element has the highest electronegativity? _

What happens to the value of electronegativity as you go down a group of elements? _

In general, what happens to the value of electronegativity as you go across a period? _

Which periodic table group has the highest values of electronegativity? _

In general, do metals have high or low electronegativities? _

Why do you think there are no electronegativity values listed for the noble

gases?

Fill in the electronegativity values for the following common elements. These are useful values to know for the rest of this study assignment.

fluorine

carbon

oxygen

hydrogen

nitrogen chlorine --

Now, put away your electronegativity table. For the rest of this exercise you should be able to compare electronegativities using the periodic table trends you have written above.

Which atom in each pair has a higher electronegativity? (circle the one with the highest

value) Mg or I B or C P or As Ge or S

2. Polarity of Bonds. When two atoms share electrons they form a covalent bond. Often the electrons are not shared equally between the two atoms. Let's look at the example of an oxygen-hydrogen bond. Oxygen has a higher electronegativity, so it will pull the bonding electrons toward itself. The electron cloud shifts toward the oxygen, making the oxygen atom slightly negative. Since the electrons are pulled away from the hydrogen, this leaves the hydrogen atom slightly positive. This partial charge can be represented with the Greek letter delta (o). We put a o- symbol next to the atom that is partially

SA 7-2

negative and a c/ symbol next to the atom that is partially positive. For the oxygen-hydrogen bond, it would look like this: 6- 6+

O-H

Because the electrons are not shared equally, the bond is partially positive on one side and partially negative on the other. A bond with an uneven distribution of charge like this is called a polar bond. The larger the difference between the electronegativities of the two elements, the more uneven the sharing of electrons, and the more polar the bond.

For each of the following covalent bonds, label the atoms with the appropriate symbol (cY or 6+) to show the polarity of the bond. Remember, the 6- symbol goes on the element with the higher electronegativity.

S-0 N-C H-Cl F-I

If two atoms in a bond have the same electronegativity, then there are no partial charges and the bond is called a nonpolar bond. In this case the electrons are shared equally between the two atoms.

3. Polarity of Molecules. Now we will look at some molecules to decide if they are polar. A molecule is polar if one side of the molecule is partially positive and the opposite side is partially negative. This type of polarity is called a dipole because it has two opposite poles.

To decide if a molecule is polar we must consider two things; the polarity of the bonds in the molecule and the shape of the molecule. In section 2 above, you learned how to predict if a bond is polar.You have previously learned how to predict the shape of a molecule using the dot structure and the VSEPR model. This exercise will lead you through 3 examples of how to use these two skills to determine if a molecule is polar.

A. The carbon dioxide molecule.Draw the dot structure for the CO2 molecule: What is the shape of the molecule?

Now consider the electronegativities of carbon and oxygen. Redraw the carbon dioxide molecule below (showing its correct shape) and label each atom with a 6- or a 6+ based on electronegativity.

You can see from your picture that each bond is polar (it is partially negative at the oxygen end and partially positive at the carbon end). However, the molecule is not polar. This is because of the linear shape. The two bond dipoles point in opposite directions and cancel each other out. Carbon dioxide is a nonpolar molecule.

SA 7-3

B. The water molecule.Draw the dot structure for the H2O molecule: What is the shape of the molecule?

Now consider the electronegativities of hydrogen and oxygen. Redraw the water molecule below (showing its correct shape) and label each atom with a/:,- or a 6+ based on electronegativity.

You can see from your picture that each bond is polar (it is partially negative at the oxygen end and partially positive at the hydrogen end). The molecule is also polar. Remember, a molecule is polar if it is partially positive on one end of the molecule and partially negative on the opposite end. You can see in your picture that the water molecule is negative at the oxygen end and positive at the hydrogen end. Because of its bent shape, the bond dipoles in the molecule do not cancel each other out. Water is a polar molecule.

(over)

SA 7-4

C. The carbon tetrafluoride molecule.Draw the dot structure for the CF4 molecule: What is the shape of the molecule?

To have a better understanding of the tetrahedral shape of the molecule, you should build a model of it. Take a black wooden ball to use as the central carbon atom. Use sticks for the bonds and yellow balls for the fluorine atoms.

Now consider the electronegativities of carbon and fluorine. Redraw the CF4 molecule below (show its correct shape- you will need to use wedges and dashes, or some other visual trick, to show the 3-dimensional tetrahedral shape) and label each atom with a 6- or a a+ based on electronegativity.

You can see from your picture that each bond is polar (it is partially negative at the fluorine end and partially positive at the carbon end). However, the molecule is not polar. This is because of the symmetry of the tetrahedral shape. All the bond dipoles point in opposite directions and cancel each other out. Carbon tetrafluoride is a nonpolar molecule.The symmetry of the shape of the molecule is very important in determining polarity. It is the perfect symmetry of the tetrahedral CF4 that causes it to be nonpolar. If you were to replace one of the fluorine atoms with something else, say a hydrogen or an iodine, the molecule would then become polar. This new molecule would be lopsided- it would be different on one side than the other. Try this with your ball and stick model. The lack of symmetry would mean that the molecule would be at least somewhat polar.

SA 7-5

4. Additional PracticeNow practice determining the polarity of the following molecules. For each molecule you will: 1) draw the dot structure, 2) write in the name of the molecular shape, 3) redraw the molecule with the correct shape, showing the partial charges with a o- or a o+ and 4) decide ifthe molecule is polar or nonpolar (note: we are asking about the polarity of the whole molecule, not the individual bonds).

Molecule Dot structure Name of molecular shape Molecule drawn withcorrect shape and partial charges shown

Is molecule polar or nonpolar?

HCN

BF3

NH3

C2F2

SA 7-6

Molecule Dot structure Name of molecular shape Molecule drawn with correct shape and partialcharges shown

Is molecule polar or nonpolar?

S02

CH20

CC4

CHC13

SA 8-1

CHEMISTRY 22 Name------------Study Assignment 8, Naming & Formulas

I. Write the chemical name for each of the following compounds:

1. CaF2

2. Na2S

3. CBr4

4. HC2H302

5. Fe2(C03)3

6. Co(OH)2

7. N204

8. K2Cr201

9. Ba02

10. Cu2S04

11. Mg(HC03)2

12. Pb3(P04)2

13. PCls

14. H2S03

15. Zn(CN)2

16. Sn(N03)2

17. Ag20 18. H2Scaql 19. NH3 20. HCl04

21. Ah03 22. bOs 23. LiClO

24. S03

25. Hg2Ch

SA 8-2

IL Write the chemical formula for each of the following compounds:

1. Strontium bromide

2. Diphosphorus pentoxide

3. Sodium dihydrogen phosphate

4. Mercury(II) phosphate

5. Potassium sulfide

6. Sulfuric acid

7. Tin(IV) oxide

8. Hydroiodic acid

9. Calcium oxalate

10. Copper(II) nitrite

11. Sodium permanganate

12. Nickel(II) hydroxide

13. Xenon hexafluoride

14. Iron(II) chlorate

15. Phosphoric acid

16. Lithium peroxide

17. Magnesium nitride

18. Ammonium sulfate

19. Chromium(III) acetate

20. Aluminum bicarbonate

21. Antimony trichloride

22. Nitrous acid

23. Gold(III) chloride

24. Platinum(IV) oxide

25. Cadmium acetate

SA 9-1

-

CHEMISTRY 22Study Assignment 9, Balancing Equations

Name ------------

Balance the following chemical equations with the lowest whole number coefficients.

1. Cu + Ag2S0 4 C uS0 4 + Ag

2. Zn + 02 - ZnO

3. C2H6 +

4. KC103 - 02 - CO 2 + H20KCI + 02

5. Li + H 20 - H2 + LiOH

6. P20s + H20 - H 3P0 4

7. Sn + N2 - Sn 3N4

8. Fe( OH)3 + H2S0 4 - H 20 + Fei( S0 4 ) 3

9. H202 - H 20 + 02

10. V + 02 - V20s

11. NH4N0 3 - N20 + H20

12. Sn0 2 + C - Sn + co

13. C 2H 50 H + 02 - CO 2 + H20

14. Hg(N0 3) 2 + K3P0 4 - Hg3(P04)2 + KN0 3

15. ZnS + 02 - ZnO + S0 2

16. NH3 + 02 - HN0 3 + H20

17. Al + Br2 -18.* Cu + HN0 3 -

A lBr3

Cu( N0 3 ) 2 + N0 2 + H20

*This is a challenging one.

CHEMISTRY 22 Name -------------Study Assignment 10, Chemical Equations

1. Write balanced chemical equations for the following reactions.

a) hydrogen gas + nitrogen gas gives ammonia (NH3)

b) sodium peroxide + water gives sodium hydroxide + oxygen gas

c) carbon disulfide + oxygen gas gives carbon dioxide + sulfur dioxide

d) silver + nitric acid gives silver nitrate + nitrogen dioxide + water

2. Complete and balance the following chemical equations. Identify the reaction type as: combination, decomposition, single replacement, double replacement, or combustion.

Reaction Type

a) MgCO 3

b) Al + 0 2 -

c) Be + HF -

SA 10-2

3. Write a complete, balanced equation for the following reactions.

b) aqueous ferric sulfate plus barium hydroxide

c) magnesium metal plus silver acetate

d) hydrobromic acid plus zinc metal

e) the heating of sodium carbonate

f) chromium(III) hydroxide plus sulfuric acid

g) aluminum metal plus chlorine gas

SA 11-1

CHEMISTRY 22Study Assignment 11, Solution Problems

Name ------------

1. A solution is made by dissolving 0.565 g of potassium nitrate in enough water to make up 250. mL of solution. What is the molarity of this solution?

2. How many moles of HCl are contained in 0.600 L of 0.120 M HCl?

3. How many mL of 0.200M KI would contain 0.0500 moles of KI?

4. How many grams of H2S0 4 are contained in 2.00 L of 6.0 M H2S0 4?

5. A solution is made by mixing 34.5 g of sugar with 75.0 g of water. What is the mass percent of sugar in this solution?

SA 11-2

6. How many grams of sodium chloride are contained in 250.0 g of a 15% NaCl solution?

7. What is the molarity of a solution made by mixing 75.0 mL of 3.00 M NH4OH with enough water to give 250. mL of solution?

8. How many grams of calcium fluoride precipitate will form if25.0 mL of 0.110 M sodium fluoride is mixed with 40.0 mL of 0.125 M calcium nitrate?

9. A 5.75 M solution of sulfuric acid has a density of 1.23 g/mL. What is the concentration of this solution in mass percent?

SA 11-3

10. Below is a table that shows the solubility of four salts in water as a function of temperature. Graph this data on the sheet of graph paper provided on page 5. After you plot the data points, connect the points for each salt with a smooth line to give a solubility curve. It is helpful to use a different color for each salt. Use your graph to answer the questions(a-h) that follow.

Solubility of Certain Salts as a Function of Temperature (in grams of salt per 100 grams of water)

Temperature(°C) NaCl KCI0 13 36 26 2810 21 36 25 3120 32 36 24 3430 46 36 23 3740 64 37 21 4050 86 37 19 4360 110 37 17 4670 138 38 15 4880 38 12 5190 39 9 54100 40 6 57

a) For which salt does the solubility change the least as the temperature changes?

b) For which salt does the solubility change the most as the temperature changes?

c) Which salt is the most soluble in water 45°C?

d) Which salt becomes less soluble as the temperature increases?

e) A solution made by mixing 35 g ofKCl with 100 g of water at 70°C would be (circle one) : unsaturated saturated supersaturated

(over)

SA 11-4

The rest of the questions require you to show some work or give some reasoning.

f) How many grams of NaCl can dissolve in 650. g of water at 50°C?

g) A solution is made by mixing 55g of KNO3 with 100 g of water at 40°C. If the solution is heated to 75°C, how many additional grams of KNO3 can be dissolved before the solution becomes saturated? (Use your graph to estimate the solubility of KNO3 at 75°C)

h) A solution of Lh SO4 in 200 g of water is saturated at 0°C. If the solution is heated to 80°C, how many grams of solute will crystallize out of the solution?

-0

150

140

130

- 120 -

' I II

' II

I

I-·· I··~ ,_ ·~· -·-

'

1--·-·

·-+-1-- I-•--

I

I...Q)

+-'ctS 110$0)

0 100,--..........±::ctSU)

0)

""

-·!·-- -··-·---- "" -

- -·

I

I I

80.Q:::J

0 70Cf)

60I

H--1,-- -1

50

40

30I

20

10I

,--- ,-- """ ,-· ·- - -I0 0 10 20 o 40 5:J 60 70 80 90 100

Temperature (Celsius)

SA II-£

90

I

SA 12-1

2

CHEMISTRY 22Study Assignment 12, Net Ionic Equations

Name -------------

When sodium hydroxide solution is mixed with hydrochloric acid, the following reaction occurs: NaOH + HCI - H20 + NaCl. This equation is called a molecular equation because it depicts all the reactants and products as neutral molecules. This type of equation is fine for balancing and stoichiometry, but it does not accurately represent what the species are really like in the solution.Actually, there are no molecules of NaOH, HCl, or NaCl present in the solution; all those species exist as dissolved ions. This study assignment will teach you how to write net ionic equations, which are a useful way of describing what is actually happening in reactions that take place in aqueous solution.

The secret to writing ionic equations is to write things the way they actually exist in the solution. If a specie is a molecule, it should be written as a molecule. If a specie dissolves by breaking apart into ions, then it should be written as ions. If a specie is a gas, it should be written as a gas, etc. As you work through this study assignment you will learn whether a substance should be written as a molecule or as ions, then you will learn how to use this information to write net ionic equations.

Part 1. How to Write the Formulas of Salts in Ionic EquationsWhen salts (ionic compounds) dissolve in water, they usually break apart into the ions they are

composed of. For example, when potassium carbonate dissolves it dissociates into two potassium ionsand a carbonate ion: K C0 - 2K\aql + C0 - caq). Because there are free ions present, this solution

2 3 3

would conduct electricity, and the salt is said to be a strong electrolyte.Some salts do not dissolve in water- they are insoluble. When you mix an insoluble salt with water,

the salt does not dissociate into ions. The salt stays together and the solid settles out of the solution.So, how you write the formula of a salt in an ionic equation depends on whether the salt dissolves in

water or not. If the salt is soluble, you should write it as dissociated ions. For example, zinc chloride is soluble so it should be written as: Zn2\ aql + 2Crcaql. If a salt is insoluble, it should be written as an intact solid. For example, cupric sulfide is insoluble so it should be written as CuS(s)• The way you determine if a salt is soluble is by looking at a solubility table. There is one provided on page 4 of this exercise.

Practice this concept by completing the following table. The first two have been done for you as an example.

Salt Soluble or Insoluble? How is it written in an ionic equation?Na2S04 s 2 Na+ + so/-

AgCI I AgCl<s)CuBr2

KMn04Fe(OH)3

CNH4)2C03BaS04

Zn(N03)2Na3P04

AhS3

SA 12-2

Part 2. How to Write the Formulas of Acids in Ionic EquationsYou can assume that any acid you see in Chem 22 is soluble. How you write an acid in an ionic

equation depends on whether the acid is strong or weak.Strong acids are acids that ionize completely in water. Hydrochloric acid is a strong acid- when it

dissolves in water, 100% of the HCl molecules ionize to make H+ ions and ci- ions. There are 6 strong acids: HCl, HBr, HI, HNO3, H2SO4, and HC1O4. You should memorize these. Because strong acids are 100% ionized in water, they should be written as separated ions in an ionic equation. For example, nitric acid should be written as H\aq) + NO3-(aq) .

All other acids (besides the six above) are weak acids. Weak acids only ionize a little bit when theyare dissolved in water. Most of the acid stays "stuck together" as molecules. Because weak acids are mostly dissolved as molecules, that's how they should be written in an ionic equation. For example, acetic acid should be written as HC2H3O2(aq).

Practice this concept by filling in the following table. For each acid indicate whether it is a weak orstrong acid, then show how it should be written in an ionic equation.

Acid Strong or Weak? How is it written in an ionic equation?HBr

HNO2H3PO4

HIH2S

HC1O4H2CO3

Part 3. How to Write the Formulas of Bases in Ionic EquationsMetal hydroxides are considered bases. Because they are ionic compounds, you should follow the

solubility guidelines that you learned in Part 1 of this exercise. If a metal hydroxide is soluble in water, it dissolves by breaking apart into ions and is considered a strong base. The strong bases are all thehydroxides of periodic table group IA, plus the hydroxides of Ca+2 , Sr+2 , and Ba+2 . Because thesecompounds dissociate completely into ions when they dissolve, they should be written as separate ions in an ionic equation. For example, barium hydroxide should be written as Ba2\ aq) + 2 OH-(aq)·

Metal hydroxides that are not soluble in water do not break up into ions and should be written as intact solids. For example, aluminum hydroxide should be written as Al(OH)3(s)·

There is one weak, soluble base you will see in Chem 22, ammonium hydroxide. OH is verysoluble in water, but it dissolves mostly as molecules, only a small percentage of the molecules ionize. Because most of the dissolved ammonium hydroxide is present as molecules, it should be written as OH(aq) in an ionic equation. (It can also be written as NH3(aq)·)

Practice this concept b1y s how· mg howeach 0 f the bases be11OW Sh OUld bewr1·tten·m an ionic equation.

Base How is it written in an ionic equation?NaOH

Mg(OH)2Ca(OH)2Z n(OH)2

OH

SA 12-3

Part 4. Writing Miscellaneous Species in Ionic EquationsFor species that are neither salts, acids, nor bases, you should still write them as they actually exist.Water: Water should be written as H2O molecules.Metals: All metals in their elemental form (except for mercury) are solids, so you should write them

as solids. For example, aluminum metal is Al(s)·Gases: Gases are common participants in chemical reactions. The gases you are most likely to

encounter are the diatomic gases (H2, 0 2, N2, F2, and Ch), plus the nonmetal oxides like CO2, NO, NO2, SO2, etc. Most of these gases are not very soluble in water and just bubble out of the solution. Because of this, you should write these species as gaseous molecules in ionic equations. For example, hydrogen gas is written as H2(g) and carbon dioxide is written as CO2(g)·

Organic compounds: Organic compounds like sugars and alcohols dissolve in water as molecules, they do not ionize. Because of this, they should be written as aqueous molecules. For example, ethanol in water exists as C2H5OH(aq).

Part 5. Writing Ionic EquationsNow you will use the skills you learned in the first 4 parts of this exercise in order to write total ionic

equations and net ionic equations. Remember, the primary rule for writing ionic equations is to write things they way they really exist in the solution.

Here are the steps you will take when writing a net ionic equation for a reaction:1. Write a balanced molecular equation for the reaction (this is just the regular kind of equation you've been doing all through this course).2. Rewrite your equation, but now write each reactant and product they way they really exist in the solution. If a specie is dissociated into ions, write it as ions. If it is a molecule, write it as a molecule. This step gives you the total ionic equation.3. Look at your total ionic equation to see ifthere are any species that don't change at all during the reaction; that is, they are exactly the same on both the left and right sides of the equation. These species are called spectators and should be canceled out. After canceling out any spectators, you will have the net ionic equation for the reaction.

Let's apply these steps to an actual reaction, the reaction between acetic acid and sodium hydroxide in aqueous solution.

1. We are mixing an acid and a base, so we can predict that the products of the reaction will be water and a salt. The molecular equation for the reactions is:

HC2H3O2 + NaOH _, H2O + NaC2H3O2

2. Now we will look at each specie and write it as it actually exists. HC2H3O2 is a weak acid, so it should be written as dissolved molecules. NaOH is a soluble metal hydroxide (a strong base), so it should be written as separate ions. Water should just be written as water molecules. NaC2H3O2 is a soluble salt, so it should be written as separate ions. This gives us the total ionic

equation:

3. If you look at the total ionic equation, you can see that the sodium ion doesn't change, it is a spectator ion. We can cancel out the Na+ from both sides of the equation to get the net ionic equation:

SA 12-4

Now you should try these steps on the reaction between magnesium metal and hydrochloric acid.1. Write the balanced molecular equation for the reaction. This reaction is a single replacement (metal+ acid), so you should be able to predict what the products of the reaction are.

2. Now look at each specie in your equation and write them as they actually exist. Refer back to parts 1-4 of this exercise as necessary. This will be your total ionic equation:

3. Look at your total ionic equation to see if there are any spectator ions. Cancel out any spectator ions, and check again to see if your equation is balanced with the lowest whole number coefficients. This will be your net ionic equation for the reaction:

Congratulations! (you can check your answer on page 6)Now practice these skills by doing the equations that start on the next page.

Simplified Solubility Guidelines for Ionic CompoundsSOLUBLE COMPOUNDS EXCEPTIONSSalts of NH+4 and periodic group IA none 2Salts of Cl-, Br-, and -I Chlorides, bromides, and iodides of Ag+, Pb+ , and

Hg/+ are insolubleSalts ofN03-, C2H302-, Cl03-, and Cl04- noneSalts of so/· Sulfates of Ba+2 and Pb+2 are insoluble

INSOLUBLE COMPOUNDS EXCEPTIONSSalts of co/·, Po/·, and 0 2· Carbonates, phosphates, and oxides of +

and periodic group IA are solubleSalts of s 2· and OH- Sulfides and hydroxides of +, all of periodic

group IA, and group IIA from Ca+2 and below aresoluble

summary ofH ow S;pecies are Wn·tten. mIo.mc Equatrnns

SALTS(Ionic Compounds) Soluble Write as ionsInsoluble Write as intact solids

ACIDS Strong (HCl, HBr, HI, HN0 3, H2S04, HCl04)

Write as ions

Weak Write as aqueous moleculesBASES Soluble metal hydroxides

(strong bases)Write as ions

Insoluble metal hydroxides Write as intact solidsAmmonium hydroxide(weak base)

Write as aqueous molecule

SA 12-5

CHEMISTRY 22Study Assignment 12, Net Ionic Equations

Name ------------

For each of the following reactions in aqueous solution, write (a) the molecular equation, (b) the total ionic equation, and (c) the net ionic equation. If products are not given, you should be able to predict the products because it is a common reaction type. Be sure to indicate the state, (aq), (g), or (s), for all neutral species. All ions are assumed to be aqueous.

1. nitrous acid plus potassium

hydroxide a)

b)

c)

2. lead(II) nitrate plus sodium

sulfate a)

b)

c)

3. zinc hydroxide plus nitric

acid a)

b)

c)

4. barium hydroxide plus hydrobromic acid

a)

b)

c)

5. sodium phosphate plus cupric

chloride a)

b)

SA 12-6

c)

2

SA 12-7

6. ammonium hydroxide plus acetic

acid a)

b)

c)

7. aluminum metal plus hydrochloric

acid a)

b)

c)

8. nickel(II) iodide plus calcium

hydroxide a)

b)

c)

9. silver carbonate plus nitric acid gives silver nitrate plus carbon dioxide plus

water a)

b)

c)

answer to equation on page 4: Mg<s> + 2H\aql - H< gl + Mg2\ aql

Periodic Table of the Elements

IA VIIII

H1.008 IIA IIIA IVA VA VIA VIIA

2

He4.00

3

Li6.94

4

Be9.01

5

B10.81

6

C12.01

7

N14.01

8

016.00

9

F19.00

10

Ne20.18

11

Na22.99

12 13

Al26.98

14

Si28.09

15

p30.97 s16

32.06

17

Cl35.45

18

Ar39.95

Mg2431

19

K39.10

20

Ca40.08

21

Sc44.96

22

Ti47.88

23

V50.94

24

Cr52.00

25

Mn54.94

26

Fe55.85

27

Co58.93

28

Ni58.69

29

Cu6355

30

Zn6539

31

Ga69.72

32

Ge72.61

33

As74.92

34

Se78.96

35

Br79.90

36

Kr83.80

37

Rb85.47

38

Sr87.62

39 40

Zr9122

41

Nb92.91

42

Mo95.94

43

Tc(99)

44

Ru101.07

45

Rh102.91

46

Pd106.42

47

Ag107.87

48

Cd112.41

49

In114.82

50

Sn118.71

51

Sb121.75

52

Te127.60

53

I126.90

54

Xe13129

y88.91

55

Cs132.91

56

Ba13733

57

La138.91

72

Hf178.49

73

Ta180.95

74

w183.85

75

Re18621

76

Os19023

77

Ir19222

78

Pt195.08

79

Au196.97

80

Hg20059

81

Tl20438

82

Pb207.2

83

Bi208.98

84

Po(209)

85

At(210)

86

Rn(222)

87

Fr(223)

88

Ra(226)

89

Ac(227)

104

Rf(261)

105

Db(262)

106

Sg(263)

58

Ce140.12

59

Pr140.91

60

Nd14424

61

Pm(147)

62

Sm15036

63

Eu151.97

64

Gd15125

65

Tb158.92

66

Dy16250

67

Ho164.93

68

Er16126

69

Tm168.93

70

Yb173.04

71

Lu174.97

90

Th232.04

91

Pa(231) u92

238.03

93

Np(237)

94

Pu(242)

95

Am(243)

96

Cm(247)

97

Bk(247)

98

Cf(249)

99

Es(254)

100

Fm(253)

IOI

Md(258)

102

No(253)

103

Lw(257)

Al

ION NAMES AND FORMULAS

Positive Ions (cations)

The charges of these ions can be figured out using the periodic table:lithium Ltsodium Na+

potassium K+

magnesium Mg2+

calcium Ca2+

barium Ba2+

aluminum Al3+

These ions must be memorized: hydrogen H+

ammomum NH4+

silver Ag+

zinc Zn2+

copper(!) or cuprous cu+

copper(II) or cupric Cu2+

mercury(!) or mercurous Hg/+

mercury(II) or mercuric Hg2+

lead(II) Pb2+

iron(II) or ferrous Fe2+

iron(III) or ferric Fe3+

manganese(II) Mn2+

nickel(II) Ni2+

tin(II) or stannous Sn2+tin(IV) or stannic

Sn4+ chromium(III) or chromic

Cr3+

A2

Negative Ions (anions)

The charges of these ions can be figured out using the periodic table:fluoride p-

chloride c1-

bromide Br-

iodide 1-

oxide oz-

sulfide sz-

These ions must be memorized: acetate

C2H3O2-

hydroxide OH-

cyanide CN-

hypochlorite c1O-(or ocn

chlorite ClO2-

chlorate ClO3-

perchlorate ClO4-

nitrite NO2-

nitrate NO3-

permanganate MnO4-

carbonate co/-

bicarbonate

or hydrogen carbonate

HCO3- chromate

Crol-dichromate Cr2ol-

oxalate C2O/-

peroxide oz2-

sulfite SO3- 2

sulfate sol-bisulfate

or hydrogen sulfate

HSO4- phosphate

PO/-

monohydrogen phosphate HPO/-

dihydrogen phosphate H2PO4-

Solubility of Salts

'1-<i::q

'"' 0M

::i::u"'-u'

N,...,

0u

N"<t

u 0' M

0z"''"<t0/:).;

Nr:fJ

N'<t0r:fJ

Ag+ I s I I I I I s I I sA l 3+ s s s s I s I I sBa2+ s s s I I s s s I s ICa2+ s s s I I s s s I s sCo2+ s s s I I s I s I I sCr3+ s s s s I s I I sCu2+ s s s I I s I s I I sFe2+ s s s I s I s I I sFe3+ s s s I I s I I sHg2+ s s s I I I I s I sK+ s s s s s s s s s s sMg2+ s s s I s s I s I I sMn2+ s s s I s I s I I sNa+ s s s s s s s s s s sNH/ s s s s s s s s s s sNi2+ s s s I I s I s I I sPb2+ I s I I I I I s I I IZn2+ s s s I I s I s I I s

S = Soluble in water I = Insoluble in water No Entry= Unstable or decomposes

Water Vapor Pressure as a Function of Temperature

Temp (OC)

VaporP (torr)

Temp (OC)

VaporP (torr)

Temp (OC)

VaporP (torr)

Temp (OC)

VaporP (torr)

0 4.6 18 15.5 24 22.4 30 31.85 6.5 19 16.5 25 23.8 40 55.310 9.2 20 17.5 26 25.2 50 92.515 12.8 21 18.6 27 26.7 60 149.416 13.6 22 19.8 28 28.3 70 233.717 14.5 23 21.1 29 30.0 80 355.l

A3