learn by doing - chemistry subject

Chemistry Book of Experiments Progress Education Tanzania 1

LEARN BY DOING

A CHEMISTRY BOOK OF LABORATORY PRACTICES AND EXPERIMENTS

FOR STUDENTS AND TEACHERS

FOR SECONDARY SCHOOL EDUCATION

Progress Education Tanzania 2009


Introduction Learning science studies is not a narrative process. It is not just enough to read the concepts and ideas and understand them. It is not just being knowledgeable about science theory. Science is a discipline that must be learnt by also participating and doing what you see in theory. You can not just count yourself as a scientist by filling your head with concepts and a lot of information without being able to apply them and bring about the results and outcomes the society need in solving and providing solutions to many fields of life. In Tanzania, science studies lag behind in performance by students than other subjects taught in our secondary education. Even those considered as taking science subjects and who happen to proceed to higher education are not fully equipped with the competency needed to make them valuable to the future national development. This is because they lack necessary practices of implementing the ideas they learn since secondary level. Many secondary schools dont have laboratories and their associated equipment for science studies. Many secondary schools lack science books that can encourage students to take science studies. There many shortcomings that needs to be addressed if we are to make progress in science. As part of a solution to minimize the problem, we, as Progress Education team have come out with this short description of some guiding experiments that are necessary for chemistry studies. They cover all the topics that involve practicing and doing experiments. It is our hope that this book can help and encourage a lot of students and teachers to engage in laboratory activities for science development.


TABLE OF CONTENTS

ACIDS & BASES 1 ............................................................................................ 4

ACIDS & BASES 2 ............................................................................................ 8

CHEMICAL KINETICS .................................................................................... 14

ELECTROCHEMISTRY 1 ................................................................................ 18



THERMODYNAMICS ..................................................................................... 27

HEAT OF SOLUTION OF A SOLID ................................................................... 30

CHEMICAL EQUILIBRIUM .............................................................................. 36

SOLUTIONS ................................................................................................. 41


SOLUTIONS ................................................................................................. 54



ACIDS & BASES 1

TESTING FOR ACIDS AND BASES

OVERVIEW Acids and bases have a number of characteristic properties. Bases, for example, typically have a slippery feel and a bitter taste whereas acids are typically sour. Tasting and feeling unknown substances in order to determine whether or not a substance might be acidic or basis is generally not a recommended approach however. Measuring pH is often a more useful way to classify a substance as an acid or base. pH, which measures H+ concentration (more specifically H3O

+ concentration) can be determined using chemical indicators. An indicator is a dye that will change colour at different pH values. By noting the colour of the indicator in a solution with an unknown pH and comparing it to the indicators known colour range, the pH of the solution may be determined. In this lab you will test the pH of a number of common substances. Another property of both acids and bases is that they neutralize each other. During this lab you will examine how a neutralization reaction may be detected by a commonly used indicator, phenolphthalein. PURPOSE To identify common substances as either

an acid or base using chemical indicators

To prepare an indicator solution and use it to determine pH

To observe how a neutralization reaction may be detected by an indicator

SAFETY Use caution when testing the HCl and NaOH solutions and the household cleaning products. All are highly corrosive substances. Avoid contact with eyes and skin. Goggles should be worn.

EQUIPMENT AND MATERIALS

Your teacher will provide you with several different indicators.

Some suggested indicators:

Ph paper Litmus paper both red and blue Phenolphthalein solution Bromothymol blue Universal indicator solution Red cabbage leaves

Suggested Solutions to Test

Antacids Apple juice Baking soda solution Clear soft drink (e.g.

Sprite) Coffee Drano

Methanol Milk Orange juice Oven cleaner Pepto Bismol Saliva Sodium chloride


Ethanol Household ammonia Javex Ketchup Lemon juice

solution Tea Tomato juice Vinegar Window cleaner

1M HCl 0.5M NaOH

Other Equipment Spot plate

Small test tubes Dropper pipettes, 2 250 ml beaker to use for water bath Bunsen burner Electronic ph meters if available

PROCEDURE Part A. Using Indicators to Determine

pH 1. Prepare a data chart in your notebook to

record your results. Include a column for each indicator you have available to you and list all test solutions you will be using. See Table 1 as a sample.

2. Place a few drops of one of the test

solution into the wells on the spot plates. Use a separate well for each of the indicators available to you.

3. For the indicator papers, dip the paper

into the solution and record the colour. For the pH paper, use the colour chart provided with the pH paper to determine the pH of the solution and record that value in your table.

1. For the indicator solutions, add one drop of indicator into each test well. Be careful not to let the dropper touch the drops of solution already in the test wells. Record the colour.

4. If an electronic pH meter is available, use

it to test the pH of all test solutions. Part B. Red Cabbage Indicator 1. Fill a 250-mL beaker roughly half-full

with torn red cabbage leaves. Add approximately 100 mL of water. Boil the leaves until the water turns a deep purple. Allow to cool and filter the liquid to use as your indicator.

Your teacher may have you prepare

indicator paper from your cabbage solution. To do this soak coffee filters or filter-paper in the solution. After drying, cut the filter into strips that can then be used for your tests.

2. Retest the samples you tested in Part A

using your cabbage indicator. Record your results in Table 2.

Part C. Using Phenolphthalein to test

Neutralization Reactions 1. Using a clean dropper pipette add 10

drops of 1M HCl to a clean test tube and add one drop of phenolphthalein. Note the colour in Table 3. Swirl the test tube to mix.

Test the pH of this solution using the pH

paper and record the pH. You may find it easier to test the pH by adding a few drops of the HCl to a spot plate test well.


2. Using a clean dropper pipette place a few

drops of 0.5 M NaOH to spot plate test well and add one drop of phenolphthalein. Note the colour.

Test the pH of this solution using the pH paper and record the pH.

3. Using the NaOH dropper pipette add

0.5M NaOH drop by drop to the test tube containing the HCl until a colour change occurs. Swirl the test tube to mix after the addition of each drop. Record the number of drops required, and record the new colour.

4. Once the colour change has occurred, use

pH paper to determine the pH of the new mixture.

RESULTS

Table 1. Indicator Tests

Solution

Red litmus

Blue litmus

phenol-

phthalein

pH paper

vinegar

lemon juice

tomato juice

1M HCl

0.5M NaOH

Table 2. Red Cabbage Indicator Tests

Solution Cabbage indicator colour

original cabbage solution

vinegar

lemon juice

etc.

Table 3. Neutralization Reaction

Solution phenolphthalein

colour pH

1.0M HCl only

0.5M NaOH only

After addition of NaOH to HCl causes

colour change


CONCLUSIONS AND QUESTIONS 1. List all substances you tested and classify

each as an acid or a base. Were any substances neutral (or nearly so)?

2. Consider the food items tested what

conclusion, if any, can you make concerning their pH? What conclusions can be made about the cleaning products in general?

3. What colour does the cabbage indicator

turn when in an acid? in a base? Did all results correspond well with your results for Part A?

4. Prepare a results table for the cabbage

indicator identify the pH range for each cabbage solution colour you noted during the lab. Youll need to refer back to your results for pH paper results (Table 1).

5. Summarize the change in colour of

phenolphthalein during the neutralization reaction how can it be used to indicate when the acid-base solution reaches the point where neutralization occurs?

How many drops of NaOH were required to neutralize the 10 drops of HCl?

Chemistry Book of Experiments 8

ACIDS & BASES 2

ACID-BASE TITRATION

OVERVIEW Often we want to determine the concentration of a solution. One way to do so is to carrying out an analytical procedure known as a titration. During a titration a carefully measured volume of the solution with the unknown concentration (called the analyte) is reacted with a second solution (the titrant) whose concentration is known (a standard solution). By knowing how much of the standard solution is required to react completely no more, no less with the solution with the unknown concentration we can calculate that solutions concentration. The point at which stoichiometrically equal amounts of the two solutions have been combined is called the equivalence point. When we neutralize an acid with a base, this will occur when [H+] = [OH-]. By using an appropriate indicator we can detect this point by noting when the indicator changes colour. This will be used to signify the end point of the titration. A balanced equation and simple calculations will then allow us to determine the concentration of the solution. PURPOSE To determine the concentration of a

solution of NaOH by titration with a standard solution of HCl.

To determine the concentration of a sample of white vinegar by titration with a standard solution of NaOH

SAFETY Acids and bases are corrosive substances.

Safety goggles must be worn. Be sure to report any spills to your teacher so they may be cleaned up properly.


Two 50-mL burettes Burette stand and clamps Erlenmeyer flask, 125-mL Erlenmeyer flask, 250-mL Wash bottle Distilled water 10-mL graduated cylinder 10-mL volumetric pipette (optional) 0.100M HCl standard solution NaOH solution with unknown concentration Vinegar (acetic acid, HC2H3O2) Phenolphthalein Distilled water PROCEDURE Part A. Titration of Base of Unknown

Concentration


1. Wash two burettes with detergent solution. Rinse them thoroughly.

2. With grease marking pencil or tape;

identify which burette is to hold each solution, the acid or base.

Rinse each burette with about 10 mL of solution that it is to hold rinse the acid-containing burette with the HCl solution and the base-containing burette with the NaOH solution. Allow the acid or base to run out of the burettete tip to rinse them.

3. Fill each burettete with the proper

solution and allow the some of each solution to run out of the burette tip. Make sure no drop remains hanging on the burettete. Be sure there are no air bubbles in the tips.

It is very important that you accurately

read and record the initial and final volumes. It is not necessary that the burettetes be filled to the very top mark (0.0 mL) at the start of the titration, but it is important that the level never go below the bottom mark (50.0 mL). Be sure to read the bottom of the meniscus at eye level. You may find it helpful to hold a white card with a large black streak or rectangle behind the burette to make it easier to read.

4. Place the 125-mL Erlenmeyer flask

beneath the acid burette. Add 10.0 mL of acid to the flask. Use your rinse bottle to make sure all drops make it to the bottom of the flask; rinse any drops that remain on the sides of the flask. Read the burette carefully and record both the

initial and final volumes from the burette into your data table.

5. Add 10-mL of distilled water to the flask. 6. Add three drops of phenolphthalein to the

flask, and swirl the flask to mix thoroughly.

7. Move the flask so it it beneath the base

burette. Place the flask on a sheet of white paper so a colour change will be more readily observed.

8. After recording the initial volume of base

in the burette, begin the titration by adding NaOH to the flask. For your initial trial you may want to add the base fairly quickly until you notice a pink colour appearing in the flask. Swirling the flask should make the pink colour disappear. At that point begin adding the NaOH more slowly, swirling the flask after each drop is added. As soon as a faint pink colour becomes permanent, stop the titration the end point has been reached. Do NOT continue until a darker pink colour has been reached if that happens youve gone past the end point.

If you do go past the end point, add a few

drops of acid (be sure to record the new volume used), then add more base.

Record the final volume of base in the

burette. 9. Repeat the titration, performing at least

four trials. Be sure to rinse the Erlenmeyer flaks well between trials.


For your other trials add the base more slowly as you near the end point in order to get more accurate readings. You do not need to refill burettetes between trials.

Part B. Titration of Vinegar 1. Using the volumetric pipette (or another

clean burette), add exactly 10.0 mL of vinegar to a clean 250-mL Erlenmeyer flask.

2. Add 100 mL of distilled water to the flask

and three drops of phenolphthalein. 3. Titrate the vinegar with the NaOH

solution used in Part A. If necessary add more NaOH to the burette before beginning the titration. Record the initial volume of base in the burette.

4. As before, the end point will be reached

as soon as a permanent, pale pink colour appears in the flask. Record the final volume of base.

5. Repeat the titration at least two more

times. RESULTS Copy Data tables 1 and 2, as shown on the last page of this lab, into your data notebook. CALCULATIONS Part A. Titration of Base of Unknown

Concentration

To calculate the concentration of the unknown base we must begin with a balanced equation. The reaction between hydrochloric acid and sodium hydroxide is:

HCl + NaOH NaCl + H2O Stoichiometrically we see that one mole of the acid reacts with one mole of the base. Because of this one-to-one relationship we can use the following formula to calculate the unknown concentration:

Macid Vacid = Mbase Vbase Rearrange the equation to solve for the unknown concentration of the base:

base

acidacidbase V

VMM

=

For each trial in Part A, determine the molarity of the NaOH solution, [NaOH]. Show your calculations in a table similar to the one shown below. Calculate the average for your trials.

Table 3. Calculating the concentration of the sodium hydroxide solution.

Trial Calculations [NaOH]

baseacidacid

base VVM

M

=

1


2

3

4

Average

-----

Collect the data from the rest of the class. Calculate the class average:

Table 4. Class data for the concentration of the sodium hydroxide solution.

Group [NaOH]

1

2

3

etc.

Average

Part B: Titration of Vinegar The reaction between sodium hydroxide and

vinegar acetic acid, HC2H3O2 is represented by:

HC2H3O2 + NaOH NaC2H3O2 + H2O Again there is a 1:1 relationship between the

acid and the base. As before we can determine the concentration of the unknown solution in this case the acetic acid if we know the volume and molarity of the base and the volume of the acid used:

acid

basebaseacid V

VMM

=

Using the molarity of the base you calculated in Part A of the lab, determine the molarity of the acetic acid. Show your calculations in Table 5, which you should copy into your notebook. Collect the data from the rest of the class. Calculate the class average:

Table 6. Class data for the concentration of the vinegar solution.

Group [NaOH]

1

2

3

etc.

Average


CONCLUSIONS AND QUESTIONS

1. How did the results for each of your trials for the titration of the sodium hydroxide compare? Were the results similar or did they vary a great deal?

2. What are some of the major sources

of error with this experiment?

3. The volume of water added during this experiment to rinse droplets of acid from the burettete or as water added to the acid in the flask does not affect the calculations and thus does not need to be accounted for. Why not?

Table 5. Calculating the concentration of the vinegar solution.

Trial Calculations [HC2H3O2]

acid

basebaseacid V

VMM

=

1

2

3

4

Average -----


DATA TABLES Table 1: Titration of NaOH with Unknown Concentration

Trial 1 Trial 2 Trial 3 Trial 4

HCl NaOH HCl NaOH HCl NaOH HCl NaOH

Initial volume

Final volume

Volume used

Table 2: Titration of Vinegar

Trial 1 Trial 2 Trial 3 Trial 4

Vinega

r NaOH

Vinegar

NaOH Vineg

ar NaOH

Vinegar

NaOH

Initial volume - - - -

Final volume - - - -

Volume used 10.0 10.0 10.0 10.0


CHEMICAL KINETICS

THE IODINE CLOCK REACTION

OVERVIEW The Iodine Clock Reaction is a classic

experiment demonstrating the effects of concentration and temperature on reaction rate.

In this experiment two solutions are

mixed. The reaction takes place in two steps.

Step 1:

IO3- (aq) + 3HSO3

- (aq) I-(aq) +3SO42- (aq) +

3H+(aq) Step 2:

5I-(aq) + 6H+

(aq) + IO3-(aq) 3I2(aq) + 3

H2O(l) The iodine, I2, produced in Step 2 will

react with starch (not shown in the equations), producing a deep blue-black solution.

The rate of the entire reaction can be

measured by timing how long it takes before the blue color appears once the two solutions are mixed.

By altering the concentration of one of the

reactants (Part A) and by changing the reaction temperature (Part B), the effects

of these factors on reaction rate can be determined.

PURPOSE

To measure the rate of a reaction To measure the effect of changing

reactant concentration on reaction rate

To measure the effect of changing temperature on reaction rate


SAFETY

Avoid getting either solution on your skin or clothes. Wash any splashes with cold water.

Wear safety goggles EQUIPMENT AND MATERIALS

Solution A (contains IO3- ions)

Solution B (contains HSO3- &

starch) distilled water ice cubes beakers, 250 mL and 100 mL (2) graduated cylinders, 10 mL (2) LARGE test tubes (2) thermometer water baths timer (stop watch or clock with

second hand) safety goggles lab apron recommended magnetic stirrer, optional

PROCEDURE Part A. Effect of Concentration 1. For Trial 1 measure exactly 10.0 mL of

solution A and pour into a 100 mL beaker 2. Use a different graduated cylinder to

measure exactly 10.0 mL of solution B and pour it into a second 100 mL beaker

3. You will begin recording reaction time as

soon as you first mix the two solutions. One person should record the time of reaction while the other partner mixes the

solutions. Pour one of the solutions into the other, then pour the solutions back and forth sever times to ensure thorough mixing. Then wait for the completion of the reaction.

4. Record the time at the instant the deep

blue-black colour first appears. 5. You will repeat Steps 1 4 using different

concentrations of solution A. Obtain the dilutions by adding distilled water to Solution A, according to the following table:

Trial Solution A (mL)

Distilled Water (mL)

1 10.0 0.0 2 9.0 1.0 3 8.0 2.0 4 7.0 3.0 5 6.0 4.0 6 5.0 5.0 7 4.0 6.0 8 3.0 7.0

Rinse and dry the beakers and graduated

cylinders between each trial to avoid contamination/early mixing of reactants.

If time is short, half the class may be

assigned trials 2, 4, 6, 8 and the other half should do trials 3, 5, 7. All groups should do the 1st trial. At the end of the lab data should be shared between groups.


Part B. Effect of Temperature 1. You will follow the same procedure as you

used for Trial 1 of Part A, mixing 10.0 mL of Solution A with 10.0 mL of Solution B. The concentrations will be kept constant, but you will vary the temperature of the solutions before mixing.

As with Part A, different groups of

students may be assigned different temperatures.

2. Class data should be obtained for the

following temperatures. Each group of students should do trials at four temperatures:

Set 1: 5C 15C 25C 35C Set 2: 10C 20C 30C 40C

3. Prepare water baths for your set of

solutions. Fill a 250 mL beaker about two third full with water of the appropriate temperature. For water baths below room temperature, use ice to chill the water. For warmer baths; use warm water from the tap or a kettle.

There should be enough water in the

beakers so that the solutions in the test tubes are well beneath the water level in the baths.

4. As before, measure out exactly 10.0 mL of

Solution A and B, pouring each into their own test tube. Place the two test tubes in your water baths, allowing them to remain for about 10 minutes to allow the solution

temperatures to reach the temperature of the water baths.

Use the thermometer to monitor water

bath temperatures. Try to keep the water

bath temperature within 0.5C of your assigned temperature. Add more ice or warm water as necessary. Record your water bath temperature just before mixing the solutions in the data table.

5. Once the solutions are at the desired

temperatures, prepare to record time and mix the two solutions. Pour the mixture back and forth between the two test tubes several times to ensure mixing. Then place the test tube containing the mixed solutions back in the water bath.

Record the time it takes from initial mixing

until the deep blue-black colour appears. At the end of the lab, share temperature

and reaction rate data between groups.

RESULTS Results may be recorded in the following data tables. Part A. Effect of Concentration


Trial Volume KIO3 (mL)

Distilled Water (mL)

Time to Completion (s)

1 10.0 0.0

2 9.0 1.0

3 8.0 2.0

4 7.0 3.0

5 6.0 4.0

6 5.0 5.0

7 4.0 6.0

8 3.0 7.0

Conclusions and Questions 1. Plot your results for Part A and Part B on

two graphs. Graph A plot mL Solution A vs reaction

time Graph B plot Temperature vs reaction

time 2. Based on your results what can you

conclude regarding the effect of concentration on reaction rate? Use the collision theory to explain this effect.

3. Based on your results what conclusion can you make regarding the effect of temperature on reaction rate? Explain this in terms of the collision theory.

Part B. Effect of Temperature

Trial Solution Temperature

(C)

Time to Completion (s)

1

2

3

4

5

6

7

8

Tem

per

ature

(

C)

Reaction Time (sec)


ELECTROCHEMISTRY 1

ELECTROCHEMICAL CELLS

OVERVIEW In an electrochemical cell, chemical

energy is converted into electrical energy. This is accomplished by using a spontaneous chemical reaction to generate an electric current, which we can simply define here as electrons traveling though a wire.

To create the electrochemical cell two

half-reactions will be set up in different containers. In one, an oxidation reaction will be used to generate a source of electrons. These free electrons will travel, through an external circuit, to the second container and will cause the reduction reaction to occur. The final requirement for our complete electrochemical cell will be a salt bridge that will permit ions to flow between the two half-cells, thus maintaining electrically neutral solutions.

PURPOSE

To build several electrochemical cells. SAFETY

Follow normal lab safety guidelines. There are no specific safety hazards for this lab.


250-mL beakers, 2 glass U-tube for the salt bridge cotton plugs for the salt bridge copper wires, 2, insulated, with

alligator clips steel wool to clean electrodes DC voltmeter metal electrodes: copper, zinc, lead

Solutions

0.5M Cu(NO3)2 0.5M Zn(NO3)2 0.5M Pb(NO3)3 0.5M KNO3 for the salt bridge

PROCEDURE

1. Each half-cell will be created by placing a

metal electrode in an electrolytic solution containing the same metals ions. For


example the copper electrode will be placed in a copper(II) nitrate solution.

For each electrochemical cell you create you will require two half-cells. Set these cells beside each other they will be connected by the U-tube.

2. Fill a beaker about two-thirds full of the electrolytic solution. Clean the electrode using the steel wool, then place the electrode in its appropriate solution.

3. Clip one end of each copper wire to the two electrodes using the alligator clips.

4. Fill the U-tube with KNO3 and stopper both ends with the cotton plugs. Turn the U-tube upside down and place one end in each half-cell.

5. Touch the other end of the copper wires to the voltmeter terminals. If the indicator on the voltmeter deflects in the wrong direction, switch the wires on the terminals. Read the highest voltage reading obtained youll need to do this quickly after connecting the wires to the voltmeter.

6. Repeat the experiment for other combinations of half-cells.

RESULTS Copy a data table similar to the one shown below into your lab notebook and use it to record your results.

Half-cells Voltage

Cu|Cu2+ Zn|Zn2+

Cu|Cu2+ Pb|Pb2+

Zn|Zn2+ Pb|Pb2+

CONCLUSIONS AND QUESTIONS 1. For each electrochemical cell you

created: a. Write out the two half-reactions for

each electrochemical cell you created. b. Identify each half-reaction as

oxidation or reduction. c. Identify each half-reaction as the

anode and cathode. d. Indicate the direction of the flow of

electrons

2. Using a Table of Standard Reduction

Potentials, calculate the theoretical voltage for each cell.

3. Compare the voltages you obtained with the theoretical voltage for each cell. What are some reasons that would account for any differences?


ELECTROCHEMISTRY 2

ELECTROLYSIS OF WATER

OVERVIEW During electrolysis, electrical energy is

used to cause a non-spontaneous chemical reaction to occur. Electrolysis is often used to obtain elements that are too chemically reactive to be found free in nature.

In this experiment electrolysis will be

used to separate water into hydrogen gas and oxygen gas. During this experiment you will perform certain tests for the products of each of the half-reactions involved in the process.

Reduction will occur at the cathode. At this electrode hydrogen gas and hydroxide ions are formed. The electrons required for this reduction will come from the power source.

4H2O + 4 e- 2 H2 + 4 OH-

Oxidation will occur at the anode,

producing oxygen gas and hydrogen ions. The electrons that are produced will return to the power source:

2H2O O2 + 4 H+ + 4 e- Adding the two half-reactions together gives us a net reaction of:

6 H2O 2 H2 + O2 + 4 H+ + 4 OH- The H+ and OH- that are produced will combine to form 4 H2O.

6 H2O 2 H2 + O2 + 4 H2O Finally we can simplify our overall equation to:

2 H2O(l) 2 H2 (g) + O2 (g) Two alternatives methods are given in this lab. One involves using a Brownlee electrolysis apparatus. If one is not available it is not difficult to assemble your own using common lab equipment. Using this method allows you to collect hydrogen and oxygen gases. The second method using simpler materials, but does not provide a way to collect the two gases.

PURPOSE

To use electrolysis to separate water into hydrogen and oxygen gas.

SAFETY

Small amounts of explosive hydrogen gas will be generated. Safety goggles


should be worn when testing for the presence of hydrogen gas.


Option 1: Brownlee Electrolysis Apparatus

Brownlee electrolysis apparatus D.C. power source or commercial

battery (9V) phenolphthalein indicator vinegar test tube holder candle wood splint

Create your own Brownlee

apparatus

insulated copper wire with alligator clips on one end

large glass jar 2 test tubes test tube clamps and ring stands to

hold the inverted test tubes in place

Option 2

In this method, easily set up as a home

experiment, hydrogen gas and oxygen gas will not be collected.

two 250 mL beakers or glass jars coffee filter insulated copper wire (alligator

clips optional) salt water solution (approximately

8 teaspoons of table salt, NaCl, in 500 mL of water)

vinegar phenolphthalein

PROCEDURE Option 1: Brownlee Electrolysis

Apparatus 1. Fill the large beaker approximately three

quarters full with water. 2. Fill both test tubes completely with

water. You may find the next step easier if you

cover the open ends of the test tube with a small piece of paper.

Invert the test tubes, placing the open

tops under water in the large jar. Both test tubes should be completely full of water with no bubbles. Allow the paper you used to cover the test tubes to drop off under water.


3. Clamp the inverted test tubes in place in the Brownlee apparatus. The electrodes will be inside the test tubes.

If creating your own set up, place the free

ends of the copper wire inside the test tubes, one wire for each tube. The alligator clips will be used to attach the wire to the power source.

4. Add a few mL of vinegar to the water,

which will help conduct the electric current.

5. Connect the wires to the power source to

begin the electrolysis. 6. Record your observations. Continue the

electrolysis until several cm3 of gas have been allowed to collect in each of the two test tubes. Note the relative amount of gas that collect in each tube. Be sure to indicate which test tube is attached to which post of the power source or battery.

7. Add a few drops of indicator solution to

the beaker. Watch for any colour changes that occur near the mouths or inside the two test tubes. Record any changes you observe, making note of which test tube produced the change.

8. You will test the gas collected in each of

the test tubes, one at a time. When testing the gas inside a test tube, carefully remove the tubes from the water, keeping the test tubes inverted so the gas doesnt escape. Use a test tube holder or

clamp to hold the test tubes while performing the tests.

Place the candle on a glass square or other

suitable support and light it. In the test tube that was attached to the

negative post of the power source, test for the presence of hydrogen gas in the test tube: a. Remove the test tube from the water

and keep it inverted. Allow the water to drain out.

b. Holding the test tube with a test tube holder, bring the open end of the tube over the lit candle.

c. A pale blue flame or a soft pop sound indicates hydrogen is present.

Next use the clamp to hold the test tube that was attached to the positive post of the power source, and test for the presence of oxygen gas: a. With the test tube still inverted, bring

its open end over the glowing (blown out) candle.

b. If the candle again bursts into flame, the gas is oxygen.

Alternate test for oxygen: a. Light a wood splint, then blow it out. b. Place the smoldering splint inside the

test tube. If the splint glows then oxygen is present.

Option 2 1. Fill one of the beakers approximately

three quarters full of the salt solution. The presence of the salt, an electrolyte, will help conduct the electric current.


This will be Beaker A and will be the cathode (site of reduction).

Add a few drops of phenolphthalein to

the water. If the solution appears red after adding the indicator (because the water is too basic), add a few drops of vinegar until the red colour disappears.

2. Fill the second beaker approximately

three quarters full of the salt solution. This will be Beaker B and will act as the anode (site of oxidation).

3. Place the beakers next to one another.

Connect the wires to the battery as follows:

Beaker A (salt solution + phenolphthalein) connect to the negative post of the battery or power source.

Beaker B (salt solution only) connect to

the positive post of the battery or power source.

4. Fold the coffee filter and place one end in

Beaker A and the other in Beaker B, forming a bridge between the two solutions.

The electrolysis will not begin until the

filter paper becomes fully wet. RESULTS Option 1: Brownlee Electrolysis

Apparatus

1. Describe what you observe occurring at the two electrodes. Make special note of the relative amounts of gas that form in each of the two test tubes are they equal amounts? If not, indicate which test tube is attached to which post of the power source or battery.

2. Describe any colour changes that

occurred after the addition of phenolphthalein.

Option 2 1. Record your observations for the two

beakers. Carefully observe the ends of the wire. Watch for any colour changes that occur

CONCLUSIONS AND QUESTIONS Option 1: Brownlee Electrolysis

Apparatus 1. Why does more gas form in one test tube

than in the other? Explain in terms of the half-reactions that occur in each test tube, identifying each test tube by which post they are attached to at the power source.

Also identify each test tube half-reaction

as either the anode or the cathode. Finally explain the flow of electrons

through the system. 2. Explain the colour changes that occurred

after the addition of phenolphthalein.


Option 2 1. What gas was formed in Beaker A? Write

the half-reaction that occurred in this beaker, identify it as either oxidation or reduction, and label it as the anode or cathode.

2. What gas was formed in Beaker B? Write

the half-reaction that occurred in this beaker, identify it as either oxidation or reduction, and label it as the anode or cathode.

3. Explain the flow of electrons through the

system. 4. What was the purpose of adding

phenolphthalein to the solution in Beaker A?


ELECTROCHEMISTRY 3

ELECTROPLATING

OVERVIEW

Electroplating is an economically important process, often used to reduce corrosion or improve the appearance of objects. During electroplating a thin layer of a desirable metal is deposited onto another object.

During electroplating, the object to be plated I attached to the negative post of a power source, causing the object to gain a negative charge. This will attract positive metallic cations from the electrolytic solution, or bath, the object is placed in. In our experiment, positive Cu2+ ions from the bath will become attracted to a nail carrying the negative charge. When the Cu2+ reaches the nail they will gain electrons and become reduced to form solid copper:

Cu2+ (aq) + 2 e- Cu(s)

The copper (II) ions removed from the bath must b replenished; this is accomplished at the anode where a solid copper plate undergoes oxidation:

Cu(s) Cu2+(aq) + 2e-

PURPOSE To use electroplating to plate copper

onto a metal object such as a nail.

SAFETY There are no safety concerns for this lab.


Cathode the metal object to be plated; an iron nail works well. Or try a brass key or a quarter Anode a copper strip Electrolytic solution 1.0 M CuSO4 Battery or power source Beaker or glass jar Insulated wire leads with alligator clips at both ends Un-insulated copper wire Popsicle sticks or other support that will cross the top of the beaker or jar used to suspend the item to be plated (optional) PROCEDURE 1. The object to be plated must be clean for

good results. Prepare by polishing with some steel wool.

2. Use the un-insulated copper wire to

suspend the item to be plated (such as the nail) into the empty beaker. Attach one end of a wire lead to the copper wire supporting the nail and the other end to the NEGATIVE post of the battery or power source.


3. Place the copper strip, the anode, into the empty beaker. Attach one end of a wire lead to the copper strip and attach the other end to the POSITIVE post of the battery or power source.

4. Carefully pour the CuSO4 solution into

the beaker until it is about two-thirds full. If the entire nail is to be plated it must be fully submerged.

5. Allow the reaction to continue for a half-

hour or so. Record your observations while electroplating is continuing.

RESULTS Record your observations during and

after the electroplating procedure. CONCLUSIONS AND QUESTIONS 1. Write the half-reaction that occurs at the

anode of the electrolytic cell. Identify the reaction as either oxidation or reduction.

2. Write the half-reaction that occurs at the

cathode of the electrolytic cell. Identify the reaction as either oxidation or reduction.

3. Write a descriptive paragraph or two that

explains both the flow of copper ions and electrons through the system.


THERMODYNAMICS

HEATING AND COOLING CURVES Overview In this lab a sample of a pure, solid substance is slowly heated in a warm water bath until the sample fully melts. The temperature of the sample is measured every 30 sec during this process and the data is recorded. Next, the liquefied sample is allowed to cool in a cool water bath until the sample solidifies (freezes). Again the temperature of the sample is measured every 30 sec and the data is recorded. Data for the two phase changes is plotted on a single graph and the graph is used to determine the melting point and freezing points of the substance. Purpose

To graph the data for both the melting and freezing of a pure substance

To determine the melting point (m.p.) and freezing point (f.p.) of the substance

To consider the energy changes that occurs during a change in phase.

Equipment, Materials, and Procedure Your teacher will provide you with a list of the materials and equipment required for this lab, and the procedures to follow.

Suitable substances to use include lauric acid, acetamide, or p-dichlorobenzene (mothballs) If you are unable to actually perform the experiment, use the following set of experimental data to plot the graphs and answer the questions at the end of the lab.


Sample Data for the Heating and Cooling of Lauric Acid, C12H22O2

Heating Cooling

Time (min

)

Temperature

(0C)

Time (min

)

Temperature

(0C) 0 30.0 0 55.0

0.5 33.2 0.5 52.6 1.0 35.5 1.0 49.5 1.5 37.5 1.5 45.5 2.0 39.0 2.0 44.3 2.5 41.0 2.5 44.1 3.0 42.0 3.0 44.0 3.5 42.6 3.5 44.0 4.0 43.0 4.0 44.0 4.5 43.4 4.5 44.0 5.0 43.6 5.0 44.0 5.5 43.7 5.5 44.0 6.0 43.8 6.0 44.0 6.5 44.0 6.5 44.0 7.0 44.0 7.0 44.0 7.5 44.1 7.5 43.7 8.0 44.2 8.0 43.5 8.5 44.5 8.5 43.3 9.0 45.2 9.0 43.0 9.5 46.0 9.5 42.6

10.0 47.5 10.0 42.3 10.5 49.0 10.5 41.9 11.0 51.4 11.0 41.5 Results Prepare a single graph that shows the results of both the heating and cooling data. The graph should be a plot of Temperature (C) versus Time (min). Students are strongly encouraged to use the graphing

capabilities of a spreadsheet such as Excel or Quatro Pro to create the graph. Computer generated graphs are preferable to hand-drawn graphs. The following items apply to creating

graphs:

All graphs require an appropriate, descriptive title. The name of the substance used in the experiment should be included in the title. Please include your name on all graphs.

Both axes must be labeled, including the units of measurement (0C and min)

Graphs should fill the page, except for 1 inch margins on all sides.

Conclusions and Questions 1. Based on your graph, determine the

melting point and freezing point of the substance used.

How do these values compare? 2. Consider the diagonal region of the

heating curve, as the sample is being heated. What does the temperature change indicate about the change in kinetic energy of the particles in the sample?


3. Describe the shape of your graph during the actual changes of state (while the substance is actually melting or solidifying)

4. During the heating process, heat is

continually being supplied to the sample throughout the entire time of the experiment even though the temperature remains constant during the actual change of phase. How can this be explained at the molecular level, in terms of what is happening to the chemical bonds holding the particles together in the solid state?

Hint remember that temperature is a

measure of the average kinetic energy of the particles in a sample of matter. If temperature remains constant, what does this say about the kinetic energy of the particles? If the energy being supplied is not changing the kinetic energy, then what form of energy is changing?

5. Consider the diagonal region of the

cooling curve, as the sample is being cooled. What does the temperature change indicate about the change in kinetic energy of the particles in the sample?

6. During the cooling process, heat is

continually being removed from the sample, yet the temperature remains constant during the actual freezing process. If the constant temperature

indicates no change in kinetic energy, then what form of energy is changing?


HEAT OF SOLUTION OF A SOLID

Overview When a solid dissolves in water to form a solution energy changes occur. In this experiment, you will determine the heats of solution for two substances - ammonium nitrate and sodium acetate - using a simple styrofoam cup as a calorimeter. Purpose

To measure experimentally the amount of heat absorbed or released during the dissolving of ammonium nitrate and of sodium acetate in water.

Materials Styrofoam cup Balance Thermometer 100 mL graduated cylinder Anhydrous sodium acetate, CH3CO2Na Ammonium nitrate, NH4NO3 Procedure 1. Accurately find the mass of about 150 mL

of tap water. Record this value and other measurements in the data table. Add the water to the Styrofoam cup calorimeter.

2. Accurately find the mass of about 15 g of solid ammonium nitrate.

3. Find and record the initial temperature of the water. Record to the nearest 0.2C.

4. Dissolve the solid ammonium nitrate in the water, stirring with the thermometer.

Record the maximum temperature difference from the initial reading.

5. Rinse out the cup, dry it thoroughly, and repeat the experiment using a sample of about 15 g of sodium acetate in place of the ammonium nitrate.

Calculations You will be calculating the molar heat of solution for each of the two solids. The molar heat of solution is the amount of heat associated with the dissolving of one mole of solute (the solute is the substance that gets dissolved). A calculation table (next page) may be used to work through the required calculations.

Summary of calculations required: 1. The total heat of solution can be

calculated by determining the amount of heat absorbed or lost by the water during the experiment. This is calculated using the formula:

Q = m c T heat absorbed/lost (joules)

= mass of water (g)

specific heat of water

(J/g C)

temperaturchange

(C)

The specific heat of water is more-or-less

a constant, with a value of 4.18 J/ g C.


Important Note: The amount of heat absorbed (or lost) by the water is equal to the amount of heat lost (or absorbed) by the solute. Therefore, to determine the value for heat gained/lost by the solute, just reverse the sign for the value of heat gained/lost by the water.

The value for the amount of heat

absorbed or lost by the solute will be converted into kJ.

2. The molar heat of solution, expressed

as kJ/mol, is determined by converting the mass of the solute used to moles:

)/(

)(molgMolarMass

gmassmoles =

Total heat of solution tells us how

much heat was gained or lost in our actual experiment with the mass of the substances actually used, while

Molar heat of solution tells us how much energy would be gained or lost for one mole of the substance.

Molar heat solution =

total heat solution moles solute used

Perform the necessary calculations by completing the calculations table.

Chemistry Book of Experiments

32

Table1. Data recording sheet Ammonium nitrate

NH4NO3

sodium acetate CH3CO2Na

Mass of cup + water (g)

Mass of empty cup (g)

Mass of water (g)

Mass of solid + container (g)

Mass of empty container (g)

Mass of solid used (g)

Initial water temperature (C)

Final water temperature (C)

Change in water temperature

(C)

Sample Data Use the following sample data if you are not able to perform the experiment. Ammonium nitrate

NH4NO3

Sodium acetate CH3CO2Na

Mass of cup + water (g) 100.6 g 103.4 g

Mass of empty cup (g) 1.6 g 1.6 g

Mass of water (g)

Mass of solid + container (g) 53.9 g 38.4 g

Mass of empty container (g) 23.3 g 23.2 g

Mass of solid used (g)

Initial water temperature (C) 20.2C 19.8C

Final water temperature (C) 8.0C 27.6C Change in water temperature

(C)


33

Table 2: Calculating Molar Heat of Solution

ammonium nitrate

NH4NO3 sodium acetate

CH3CO2Na

1. Total mass of water used (from Table 1)

2. Change in water temperature (from Table 1)

3. Specific heat of water 4.18 J/ g C 4.18 J/ g

4. Energy absorbed/lost by the water (J) [see formula 1: Q = mcT]

5. Energy absorbed/lost by the solute (J) [same value, but opposite sign, as in calculation step 4]

6. Energy absorbed/lost (kJ) [convert j into kJ]

7. Molar mass of solute (g/mol)

8. Moles of solute actually used (mol)

9. Molar heat of solution (kJ/mol) [see formula 2]


34

QUESTIONS 1. Both of the substances used in this

experiment consist of ions. Give the chemical formulas of the two ions that make up each of the compounds used in this experiment:

(Example: sodium chloride, NaCl, consists

of: Na+ and Cl- ions) Ammonium nitrate, NH4NO3:

________________ Sodium acetate, CH3CO2Na:

________________ 2. For each of the dissolving processes, state

whether the overall process was endothermic (absorbed energy) or exothermic (released energy)

Ammonium nitrate, NH4NO3: Sodium acetate, CH3CO2Na:

3. Write balanced equations for the

dissociation of each ionic compound. Include both the physical states of all participants (s, l, g, aq) and the energy term. For the energy term - be sure to write it on the appropriate side of the equation, and give the value for this term that you obtained during your experiment.

Example: The equation for the solution process of

sodium chloride, an endothermic solution process, is:

NaCl(s) + energy Na+(aq) + Cl-(aq) You calculated the value for the energy term

(the final row in the calculation table). Include this value in your equation:

Ammonium nitrate, NH4NO3: Sodium acetate, CH3CO2Na:

Chemistry Book of Experiment Progress Education Tanzania 35

4. The percentage error, Er, is a way to measure the accuracy of your experimental work. Percent error is calculated as follows:

%100= Expected

E ExpectedObservedr

Where: observed = value you obtained during your experiment Expected = actual or expected value. Using the data obtained from your experiment and the following actual values for molar

heats of solution, calculate your percentage errors for this experiment.

Observed Value Expected Value Percentage E

ammonium nitrate

+25.8 kJ/mol

sodium acetate

-17.4 kJ/mol


CHEMICAL EQUILIBRIUM

LE CHATELIERS PRINCIPLE

OVERVIEW Some chemical reactions are reversible;

that is, not only do the reactants react to form products, but the products can in turn reform into the original reactants. When a reversible system reaches a point at which the rate of the forward reaction equals the rate of the reverse reaction, the system is said to be at equilibrium. At equilibrium, no observable changes in the system can be noted. It is important to understand that at equilibrium all reaction participants are present all reactant particles, as well as all product particles. This does not mean, however, that all are present in equal amounts.

Le Chataliers Principle tells us that if a

system at equilibrium is subjected to a stress, the system will shift in order to minimize the effect of that stress. In this laboratory exercise you will examine how various stresses cause equilibrium systems to shift. These particular equilibrium systems undergo color changes as the equilibrium shifts between the reactants and products, allowing you to see which side of the reaction becomes favored.

PURPOSE

To observe the effect of various stresses (ion concentration; temperature) on equilibrium systems.

SAFETY 1. Use extreme caution when handling the

acidified cobalt(II) chloride hexahydrate, CoCl2

. 6 H2O. HCl, used to make the acidified solution, is highly corrosive.

Report any spills immediately to the

teacher. Spills on the skin should be flushed with cold water.

2. Potassium chromate, K2CrO4 is a hazardous substance. Use with care.

3. Use caution when using the hot water bath.

THE REACTIONS Part 1: Chromate Dichromate

Equilibrium 2 CrO4

2-(aq)+2 H3O

+(aq) Cr2O7

-(aq)+ 3H2O(l)

yellow orange Part 2: Iron(III) Thiocyanate Ion

Complex Fe3+(aq) + SCN

-(aq) Fe(SCN)

2+(aq)


light brown red Part 3: Cobalt(II) Chloride Complex; Effect of Temperature [CoCl4]

2- + 6 H2O [Co(H2O)6]2+ + 4 Cl-

blue pink EQUIPMENT AND MATERIALS

Amounts needed per group are approximate Part 1: Chromate Dichromate

Equilibrium

0.1 M K2CrO4, 5 mL per group 0.1 M K2Cr2O7, 5 mL per group 1 M HCl, dropper bottle 1 M NaOH, dropper bottle 2 test tubes in a test tube rack

Part 2: Iron(III) Thiocyanate Ion

Complex

0.1 M FeCl3, 10 mL per group 0.1 M KSCN, 10 mL per group 0.1 M KCl, 5 mL per group 10 mL graduated cylinder 250 mL or larger beaker distilled water, approx 100 mL 4 test tubes in a test tube rack dropper pipette

Part 3: Cobalt(II) Chloride Complex; Effect of Temperature

0.2 M acidified COCl2. 6 H2O, 15 mL per group

hot water bath (approx. 90 C) ice water bath 3 test tubes in a test tube rack

PROCEDURE Part 1. Chromate Dichromate

Equilibrium 1. Fill a test tube approximately half-full

with potassium chromate, K2CrO4, (Tube 1).

2. Fill another test tube approximately half-

full with dichromate, K2CrO4 (Tube 2). 3. To Test Tube 1 add several drops of HCl.

HCl is an acid; adding HCl increases the concentration of H3O

+ ions in the equilibrium system. Note the color change.

4. After recording the color change in Test

Tube 1, add several drops of NaOH. NaOH is a base; adding a base decreases the concentration of H3O

+ ions in the equilibrium system. Record the color change.

5. To Test Tube 2 add several drops of

NaOH until a color change is observed. 6. After recording the color change, add

several drops of HCl to Test Tube 2. Again note the change in color.


Part 2: Iron(III) Thiocyanate Ion Comlex 1. Pour 5 mL of 0.1 M FeCl3 into the

beaker. 2. After rinsing the graduated cylinder,

measure 5 mL of 0.1 M KSCN. Add to the beaker containing the FeCl3. Note the color change.

3. Add enough distilled water to the beaker

to dilute the solution to a light brown color. Pour some into a test tube to check the color.

4. Pour about 10 mL of this solution into

each of the four numbered test tubes. The first test tube will serve as a control.


FeCl3 until a color change is observed. Adding more FeCl3 increases the concentration of Fe3+ in solution. Record the color change.


KSCN until a color change is observed. Adding more KSCN increases the concentration of SCN- in solution. Record the color change.

7. To Test Tube 4 add several drops of KCl

until a color change is observed. Adding KCl causes the concentration of Fe3+ to decrease because the Fe3+ and Cl- react to form FeCl4

-.

Part 3: Cobalt (II) Chloride Complex; Effect of Temperature 1. Fill three test tubes approximately half-

full with the acidified CoCl2 . 6 H2O

solution. 2. Test tube 1 will serve as the control.

Keep this test tube at room temperature. Record the initial colour of the solution.

3. Place the second test tubes in the hot

water bath. After a few minutes a colour change will occur. Record the colour.

4. Place the third test tube in the cold water

bath. Record any colour change. 5. Reverse tubes 2 and 3. Observe any

colour changes that occur. RESULTS Part 1: Chromate Dichromate

Equilibrium

Solution

K2CrO4 initial colour

HCl added

NaOH added

K2Cr2O7 initial colour

NaOH added


HCl added

Part 2: Iron (III) Thiocyanate Ion Complex Test Tube

Stress Applied

Initial Color

Final Color

1 Control --

2 Fe3+ added

3 SCN- added

4 Cl- added: decreases [Fe3+]

Part 3: Cobalt (II) Chloride Complex; Effect of Temperature

Temperature Solution Colour

room temperature

hot water bath

cold water bath

CONCLUSIONS AND QUESTIONS Part 1: Chromate Dichromate

Equilibrium

1. Use Le Chateliers Principle to explain

the color changes observed in both test tubes with the addition of both HCl and NaOH.

Write your answer as a clear, descriptive

paragraph.

2 CrO42-

(aq) + 2 H3O+

(aq) Cr2O7-(aq) +

3H2O(l) yellow orange Adding HCl, an acid, increases the

concentration of H3O+ ions; adding

NaOH, a base, decreases the concentration of H3O

+ ions. Part 2: Iron (III) Thiocyanate Ion Complex 2. Use Le Chateliers Principle to explain

the color changes observed in Test Tubes 2 3 upon the addition of FeCl3, KSCN, and KCl.

Write your answer as a clear, descriptive

paragraph.

Fe3+(aq) + SCN-(aq) Fe(SCN)

2+(aq)

light brown red Part 3: Cobalt (II) Chloride Complex; Effect of Temperature [CoCl4]

2- + 6 H2O [Co(H2O)6]2+ + 4 Cl-


blue pink 3. Based on the colour changes observed in

the hot water and cold water baths, determine whether the forward reaction is endothermic or exothermic.

Rewrite the equation with a simple

energy term ( + heat) included on the appropriate side of the equation. You may find it easier if you begin by using only the terms heat, pink, and blue in your equation.


SOLUTIONS

PRECIPITATION REACTIONS

OVERVIEW When two aqueous solutions of ionic compounds are combined, a solid precipitate may form. This occurs when a positive cation from one solution and a negative anion from the other solution form an insoluble compound. The attraction between the oppositely charged ions is stronger than the attraction of the individual ions to the polar water molecules, the solutions solvent. The result is a solid precipitate that rapidly comes out of solution.

For example, when solutions of silver nitrate, AgNO3, and sodium chloride, NaCl are combined, a double displacement reaction occurs and a white precipitate, AgCl, immediately forms:

AgNO3 (aq) + NaCl(aq) AgCl(s) + NaNO3(aq)

The net ionic equation, which removes the un-reacting spectator ions, shows more clearly the ions of interest:

Ag+(aq) + Cl-(aq) AgCl(s)

If no insoluble combination between anions and cations exist, no precipitate will form. Instead, all ions remain in solution and no reaction occurs.

In this lab you will use your knowledge of precipitation tables to predict precipitation reactions. Examine the lists of solutions you will be using for this experiment. You will be mixing solutions from Set A with Set B. Which combinations do you predict will result in a precipitate? Record your predictions. You will then test your predictions by combining pairs of solutions to see if a precipitate forms. Additionally you may be asked to prepare the standard 0.10M solutions used for this lab. Your teacher may have you do this in advance of the precipitation tests. PURPOSE

To predict precipitation reactions. To observe a variety of precipitation

reactions To write net ionic equations for

precipitation reactions To prepare 0.10M standard solutions

SAFETY

Follow general lab safety rules for this experiment.



Clear acetate overhead sheet or spot plate

Grease pencil to draw a grid on acetate sheet

Solutions as assigned by your teacher. You may have any or all of the following 0.10 M solutions:

Set A

Set B

sodium nitrate, NaNO3

potassium nitrate, KNO3

silver nitrate, AgNO3

ammonium nitrate,

NH4NO3

lead(II) nitrate,

Pb(NO3)2

calcium nitrate,

Ca(NO3)2

magnesium nitrate,

Mg(NO3)2

barium nitrate, Ba(NO3)2

copper(II) nitrate,

Cu(NO3)2

iron(III) nitrate,

Fe(NO3)3

sodium chloride, NaCl

sodium hydroxide,

NaOH

sodium bromide,

NaBr

sodium sulfide, Na2S

sodium iodide, NaI

sodium phosphate,

Na3PO4

sodium sulfate,

Na2SO4

sodium carbonate,

Na2CO3

Alternates

Alternates

sodium acetate, NaC2H3O2

zinc acetate,

Zn(C2H3O2)2

potassium carbonate,

K2CO3

potassium phosphate,

K3PO4

ammonium sulfate, (NH4)2SO4

Optional equipment if preparing standard solutions

centigram or electronic balance scoopula 100 mL volumetric flask or graduated

cylinder 250 mL or smaller beaker dropper bottle to store your solution

PROCEDURE 1. If your teacher has assigned you certain

standard solutions to prepare, complete Step 1. If these solutions have already been prepared, continue to Step 2.

To prepare a standard solution you will

need to measure a precise mass of the solid and add to it just enough water to make 100 mL (0.100 L) of solution.

Calculate the mass of solid needed. Using

unit analysis we can easily do this if we first determine the molar mass of the compound (units for molar mass are gmole1):

g = 110.010.0 L

Lmole

moleg

mass

= molar

mas

desired concentra

tion

desired

volu


s me For example, to prepare 100 mL of a

0.10M solution of NaCl we first calculate the molar mass of NaCl and find it to be

58.5 gmole-1. Since we want to prepare 0.100 L of a

0.10 M solution we can determine the mass of NaCl needed:

g = gLLmole

moleg 59.0

110.010.05.58

=

Check your calculations with your

teacher before proceeding. Prepare the standard solution. To prepare

the solution, accurately measure out the required mass and place it in a beaker. Add enough distilled water (about 20 mL or less) to dissolve the salt.

Pour this solution into a 100 mL

volumetric flask (if available) or a 100 mL graduated cylinder. Add enough distilled water to bring the total volume of solution up to the 100 mL mark on the flask or cylinder.

Store your solution in the dropper bottle

or other container provided by your teacher and label it with the formula of the compound and the molarity of the solution (e.g. 0.10 M NaCl)

2. Your teacher may ask to see your

predictions for the precipitation reactions. You may record these on the chart provided.

3. On the blank Test Grid provided, list

along the top row the Set A solutions you will be testing by giving the formula and charge of the ions present in the solution. An example is shown on the grid. Similarly, list the Set B solutions in the first column.

Prepare a second, identical test grid

which you will use to record your results. Place one of the test grids underneath the

overhead sheet. You may find it useful to trace the grid on the overhead using the grease pencil. This will help prevent drops from one test reaction from running into another test reaction.

If using a spot plate instead of an

overhead, still prepare the Test Grid, and keep it beside your spot plate for reference.

3. On the overhead or the spot plate carry

out your test reactions. Place one drop of the solution containing the Set A solution in the first cell; add a drop of the solution containing the Set B solution to that same cell. Be careful not to let the dropper bottle touch the drop already in the cell in order to avoid contaminating the solutions!

Observe if a reaction occurs. If a

precipitate forms, record this on your test grid data sheet as PPT, also noting the colour of the precipitate. If no


precipitate forms, write NR (for No Reaction) on your data sheet.

Continue testing all possible

combinations of Set A and Set B solutions.

QUESTIONS AND CONCLUSIONS 1. For every reaction in which a precipitate

occurred, write both the full reaction equation and also the net ionic equation.

In both equations be sure to identify the

precipitate as a solid, by (s) after the formula.

Be sure to balance all equations. 2. Comment on the accuracy of your

predictions.


RESULTS: PREDICTIONS The following table lists the solutions you may be mixing together. Do you predict that a precipitate will form? If so, what would be the formula of the precipitate?

If no precipitate is expected, write NR for no reaction. In the lab you will test your predictions.

PREDICTIONS (continued)

Reactants Predicted

Precipitate Predicted

Precipitate

NaCl + NaNO3 NaBr + NaNO3

NaCl + KNO3 NaBr + KNO3

NaCl + AgNO3 NaBr + AgNO3

NaCl + NH4NO3 NaBr + NH4NO3

NaCl + Pb(NO3)2 NaBr + Pb(NO3)2

NaCl + Ca(NO3)2 NaBr + Ca(NO3)2

NaCl + Mg(NO3)2 NaBr + Mg(NO3)2

NaCl + Ba(NO3)2 NaBr + Ba(NO3)2

NaCl + Cu(NO3)2 NaBr + Cu(NO3)2

NaCl + Fe(NO3)3 NaBr + Fe(NO3)3

NaOH + NaNO3 Na2S + NaNO3

NaOH + KNO3 Na2S + KNO3

NaOH + AgNO3 Na2S + AgNO3

NaOH + NH4NO3 Na2S + NH4NO3

NaOH + Pb(NO3)2 Na2S + Pb(NO3)2

NaOH + Ca(NO3)2 Na2S + Ca(NO3)2

NaOH + Mg(NO3)2 Na2S + Mg(NO3)2

NaOH + Ba(NO3)2 Na2S + Ba(NO3)2

NaOH + Cu(NO3)2 Na2S + Cu(NO3)2

NaOH + Fe(NO3)3 Na2S + Fe(NO3)3


Reactants Predicted

Precipitate Predicted

Precipitate

NaI + NaNO3 Na2SO4 + NaNO3

NaI + KNO3 Na2SO4 + KNO3

NaI + AgNO3 Na2SO4 + AgNO3

NaI + NH4NO3 Na2SO4 + NH4NO3

NaI + Pb(NO3)2 Na2SO4 + Pb(NO3)2

NaI + Ca(NO3)2 Na2SO4 + Ca(NO3)2

NaI + Mg(NO3)2 Na2SO4 +

Mg(NO3)2

NaI + Ba(NO3)2 Na2SO4 + Ba(NO3)2

NaI + Cu(NO3)2 Na2SO4 + Cu(NO3)2

NaI + Fe(NO3)3 Na2SO4 + Fe(NO3)3

Na3PO4 + NaNO3 Na2CO3 + NaNO3

Na3PO4 + KNO3 Na2CO3 + KNO3

Na3PO4 + AgNO3 Na2CO3 + AgNO3

Na3PO4 + NH4NO3 Na2CO3 + NH4NO3

Na3PO4 + Pb(NO3)2 Na2CO3 + Pb(NO3)2

Na3PO4 + Ca(NO3)2 Na2CO3 +

Ca(NO3)2

Na3PO4 + Mg(NO3)2 Na2CO3 +

Mg(NO3)2

Na3PO4 + Ba(NO3)2 Na2CO3 + Ba(NO3)2

Na3PO4 + Cu(NO3)2 Na2CO3 +

Cu(NO3)2

Na3PO4 + Fe(NO3)3 Na2CO3 + Fe(NO3)3


SAMPLE DATA

SET A SOLUTIONS

SET

B S

OLU

TIO

NS

Na+

NO3

-

K+

NO3

-

Ag+

NO3

-

Na+

Cl-

Na+

OH-


SAMPLE DATA

SET A SOLUTIONS

SET B SO

LUT

ION

S

Na+

NO3

-

K+

NO3

-

Ag+

NO3

-

NR NR white PPT

Na+

Cl-

Na+

OH-


CHEMICAL KINETICS 1

FACTORS AFFECTING REACTION RATE OVERVIEW Chemical reactions occur at different rates. In this experiment you will consider some of the key factors that influence the rate of a reaction: Nature of reactants - particle size Temperature Concentration Catalysts

According to the collision theory, the rate of a reaction depends on the frequency of collisions between reacting particles. The more frequent the collisions, the faster the rate of the reaction. However, in order for the collisions to be effective, the particles must collide with sufficient energy (activation energy). Furthermore, the particles must collide with the proper orientation. The factors that will be examined in this lab influence reaction rate by either increasing how often collisions occur or by making collisions more effective. PURPOSE To examine factors that increase

reaction rate

SAFETY Portions of this lab may only be carried

out under the supervision of a teacher Safety goggles must be worn when

working with acids. EQUIPMENT AND MATERIALS Part 1: Effect of Particle Size solid zinc, approx. 0.5 cm 2 cm, or

solid marble chips zinc powder or calcium carbonate

powder balance 2 test tubes 1M HCl (approx 10 mL per group)

Part 2: Effect of Temperature

3 Alka Seltzer tablets 3 250-mL beakers water at three temperatures

with ice, room temperature, warm (around 70C)

Part 3: Effect of Concentration

1M HCl, 5 mL per group 3M HCl, 5 mL per group 6M HCl, 5 mL per group


3 pieces of zinc metal, each approx 1 cm 1 cm

3 test tubes

Part 4: Effect of a Catalyst

3% hydrogen peroxide, H2O2 10 mL per group

0.1 M iron(III) nitrate, Fe(NO3)3 0.1 M sodium chloride, NaCl 0.1 M calcium chloride, CaCl2 0.1 M potassium nitrate, KNO3 0.1 M manganese chloride,

MnCl2 100-mL graduated cylinder 10-mL graduated cylinder 7 test tubes per group

PROCEDURES Part 1: Effect of Particle Size on

Reaction Rate Powdered calcium carbonate and marble chips may be used instead of zinc

Zn + 2 HCl ZnCl2 + H2

CaCO3 + HCl CaCl2 + H2CO3 1. Obtain a piece of solid zinc metal,

approximately 0.5 cm 2 cm. Find the mass of this sample, and place it in a test tube.

2. Using the balance obtain a sample of powered zinc that is close to the mass of your piece of solid zinc. Place this sample in the second test tube. Caution: powdered zinc is flammable.

3. Place both test tubes in a test tube rack.

Add 5 mL of 1M HCl to both test tubes. Be sure to wear your safety goggles.

4. Observe both test tubes and record your

observations and record your observations in the data table.

Part 2: Effect of Temperature 1. Half fill three 250-mL beaker with

water. In one beaker add several ice cubes. A second beaker will contain water at room temperature. In the third beaker add water that has been heated to about 70C.

2. Record the water temperature in the

three beakers, and then add an Alka Seltzer tablet to each.

3. Record the time it takes for the Alka

Seltzer tablet to completely dissolve.


Part 3: Effect of Concentration

Zn + 2 HCl ZnCl2 + H2 1. Pour 5 mL of each of the three HCl

solutions into separate test tubes. Place the test tubes in a test tube rack.

2. Add one piece of zinc to each test tube. 3. Record the time you added the zinc to

the tubes, and the time each reaction stops. Also record your observations for each tube.

Part 4: Effect of a Catalyst In this part of the lab you will determine which substance/substances act as a catalyst for the decomposition of hydrogen peroxide.

2 H2O2 2 H2O + O2 1. Dilute the hydrogen peroxide by adding

10 mL of 3% H2O2 to a 100-mL graduated cylinder. Add 90 mL of distilled water to obtain 100 mL of diluted (0.3%) hydrogen peroxide.

2. Use a small amount of this solution to

rinse out a 10-mL graduated cylinder and 7 test tubes. Pour the rinses away.

3. Place 5-mL of the 0.3% H2O2 solution

into each of the 7 test tubes.

4. Add 5 drops of each of the following

solutions to separate test tubes:

0.1 M FeCl3 0.1 M NaCl

0.1 M Fe(NO3)3 0.1 M CaCl2

0.1 M KNO3 0.1 M MnCl2

5. Mix each tube by swirling the test tube or gently stirring with a clean stirring rod.

6. Observe each solution, noting the

production of any gas bubbles that form. Record each reaction rate as fast, slow, very slow, or none in your data table.

RESULTS Record your results for each part of the lab in the data tables provided on the following page. CONCLUSIONS For each part of this lab present your results and conclusions in formal lab report format. A paragraph will be sufficient for each of the four parts to this lab. In clearly worded and complete sentences describe what reactions were carried out and present your observations findings. When possible include the chemical equation involved. Refer the reader to the appropriate data table (for example See Table 1.). Use the collision theory to explain your findings.


RESULTS Table 1: Effect of Particle Size on Reaction Rate

Substance Tested Observations

powdered zinc or calcium carbonate

solid zinc or marble chips

Table 2: Effect of Temperature

Water Condition

Water Temperature (C)

Time to Completion

cold

room temperature

warm

Table 3: Effect of Concentration

Acid Concentration

Start Time Time at

Completion Observations

1 M HCl

3 M HCl

6 M HCl

Table 4: Effect of a Catalyst


Possible Catalysts

FeCl3 NaCl Fe(NO3)3 CaCl2 KNO3 MnCl2

Reaction Rate


SOLUTIONS

THE SOLUBILITY OF A SALT

OVERVIEW A key factor affecting the solubility of a substance how much solute can be dissolved in a solvent is temperature. For most substances increasing temperature will increase solubility - more solute will be able to dissolve in the same volume of solvent. A solubility curve illustrates how the solubility of a substance varies with temperature. By determining the mass of solute that can be dissolved in a volume of solvent under a variety of temperatures we can easily construct a solubility curve. In this lab exercise you will create a solubility curve for an ionic compound, potassium nitrate, KNO3. PURPOSE

To calculate the solubility of a substance under a variety of temperatures.

To construct a solubility curve based on experimental data.

SAFETY Use caution when using the hot water

bath to avoid hot water and steam burns.


Balance Four test tubes per group Hot water baths 400 ml or 600 ml

beaker, ring stand, test tube clamp, gas burner or alternative

Four thermometers per group 10 ml graduated cylinder Scoopula Grease pencil Test tube clamp Test tube rack Solid potassium nitrate distilled water

PROCEDURE 1. Prepare a water bath by filling a large

beaker approximately 2/3 full with water. Place the beaker on a ring stand above a gas burner and begin heating the water to just below boiling. While this is heating continue with Step 2.

2. Using a grease pencil number your test

tubes 1 through 4. 3. Accurately measure out the following

masses of solid potassium nitrate, placing the salt in the appropriate test tube. It is not necessary that you measure out exactly


the masses given below, but you must record the precise masses you actually use.

Test tube 1: 2 grams KNO3

Test tube 2: 4 g

Test tube 3: 6 g

Test tube 4: 8 g 4. Add exactly 5.0 mL of water to each of

the test tubes. 5. Place each of the tubes into the water

bath in order to dissolve the solid KNO3 in each test tube. You may find it necessary to use a stirring rod to help the dissolving process, particularly for est tubes 3 and 4.

6. Remove test tube 1 from the hot water

bath once the KNO3 has fully dissolved and place a thermometer in the tube. Watch the solution carefully. Record the temperature as soon as you see crystals forming within the test tube (you will need to wait awhile for crystals to form in this first test tube).

7. Repeat Step 6 for the other test tubes

once the KNO3 dissolves. Be prepared to act quickly for test tube 4

crystallization may occur very soon after you remove the test tube from the hot water bath. It may be necessary to return the test tube to the water bath to re-dissolve the salt and allowing it to recrystallize again.

RESULTS 1. Copy the following table into your data

book and fill it in with your own set of mass and volume data.

Convert the mass/volume ratio you used for each test tube into mass/100 mL ratio.

For example in test tube 1 you had 2.0 g KNO3 dissolved in 5.0 mL. This is equivalent to what mass of KNO3 per 100 mL of water?

2. Also record the temperature at which

recrystallization occurred. This is the solubility of the substance at that temperature - the maximum amount of solute that can be dissolved in 100.0 mL of water at that temperature.


Test Tub

e

Mass KNO

3 (g) used

Volume

Water (mL)

Convert to

g/100 mL

Saturation

Temp

(C) 1 2 3 4

3. Using graph paper construct a solubility

curve for KNO3 based on your data. CONCLUSIONS AND QUESTIONS 1. The line on your graph represents the

concentration of a saturated solution of KNO3 for various temperatures. Be sure to add values on both the X- and Y-axes.

Connect your data points by a smooth

curve. You may add this additional data to your

curve:

Solubility at 0C: 13 g/100 mL Solubility at 100C: 247 g/100 mL 2. Based on your solubility curve, predict

the solubility of KNO3 at the following temperatures:

a) 50C

b) 70C 3. Based on your solubility curve, would

you best describe the following solutions as unsaturated, saturated, or supersaturated?

a) 70 g / 100 mL H2O at 45C

b) 80 g / 100 mL H2O at 70C

c) 80 g / 100 mL H2O at 30C

The Solubility of Potassium Nitrate


0 10 20 30 40 50 60 70 80 90

Temperature (C)

Sol

ubili

ty g

KN

O3 /

100

mL

H2O

0

60

80

100

40

160

140

120

20

180


CHEMICAL KINETICS 2

RATE OF COOLING OF WATER Overview In this lab a sample of hot water is allowed to cool, with the temperature of the water being recorded every 30 s until the water reaches room temperature. A graph of the data is prepared. The rate of cooling is calculated for various time periods over the full course of the experiment; during the first ten minutes; and during the last ten minutes. The calculated rates of cooling are then compared. Purpose To graph the data for the cooling of

water. To calculate rate of cooling To compare rates of cooling for

different time periods. Equipment and Materials Thermometer Heat source Bunsen burner or hot

plate Large beaker or other container for the

water Approximately 500 ml of water Procedure

1. Heat the water until it begins to boil, and then carefully remove it from the heat source.

2. Record the water temperature every 30 s as it cools. Continue recording until the water temperature reaches room temperature or slightly above room temperature.

Sample Data Table for the Cooling of Water

Cooling Data for Water Time (min) Temperature (C) 0 0.5 1.0 1.5 2.0 2.5 3.0 3.5 4.0 4.5 5.0 5.5 6.0 6.5 7.0 7.5 8.0


8.5 9.0 9.5 . . continue Results

Prepare a graph that shows the results of the data. The graph should be a plot of Temperature (C) versus Time (min). Students are strongly encouraged to use the graphing capabilities of a spreadsheet such as Excel or Quatro Pro to create the graph. Computer generated graphs are preferable to hand-drawn graphs.

The following items apply to creating graphs: All graphs require an appropriate,

descriptive title. The name of the substance used in the experiment should be included in the title. Please include your name on all graphs.

Both axes must be labeled, including the units of measurement (C and min)

Graphs should fill the page, except for 1 inch margins on all sides.

Conclusions and Questions 1. You will calculate the rate of cooling for

three time periods: The entire cooling period

The first ten minutes The last ten minutes

To calculate a rate, determine how much the temperature changed within the desired time period, and divide by the time. For example, if the water temperature dropped 15 degrees during the first five minutes, the rate would be calculated as:

RATE = Cmin/3min5min15 0=

2. How do the three rates compare? Why?

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