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Equilibrium Labs Name: _____________________________________ of 19

Equilibrium and LeChatelier’s PrinciplePreliminary Lab Assignment (Answer the questions on this lab and email the file to me.

Send page 1 and your discussion of each part. You and partner collaborate on each part. Partner 1 do part 2&4 and Partner 2 do part 1&2.)

1. Define equilibrium.

2. State LeChatelier’s Principle.

3. 6 CO2(g) + 6 H2O(1) C6H12O6(s) + 6 O2(g) ΔH = 2820 kJ

For the above reaction tell how the amount of C6H12O6(s) present at equilibrium would be affected by each of the following: Partner 1 do aceg and partner 2 do bdfh

(a) Some CO2(g) is added.

(b) The temperature is raised.

(c) The volume is decreased.

(d) Some O2(g) is removed.

(e) Some of the C6H12O6(s) is removed.

(f) A catalyst is added.

(g) Some H2O is removed.

(h) Ice is added.


Equilibrium and LeChatelier’s PrincipleLeChatelier’s Principle states that: If an equilibrium system is subjected to a stress, the system will react to remove the stress. To remove a stress, a system can only do one of two things: form more products using up reactants, or reverse the reaction and form more reactants, using up products. In this experiment you will form several equilibrium systems. Then, by putting different stresses on the systems, you will observe how equilibrium systems react to a stress.

Before you carry out each section, predict which way you think the equilibrium will shift. Then, carry out the reaction to verify your prediction.

ChemicalsSodium chloride, NaCl(s) Hydrochloric acid, HCl, 12 MPotassium thiocyanate, KSCN, 0.002 M Hydrochloric acid, HCl, 0.1 MBromthymol blue indicator solution Sodium hydroxide, NaOH, 0.1 MPotassium thiocyanate, KSCN(s) Iron(III) nitrate, Fe(NO3)3, 0.2 MSilver nitrate, AgNO3, 0.1 M Ethanol, C2H5OH(l)Cobalt(II) chloride, CoC12• 6H2O(s) Disodium hydrogen phosphate, Na2HPO4(s)EquipmentTest tubes, 13- x 150-mm Test tube rackBeaker, 100 mL Graduated cylindersStirring rod Funnel, filter paper, and holder for funnel

Procedure

Safety AlertYou will be using a concentrated solution of hydrochloric acid, as well as 0.1 M hydrochloric acid. Hydrochloric acid is hazardous. Use it with care. It has strong vapors. Avoid breathing them. Wash spills off yourself with lots of water. Neutralize spills on the lab bench with baking soda.

You will also be using 0.1 M sodium hydroxide. This is also hazardous. Wash spills off with water; neutralize spills on the lab bench with vinegar.

The alcohol is flammable. Keep it away from flames.

Wear Chemical Splash Goggles and a Chemical-Resistant Apron.


1. Equilibrium in a Saturated Solution

You will investigate the equilibrium in saturated sodium chloride solution:

NaCl(s) Na+(aq) + Cl-

(aq)

Pour some solid NaCl into a 13- x 100-mm test tube and fill the tube 3/4 full of distilled water. Cork and shake to form a saturated solution. If all the NaCl dissolves, pour some additional NaCl in the tube and shake until a saturated solution with some excess solid is obtained.

Filter the solution into a second test tube. To this saturated solution of NaCl, add some Cl- ions in the form of concentrated HCl. Record and explain the results.

2. An Acid—Base Indicator Equilibrium

Acid—base indicators are large organic molecules that can gain and lose hydrogen ions to form substances that have different colors. The reaction of the indicator bromthymol blue can be illustrated as follows:

Hln(aq) H+(aq) + 1n-

(aq)

yellow blue

In this reaction Hln is the neutral indicator molecule, and ln- is the indicator ion after the molecule has lost a hydrogen ion. Equilibrium reactions can easily be forced to go in either direction. Reactions like this are said to be reversible.

Fill a small test tube about half-full of distilled water. Add several drops of bromthymol blue indicator solution. Add 5 drops of 0.1 M HCl and stir. This will increase the amount of H+ in solution. Note the color of the indicator.

Next add 0.1 M NaOH drop by drop with stirring until no further color change occurs. Adding OH- ions causes the H ion concentration to decrease as the ions combine to form water molecules. Again, note the color. See if you can add the right amount of acid to this test tube to cause the solution to be green in color after it is stirred (half of the indicator is blue and half is yellow).

3. A Complex Ion Equilibrium

An equilibrium system can be formed in solution with the following ions:

Fe+3(aq) + SCN-

(aq) FeSCN+2(aq)

colorless colorless red-brown

The iron ion (Fe+3) and the thiocyanate ion (SCN-) are both colorless; however, the ion that forms from their combination, the FeSCN+2 ion, is colored a dark red-brown. It is the color of this ion that will indicate how the equilibrium system is being affected.

Pour about 25 mL of 0.0020 M KSCN solution (a source of SCN- ion) into a beaker. Add 25 mL of distilled water and 5 drops of 0.20 M Fe(NO3)3 solution. Swirl the solution and note the following: the color of the KSCN solution, the color of the Fe(NO3)3 solution, and the color of the resulting complex ion.

You will stress the equilibrium system that has resulted in several ways. Pour equal amounts of the solution from the beaker into four test tubes. The solution in the first test tube will be the reference solution.

To the second test tube add 2-3 crystals of solid KSCN. Describe the results.

To the third test tube add 6 drops of Fe(NO3)3 solution. Stir and describe the results.

To the fourth test tube add small crystals of Na2HPO4, a few at a time. Stir and note the results. Phosphate ions, P043-, have the ability to form complex ions with Fe3t which has the same effect as removing Fe3

Equilibrium Labs Name: _____________________________________ of 19from solution.

4. An Equilibrium with Cobalt Complex Ions

In this section we will investigate the equilibrium between two different complex ions of cobalt. The reaction is endothermic:

Co(H2O)6+2

(aq) + 4 Cl-(aq) CoCl4

-2(aq) + 6 H2O(l) ΔH = +50 kJ/mol

pink blue

Measure about 10 mL of ethanol into a beaker.

Examine solid cobalt(II) chloride, noting both its color and the formula of the compound. Dissolve small amount of cobalt(II) chloride (about half the size of a pea) in the beaker of ethanol. The solution should be purple. If it is pink, add a little concentrated HCl until it is purple.

Put about 2 mL of the alcoholic cobalt solution into each of three small test tubes. To one of the test tubes, add 3 drops of distilled water, one drop at a time with stirring, noting what happens with each drop. Add 3 drops of distilled water to each of the other two test tubes. Make a note of the effect of this stress on the system.

The first test tube is the control. To the second test tube, add 5 drops of concentrated HCl, 12 M, one drop at a time with stirring. Note the results.

To the third test tube add a few crystals of solid sodium chloride. Stir and note the results.

Put the remainder of alcoholic cobalt solution from the beaker into a fourth test tube. Add 10 drops of 0.1 M silver nitrate solution, one drop at a time. Silver and chloride ions combine to form a precipitate of AgCl. Note the color of the solution as the chloride ions precipitate. You may wish to let the precipitate settle to observe the solution color more easily, or you may centrifuge the test tube.

Obtain a sealed Beral pipet containing some of the alcoholic cobalt chloride—water system. Note its color. Immerse the large end of the pipet in some hot water (about 60°C) and see if there is a color change.

Lastly, chill the Beral pipet in an ice bath to see if the color change in the previous step is reversible.

Explain the effect of the temperature change on the equilibrium in terms of the fact that the value of

ΔH for the reaction is +50 kJ/mol.

DisposalSolutions from Parts 1, 2 and 3 can be safely washed down the sink with excess water. Dispose of the solutions containing cobalt from Part 4 according to the Flinn Chemical Catalog/Reference Manual, suggested disposal method #27f. See the appendix.

DiscussionIn your laboratory manual describe the results of each part of the experiment and interpret them using LeChatelier’s Principle.

Safety AlertEthanol is flammable. Turn off all flames.Silver nitrate causes stains on skin and clothing. Wash spills off with soap and water immediately.


Partner 1: Determining the Ksp

of Calcium HydroxideCalcium hydroxide is an ionic solid that is sparingly soluble in water. A saturated, aqueous, solution of Ca(OH)2 is represented in equation form as shown below.

Ca(OH)2 (s) ↔ Ca2+ (aq) + 2OH– (aq)

The solubility product expression describes, in mathematical terms, the equilibrium that is established between the solid substance and its dissolved ions in an aqueous system. The equilibrium expression for calcium hydroxide is shown below.

Ksp = [Ca2+][OH–]2

The constant that illustrates a substance’s solubility in water is called the Ksp. All compounds, even the highly soluble sodium chloride, have a Ksp. However, the Ksp of a compound is commonly considered only in cases where the compound is very slightly soluble and the amount of dissolved ions is not simple to measure.

Your primary objective in this experiment is to test a saturated solution of calcium hydroxide and use your observations and measurements to calculate the Ksp of the compound. You will do this by titrating the prepared Ca(OH)2 solution with a standard hydrochloric acid solution. By determining the molar concentration of dissolved hydroxide ions in the saturated Ca(OH)2 solution, you will have the necessary information to calculate the Ksp.

OBJECTIVESIn this experiment, you will

Titrate a saturated Ca(OH)2 solution with a standard HCl solution. Determine the [OH–] for the saturated Ca(OH)2 solution. Calculate the Ksp of Ca(OH)2.

Figure 1

Equilibrium Labs Name: _____________________________________ of 19MATERIALSVernier computer interface saturated calcium hydroxide, Ca(OH)2, solutioncomputer 0.050 M hydrochloric acid, HCl, solutionVernier pH Sensor (optional) indicator solutionring stand and ring buret clamptwo 100 mL beakers filter papertwo 250 mL beakers filter funneltwo 50 mL graduated cylinders magnetic stirrer and stirring bar50 mL buret

PROCEDURE1. Obtain and wear goggles.

2. Obtain about 70 mL of a saturated calcium hydroxide solution. CAUTION: Calcium hydroxide solution is caustic. Avoid spilling it on your skin or clothing.

3. Set up a ring stand, ring, filter funnel, and filter paper. Filter your sample of Ca(OH)2 solution into a clean beaker. Measure out about 15 mL of the filtered solution into a 250 mL beaker. Record the precise volume of Ca(OH)2 solution that you are using.

4. Obtain about 200 mL of 0.050 M HCl solution. CAUTION: Handle the hydrochloric acid with care. It can cause painful burns if it comes in contact with the skin.

5. Connect a pH Sensor to Channel 1 of the Vernier computer interface. Connect the interface to the computer using the proper interface cable.

6. Set up the beaker of Ca(OH)2 solution on a magnetic stirrer. If you are not using a magnetic stirrer, use a glass stirring rod to stir the solution throughout the titration.

7. Use a utility clamp to suspend the pH Sensor on a ring stand as shown in Figure 1. Position the pH Sensor in the Ca(OH)2 solution, and adjust its position so that the sensor is not struck by the magnetic stirring bar.

8. Connect a buret to the ring stand. Rinse and fill the buret with the 0.050 M HCl solution.

9. Start the Logger Pro program on your computer. Open the file “23 Ksp” from the Advanced Chemistry with Vernier folder.

10. Conduct the titration carefully. The guidelines below are general suggestions; use your judgment in conducting the titrations to get the best results.a. Before you have added any of the HCl titrant, click and monitor pH for

5–10 seconds. Once the displayed pH reading has stabilized, click . In the edit box, type 0 (for 0 mL added). Press the ENTER key to store the first data pair.

b. Add a small amount of the titrant, up to 0.50 mL. When the pH stabilizes click . Enter the current buret reading and press ENTER to save the second data pair.

c. Continue adding the HCl solution in increments that lower the pH consistently, and enter the buret reading after each increment.

d. When you reach the equivalence point, continue adding the HCl solution until the pH value remains constant.

11. When you have finished collecting data, click . Choose Store Latest Run from the Experiment menu to save the results of the first trial.

Equilibrium Labs Name: _____________________________________ of 1912. Follow the steps below to find the equivalence point, which is the largest increase in pH upon the

addition of a very small amount of NaOH solution. A good method of determining the precise equivalence point of the titration is to take the second derivative of the pH-volume data, a plot of 2pH/vol2.a. View a plot of the second derivative on Page 3 by clicking on the Next Page button, .b. Analyze the second derivative plot and record the volume of NaOH at the equivalence point.

13. Dispose of the reaction mixture as directed. Rinse the pH Sensor with distilled water in preparation for the second titration.

14. Repeat the necessary steps to titrate a second sample of the filtered Ca(OH)2 solution. Conduct a third trial if directed by your instructor. Print a copy of the graph (or graphs) that you will use in your data analysis.

DATA TABLE email with graphs (use print screen or print as PDF

Trial Equivalence point (mL)

1

2

DATA ANALYSIS email1. Calculate the [OH–] from the results of your titrations. Explain your calculations.2. Calculate the [Ca2+]. Explain your calculations.3. Calculate the Ksp for calcium hydroxide. Explain your calculations.

4. Find the accepted value of the Ksp for calcium hydroxide and compare it with your value. Discuss the discrepancy and suggest possible sources of experimental error.


Partner 2: The Determination ofan Equilibrium Constant (email)

Chemical reactions occur to reach a state of equilibrium. The equilibrium state can be characterized by quantitatively defining its equilibrium constant, Keq. In this experiment, you will determine the value of Keq for the reaction between iron (III) ions and thiocyanate ions, SCN–.

Fe3+ (aq) + SCN– (aq) → FeSCN2+ (aq)

The equilibrium constant, Keq, is defined by the equation shown below.

To find the value of Keq, which depends only upon temperature, it is necessary to determine the molar concentration of each of the three species in solution at equilibrium. You will use a colorimeter to help you measure the concentrations (see Figure 1). The amount of light absorbed by a colored solution is proportional to its concentration. The red FeSCN2+ solution absorbs blue light, and it will be analyzed at 470 nm (blue light).

Figure 1

In order to successfully evaluate this equilibrium system, it is necessary to conduct three separate tests. First, you will prepare a series of standard solutions of FeSCN2+ from solutions of varying concentrations of SCN– and constant concentrations of H+ and Fe3+ that are in stoichiometric excess. The excess of H+ ions will ensure that Fe3+ engages in no side reactions (to form FeOH2+, for example). The excess of Fe3+ ions will make the SCN– ions the limiting reagent, thus all of the SCN– used will form FeSCN2+ ions. The FeSCN2+ complex forms slowly, taking at least one minute for the color to develop. It is best to take absorbance readings after a specific amount of time has elapsed, between two and four minutes after preparing the equilibrium mixture. Do not wait much longer than four minutes to take readings, however, because the mixture is light sensitive and the FeSCN2+ ions will slowly decompose.

In Part II of the experiment, you will analyze a solution of unknown [SCN–] by using the same procedure that you followed in Part I. In this manner, you will determine the molar concentration of the SCN– solution.

Third, you will prepare a new series of solutions that have varied concentrations of the Fe3+ ions and the SCN– ions, with a constant concentration of H+ ions. You will use the results of this test to accurately evaluate the equilibrium concentrations of each species.

Equilibrium Labs Name: _____________________________________ of 19OBJECTIVESIn this experiment, you will

Prepare and test standard solutions of FeSCN2+ in equilibrium. Test solutions of SCN– of unknown molar concentration. Determine the molar concentrations of the ions present in an equilibrium system. Determine the value of the equilibrium constant, Keq, for the reaction.

MATERIALSVernier computer interfacecomputer

0.200 M iron (III) nitrate, Fe(NO3)3, solution in 1.0 M HNO3

Vernier ColorimeterTemperature Probe (optional)

0.0020 M iron (III) nitrate, Fe(NO3)3, solution in 1.0 M HNO3

plastic cuvette potassium thiocyanate, KSCN solution offour 10.0 mL pipettes unknown concentrationpipet pump or bulb 0.0020 M thiocyanate, SCN–

six 20 × 150 mm test tubes test tube rack50 mL volumetric flask distilled watereight 100 mL beakers tissue plastic Beral pipets

PRE-LAB EXERCISE Print this data tableFor the solutions that you will prepare in Step 2 of Part I below, calculate the [FeSCN2+]. Presume that all of the SCN– ions react. In Part I of the experiment, mol of SCN– = mol of FeSCN2+. Thus, the calculation of [FeSCN2+] is: mol FeSCN2+ ÷ L of total solution. Record these values in the table below.

Beaker number [FeSCN2+]

1 0.00 M

2

3

4

5

Equilibrium Labs Name: _____________________________________ of 19PROCEDUREPart I Prepare and Test Standard Solutions

1. Obtain and wear goggles.

2. Label five 100 mL beakers 1–5. Obtain small volumes of 0.200 M Fe(NO3)3, 0.0020 M SCN–, and distilled water. CAUTION: Fe(NO3)3 solutions in this experiment are prepared in 1.0 M HNO3 and should be handled with care. Prepare four solutions according to the chart below. Use a 10.0 mL pipet and a pipet pump or bulb to transfer each solution to a 50 mL volumetric flask. Mix each solution thoroughly. Measure and record the temperature of one of the above solutions to use as the temperature for the equilibrium constant, Keq.

Beaker number

0.200 M Fe(NO3)3

(mL)0.0020 M SCN–

(mL)H2O(mL)

1 5.0 0.0 45.02 5.0 2.0 43.03 5.0 3.0 42.04 5.0 4.0 41.05 5.0 5.0 40.0

3. Connect a Colorimeter to Channel 1 of the Vernier computer interface. Connect the interface to the computer with the proper cable.

4. Start the Logger Pro program on your computer. Open the file “10 Equilibrium” from the Advanced Chemistry with Vernier folder.

5. Calibrate the Colorimeter.a. Prepare a blank by filling an empty cuvette ¾ full with distilled water. Place the blank in the cuvette

slot of the Colorimeter and close the lid.b. If your Colorimeter has a CAL button, set the wavelength on the Colorimeter to 470 nm, press the CAL

button, and proceed directly to Step 6. If your Colorimeter does not have a CAL button, continue with this step to calibrate your Colorimeter.

c. Choose Calibrate CH1: Colorimeter from the Experiment menu, then click .d. Turn the wavelength knob on the Colorimeter to the “0% T” position.e. Type 0 in the edit box.f. When the displayed voltage reading for Reading 1 stabilizes, click .g. Turn the knob of the Colorimeter to the Blue LED position (470 nm).h. Type 100 in the edit box.i. When the voltage reading for Reading 2 stabilizes, click , then click .

6. You are now ready to collect absorbance data for the standard solutions. Click to begin data collection. Note: Take readings within 4 minutes of preparing the mixtures.c. Empty the water from the cuvette. Using the solution in Beaker 1, rinse the cuvette twice with ~1 mL

amounts and then fill it ¾ full. Wipe the outside with a tissue, place it in the Colorimeter, and close the lid. Wait for the absorbance value displayed in the Meter window to stabilize. Click , type the concentration of FeSCN2+ (from your pre-lab calculations) in the edit box, and press the ENTER key.

d. Discard the cuvette contents as directed. Rinse and fill the cuvette with the solution in Beaker 2. Follow the procedure in Part a of this step to measure the absorbance, and enter the concentration of this solution.

e. Repeat Part b of this step to measure the absorbance of the solutions in Beakers 3, 4, and 5.f. Click when you have finished collecting data. Click the Examine button, , and record the

absorbance values for each data pair.

Equilibrium Labs Name: _____________________________________ of 197. Click the Linear Fit button, . A best-fit linear regression line will be shown for your five data points.

This line should pass near or through the data points and the origin of the graph. (Note: Another option is to choose Curve Fit from the Analyze menu, and then select Proportional. The Proportional fit has a y-intercept value equal to 0; therefore, this regression line will always pass through the origin of the graph).

Leave the graph and best fit line displayed and proceed to Step 8.

Part II Test an Unknown Solution of SCN–

8. Obtain about 10 mL of the unknown SCN– solution. Use a pipet to measure out 5.0 mL of the unknown into a clean and dry 100 mL beaker. Add precisely 5.0 mL of 0.200 M Fe(NO3)3 and 40.0 mL of distilled water to the beaker. Stir the mixture thoroughly.

9. Using the solution in the beaker, rinse a cuvette twice with ~1 mL amounts and then fillit ¾ full. Wipe the outside with a tissue, place it in the Colorimeter, and close the lid. Watch the absorbance readings in the Meter window. When the readings stabilize, record the absorbance value for your unknown in your data table. Remove and clean the cuvette.

10. Determine the concentration of the unknown SCN– solution.a. With the linear-regression curve still displayed on your graph, choose Interpolate from the Analyze

menu.b. A vertical line with a cursor now appears on the graph. The cursor’s concentration and absorbance

coordinates are displayed in the floating box.c. Move the cursor along the regression line until the absorbance value is approximately the same as the

absorbance value you recorded in Step 9. The corresponding concentration value is the concentration of the unknown solution, in mol/L. Record this value in your data table.

Part III Prepare and Test Equilibrium Systems

11. Prepare five test tubes of solutions, according to the chart below. Follow the necessary steps from Part I to test the absorbance values of each mixture. Record the test results in your data table. Note: You are using 0.0020 M Fe(NO3)3 in this test.

Test tubenumber

0.0020 M Fe(NO3)3

(mL)0.0020 M SCN–

(mL)H2O(mL)

1 3.00 0.00 7.002 3.00 2.00 5.003 3.00 3.00 4.004 3.00 4.00 3.005 3.00 5.00 2.00

12. To get good data for the calculation of Keq, you must determine the net absorbance of the solutions in Test Tubes 2–5. To do this, subtract the absorbance reading for Test Tube 1 from the absorbance readings of Test Tubes 2–5, and record these values as net absorbance in your data table.

Equilibrium Labs Name: _____________________________________ of 19DATA TABLE email the last two pages to me and any graphs generated by this lab.Parts I and II

Beaker Absorbance

1

2

3

4

5

Unknown, Part II

Best-fit line equation for the Part I standard solutions:

Part III

Test tube number Absorbance Net absorbance

1

2

3

4

5

Equilibrium Labs Name: _____________________________________ of 19DATA ANALYSIS1. (Part II) Use the calibration equation from Item 1 and the absorbance reading for your unknown solution

to determine [SCN–].

2. (Part II) Compare your experimental [SCN–], of your unknown, with the actual [SCN–]. Suggest reasons for the disparity.

3. (Part III) Use the net absorbance values, along with the best fit line equation of the standard solutions in Part I to determine the [FeSCN2+] at equilibrium for each of the mixtures that you prepared in Part III. Complete the table below and give an example of your calculations.

Test tube number 2 3 4 5

[FeSCN2+]

4. (Part III) Calculate the equilibrium concentrations for Fe3+ and SCN– for the mixtures in Test tubes 2-5 in Part III. Complete the table below and give an example of your calculations.

Test tube number 2 3 4 5

[Fe3+]

[SCN–]

5. Calculate the value of Keq for the reaction. Explain how you used the data to calculate Keq.


LAB: Heat of Reaction between Vinegar and Baking SodaIntroduction (email the last two pages of this lab)

The chemical change that occurs when vinegar (5% acetic acid) and baking soda (sodium bicarbonate) are mixed together is an example of a gas forming acid/base reaction. The weak acid, acetic acid, reacts with the weak base, sodium bicarbonate, to form sodium acetate, carbon dioxide gas and liquid water. In this experiment, we will be combining solid sodium bicarbonate with vinegar.

CH3COOH(aq) + NaHCO3(s) HNa(s) + CH3COO(aq) + CO3(g)

CH3COOH(aq) + NaHCO3(s) NaCO2(g) + H2O(l) + CH3COO(aq)

CH3COOH(aq) + NaHCO3(s) NaCH3COO(aq) + H2O(l) + CO2(g)

CH3COOH(aq) + NaHCO3(s) NaH2O(aq) + COOH(l) + CO2(g)

Experimental The chemical change that occurs when vinegar (5% acetic acid) and baking soda (sodium bicarbonate) are mixed together is an example of a gas forming acid/base reaction. The weak acid, acetic acid, reacts with the weak base, sodium bicarbonate, to form sodium acetate, carbon dioxide gas and liquid water. In this experiment, we will be combining solid sodium bicarbonate with vinegar. Recall from the introduction that the equation describing this process is:

Materials

Baking Soda {NaHCO3(s)}

Vinegar (5% by wt CH3COOH)

Foam cups (calorimeter)

Measuring spoons

Thermometer


Procedure1. Accurately measure out 60 ml of vinegar and pour it into the calorimeter. Determine the mass

of the liquid by assuming the density of vinegar to be the same as that of pure water, 1.0 g/ml. Compute the mass of acetic acid and the number of moles of acetic acid present in the calorimeter. Enter this information Data Sheet on the next page.

2. Carefully measure 1 teaspoon of baking soda (recall that 1 teaspoon = 5 ml). Assume baking soda to be pure sodium bicarbonate (density = 2.16 g/ml). Determine the mass and the number of moles sodium bicarbonate in the teaspoon. Enter this information in Table I of your data sheet.

3. Immerse your thermometer into the vinegar. Record the temperature in Table II of your data sheet. Next, quickly add all of the baking soda and agitate the mixture (note: if you are using a 12 oz cup the mixture may bubble over, carefully add the baking soda or use a larger foam cup). Observe the temperature continuously until the it stops changing. Record the final temperature in Table II of your data sheet.

4. Measure the volume of the solution by pouring it into a graduated cylinder. Assuming that the mixture has a density of 1 g/ml, determine its mass. Record the solution's mass and volume in Table II of your data sheet.

5. Compute the change in temperature (T), qsolution, qrxn and H for the reaction as written in equation 2. Enter these data in Table III of your data sheet.

Substance volume (ml) mass (g) moles Tinitial (oC) =

vinegar xxxxxxx Tfinal (oC)=

acetic acid xxxxxxx Vsolution (ml)=

baking soda ng (g)=

T (oC)=

qsolution (J)=

qreaction (J)=

Enthalpy of Reaction II: Vinegar and Baking Soda

Table IQuantities of Materials

Table IILaboratory Measurements

Table IIIComputations


H (kJ) =


Results and DiscussionOnce you have compared your results with the other members of your lab team, collaborate with your lab team to write a lab report. Your report should include a summary of the lab procedure completed in your kitchen, examples of the calculations you made, an explanation of the results of the calculations, and a conclusion. The conclusion should use key terms related to the lab, such as endothermic, exothermic, heat transfer, enthalpy, etc. Your conclusion should also address the following questions.

1. Why do we use the limiting reactant data to calculate H?

2. Why might it be important to know H prior to running a reaction?

3. If we increased the amount of acetic acid and sodium bicarbonate to 2.0 moles, what effect would this have on the reaction? The enthalpy change? Explain.

4. If we increased the amount of acetic acid and sodium bicarbonate to 50 moles, what effect would this have on the reaction? The enthalpy change? Explain.

5. Paragraph discussion of conclusions and Heat of Reaction.

Equilibrium Labs Name: _____________________________________ of 19Solution Equilibrium and Ksp (Email your word doc. to me.) Use PhET on the desktop or go to http://phet.colorado.edu/simulations/index.phpNon-Obvious Controls for the simulation:

Be sure to try all the different tabs at the top of the simulation. The model increases in difficulty as you go from Table Salt to the right

You can Pause the simulation and then use Step in incrementally analyze.

Launch Soluble Salts and start a Word document to write a lab report.

Learning Goals: Students will be able to: Describe the equilibrium of a saturated solution macroscopically and microscopically with supporting

illustrations. Write equilibrium expressions for salts dissolving Calculate K sp from molecular modeling.

1. Observe what happens as you add one shaker of salt to the water. Talk about your observations and then investigate salts dissolving in water further. When you feel like you understand what equilibrium means for a salt dissolving in water, write an introduction for your lab report that explains your understanding of equilibrium. Illustrate your introduction with “test tube” size drawings and “close-up” views to show the ions and crystals. Some things to think about are:

a. In general terms, what the reactant is and what the products are when you put a salt in water.b. What would a test tube of the salt/water equilibrium look like?c. What is happening on a molecular scale when equilibrium is established? d. How does the speed at which you add the salt effect the equilibrium? e. How does the volume of water or amount of salt added affect the equilibrium?

2. Design experiments to determine the value of the constant for each salt include the effect of varying volume.

a. Write your procedureb. Show sample calculations.c. Make a data table for each salt that demonstrates good experimental design.d. Write the equilibrium expressions for each salt with the determined constant.

3. Write a conclusion for your experiment that includes addressing these questions.a. Which salt gave you the best data? Explain your reasoning.b. Which salt gave you the poorest data? Explain your reasoning.c. How do your values compare to the published ones? Cite your references. d. How do the solubility rules relate to the Ksp values that you determined? e. How could you use Ksp values to predict solubility?

Equilibrium Labs Name: _____________________________________ of 19Experiment: Warming Curve

Purpose-determine heat required to heat water. Take approximately 100 g of ice and put the ice in a beaker. Record the temperature every minute. (The first temperature should be taken before the beaker is

put on the hot plate and continue every minute thereafter.) STIR THE Water/Ice BEFORE EACH TEMPERATURE READING.

Heat on a hot plate (on medium heat) until the water has come to a rolling boil for 5 minutes. Mass the water after cooling. Calculate the calories and joules of heat required for this lab. See below. DO A FORMAL LAB REPORT. Use Excel for your data and graph, Purpose, analysis section (see next step), and conclusion should be done as a wiki. Take do a print screen to put your data into your wiki page. In the conclusion express the significance of the graph shape (flat part what was happening when the graph was flat.) The math analysis is summarized above. You have to melt the ice, warm the water, and measure the amount of water that is converted to steam. So do three calculations and add them together.

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