Chapter 13. Laboratory: Chemical Equilibrium and Le Chatelier’s Principle
It’s important to understand two things about Le Chatelier’s Principle. First, the equilibrium change can (almost) never return the system to its original equilibrium state. For example, if you add reactants to a system that has achieved dynamic chemical equilibrium, the concentration of reactants at the new equilibrium state will be higher than it was at the original equilibrium state (although lower than it would have been without the change in equilibrium state). Second, Le Chatelier’s Principle is qualitative, rather than quantitative. That is, applying Le Chatelier’s Principle tells you whether the reaction will be forced to the left (more reactants) or to the right (more products), but it cannot tell you how far the reaction will be forced or what the eventual concentrations of reactants and products will be.
In this chapter, we’ll examine chemical reactions and Le Chatelier’s Principle.
Laboratory 13.1: Observe Le Chatelier’s Principle in Action
Le Chatelier’s Principle states that a forced change in concentration, temperature, volume, or pressure results in change to the dynamic equilibrium state. In this lab, we’ll examine various reactions that are in a state of dynamic chemical equilibrium, force changes to each of these environmental characteristics, and observe the results.
Procedure
This lab has four parts, each of which examines one of the effects of concentration, temperature, volume, and pressure on dynamic equilibrium states. As you observe each experiment, record your observations in the Observations section that follows the procedure sections.
Caution
With the exception of concentrated hydrochloric acid, the chemicals used in this laboratory are reasonably safe. Hydrochloric acid is toxic, corrosive, and produces irritating fumes. This experiment uses an open flame, so use caution and have a fire extinguisher handy. Wear splash goggles, gloves, and protective clothing.
Part I: Examine the Effect of Concentration on Equilibrium
In this part of the lab, we examine the effect of changing the concentration of the reactant ions on a saturated solution of sodium chloride (NaCl). That solution contains sodium ions (Na+) and chloride ions (Cl–). To saturated sodium chloride solution samples, we’ll add concentrated hydrochloric acid (which contains Cl– ions but not Na+ ions), saturated sodium carbonate solution (which contains Na+ ions but not Cl– ions), and saturated magnesium sulfate solution (which contains neither), and observe the effects on the saturated sodium chloride solution.
goggles, gloves, and protective clothing | |
dropper or Beral pipette (3) | |
test tubes (6) and rack | |
beaker, 150 mL (2) | |
alcohol lamp or other heat source | |
ring stand | |
support ring | |
wire gauze | |
stirring rod | |
tap water | |
sodium chloride (~40 g) | |
sodium carbonate (~1.6 g) | |
magnesium sulfate (~1.3 g) | |
hydrochloric acid, concentrated (~5 mL) | |
small bottle of carbonated soft drink (2, one chilled and one at room temperature) |
Disposal
All of the solutions used in this laboratory can safely be flushed down the drain with copious water.
If you have not already done so, put on your splash goggles, gloves, and protective clothing.
Prepare a saturated solution of sodium carbonate by adding about 1.6 g of sodium carbonate to 5 mL of water in a test tube. Stir or shake the solution to ensure that the sodium carbonate dissolves completely, and continue adding sodium carbonate until some of it remains undissolved at the bottom of the test tube.
Prepare a saturated solution of magnesium sulfate by adding about 1.3 g of magnesium sulfate to 5 mL of water in a test tube. Stir or shake the solution to ensure that the magnesium sulfate dissolves completely, and continue adding magnesium sulfate until some of it remains undissolved at the bottom of the test tube.
Prepare a saturated solution of sodium chloride by adding about 40 g of sodium chloride to 100 mL of water in a small beaker or flask. Stir or shake the solution to ensure that the sodium chloride dissolves completely, and continue adding sodium chloride until some of it remains undissolved at the bottom of the beaker or flask.
Transfer about 5.0 mL of the saturated sodium chloride solution to each of four test tubes, labeled A through D, one of which will serve as the control. The exact amount of solution is not critical, but make sure that each test tube contains the same amount.
Place the test tubes adjacent to each other in a test tube rack under a strong light. (It may be easier to observe the reactions if you place a sheet of black construction paper or a similar material behind the test tube rack.)
Add concentrated hydrochloric acid dropwise to test tube A, observing any change that occurs as each drop is added. Continue adding hydrochloric acid until you have added about 5 mL.
Add saturated sodium carbonate solution dropwise to test tube B, observing any change that occurs as each drop is added. Continue adding the saturated sodium carbonate solution until you have added about 5 mL.
Add saturated magnesium sulfate solution dropwise to test tube C, observing any change that occurs as each drop is added. Continue adding the saturated magnesium sulfate solution until you have added about 5 mL.
Add water dropwise to test tube D, the control test tube, observing any change that occurs as each drop is added. Continue adding water until you have added about 5 mL.
Part II: Examine the Effect of Temperature on Equilibrium
In our saturated solution of sodium chloride, solid sodium chloride exists in equilibrium with aqueous sodium ions and chloride ions. Le Chatelier’s Principle tells us that changing the temperature of the reaction environment will cause a change in that dynamic chemical equilibrium, but we have no way of knowing the direction of that change. Intuitively, we expect that more sodium chloride will dissolve in a given volume of water at a higher temperature, so let’s test that hypothesis:
If you have not already done so, put on your splash goggles, gloves, and protective clothing.
Transfer about half of the remaining saturated sodium chloride solution (~40 mL) to a small beaker or flask.
Add a small amount of sodium chloride to the beaker or flask; just enough so that undissolved sodium chloride crystals are visible on the bottom of the beaker or flask.
Set up your tripod stand, wire gauze, and alcohol burner.
With stirring, gently heat the beaker or flask that contains the saturated sodium chloride solution, keeping an eye on the undissolved sodium chloride.
Part III: Examine the Effect of Volume on Equilibrium
In our saturated solution of sodium chloride, solid sodium chloride exists in equilibrium with aqueous sodium ions and chloride ions. Le Chatelier’s Principle tells us that changing the volume of the solvent will cause a change in that dynamic chemical equilibrium. Intuitively, we expect that if X g of sodium chloride dissolves in Y mL of water, increasing the amount of water will increase the amount of sodium chloride that dissolves. Let’s test that hypothesis:
If you have not already done so, put on your splash goggles, gloves, and protective clothing.
Make sure that the remaining saturated sodium chloride solution (~40 mL) has some undissolved sodium chloride visible in the bottom of the beaker or flask. Note the approximate amount of undissolved sodium chloride.
Add 10 drops of water to the beaker or flask, and stir or swirl the solution to determine whether any additional sodium chloride dissolves.
Continue adding water 10 drops at a time until all of the visible undissolved sodium chloride has dissolved.
PART IV: Examine the Effect of Pressure (and Temperature) on Equilibrium
Le Chatelier’s Principle says that changing the pressure of the reaction environment will change the dynamic chemical equilibrium. Carbonated soft drinks contain dissolved carbon dioxide gas, which is more soluble at higher pressures. If we reduce the pressure, we expect the soft drink solution to reduce the concentration of dissolved carbon dioxide by liberating carbon dioxide gas. (You may want to do this experiment outdoors to avoid making a mess indoors.) Let’s test that hypothesis:
If you have not already done so, put on your splash goggles, gloves, and protective clothing.
Leaving it capped, shake the chilled sealed bottle of carbonated soft drink. (Agitating the sealed bottle mixes small bubbles of the carbon dioxide gas in the “empty” part of the bottle into the liquid, providing loci for the dissolved carbon dioxide to come out of solution.)
Immediately after you give the carbonated soft drink bottle a thorough shaking, point the mouth in a safe direction and remove the cap.
Repeat steps 2 and 3 with the bottle of carbonated soft drink that’s at room temperature.
Clean up.
Observations
If you are not maintaining a lab notebook, record your observations narratively here.
Part I
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Part II
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Part III
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Part IV
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Review Questions
Laboratory 13.2: Quantify the Common Ion Effect
In Part I of the preceding laboratory, we examined the effect on equilibrium of changing the concentration of one of the reactant ions. We found that increasing the concentration of either reactant ion (Na+ or Cl–) by adding sodium carbonate or hydrochloric acid, respectively, forced solid sodium chloride to precipitate from the saturated solution. This phenomenon is called the common ion effect.
In addition to affecting solubility, the common ion effect reduces the ionization of weak acids and weak bases in solution. Acetic acid is a weak organic acid that dissociates only partially in aqueous solutions, yielding hydronium (H3O+) ions and acetate (CH3COO–) ions:
CH3COOH(l) + H2O ⇔ CH3COO–(aq) + H3O+(aq)
Sodium acetate is an ionic compound that dissociates essentially completely in aqueous solution, yielding sodium (Na+) ions and acetate (CH3COO–) ions.
CH3COONa(s) → CH3COO–(aq) + Na+(aq)
The common ion effect tells us that acetic acid should dissociate less completely in sodium acetate solution than in pure water, because the sodium acetate increases the concentration of acetate ions. Less dissociation means fewer protons, and an accordingly higher pH. In this lab, we’ll test that hypothesis by measuring the pH of a fixed concentration of acetic acid in solutions that contain various concentrations of sodium acetate.
Caution
Glacial acetic acid is corrosive, toxic, and has a very strong, biting odor. Wear splash goggles, gloves, and protective clothing.
Procedure
If you have not already done so, put on your splash goggles, gloves, and protective clothing.
Label six test tubes A through F.
Prepare a concentrated (nearly saturated) solution of sodium acetate by adding 4.0 g of anhydrous sodium acetate to 7 mL of water in the graduated cylinder and then adding water to bring the total solution volume to 10 mL. Stir or shake the solution to ensure that the sodium acetate dissolves completely. We’ll define this concentration of sodium acetate as 100%.
Transfer 5.0 mL of the 100% sodium acetate solution to test tube A and 5.0 mL to test tube B.
Add 5.0 mL of water to test tube B and mix thoroughly.
Transfer 5.0 mL of the solution in test tube B to test tube C, add 5.0 mL of water to test tube C, and mix thoroughly.
Repeat these steps until the five test tubes labeled A through E contain, respectively, 5.0 mL each of sodium acetate solutions of 100%, 50%, 25%, 12.5%, and 6.25%.
Transfer 5.0 mL of water to test tube F.
Add 1.0 mL of glacial acetic acid to test tube A and mix thoroughly.
Following the directions supplied with your pH meter, test the pH of the solution and record the data in Table 13-1.
Repeat steps 9 and 10 for test tubes B through F.
After 15 minutes have passed, retest the pH of each of the six test tubes. Repeat the test again after 60 minutes. Record the data in Table 13-1.
Review Questions
Laboratory 13.3: Determine a Solubility Product Constant
The coexistence of an ionic solid and its component ions in solution is one example of a chemical equilibrium. This equilibrium can be quantified using the solubility product principle, which states that in a saturated solution of an ionic compound, the product of the molar activities, called the solubility product constant or Ksp, has a constant value at any particular temperature and pressure.
Because ionic activities are very difficult to determine accurately and molarities are easy to determine, solubility product calculations are normally performed using molarity values. For dilute solutions, using molarities rather than activities introduces only tiny errors, because molarity and activity are nearly identical in dilute solutions. For concentrated solutions, molarity and activity may differ significantly, so solubility product calculations, although useful for sparingly soluble ionic compounds, are of limited use for very soluble ionic compounds.
Consider the sparingly soluble salt silver chloride. In a saturated aqueous solution of silver chloride, the silver and chlorine are present primarily as solvated silver ions and chlorine ions. (Solvated ions are ions that have bound to solvent molecules.) Undissociated molecular silver chloride exists in the solution, but it is present in such a tiny amount that it can be ignored. The following equilibrium equation describes the saturated silver chloride solution:
AgCl(s) ⇔ Ag+(aq) + Cl–(aq)
The solubility product constant for this equilibrium is:
Ksp = [Ag+] · [Cl–]
Square brackets in equilibrium expressions indicate molar concentration, so this expression is a short way to say that the solubility product constant is equal to the molar concentration of the silver ions multiplied by the molar concentration of the chloride ions. Because one molecule of silver chloride dissociates into one silver ion and one chloride ion, the concentrations of the silver ions and chloride ions are identical. Assigning that concentration the value x simplifies the expression to:
Ksp = x · x = x2
goggles, gloves, and protective clothing | |
balance and weighing paper | |
beaker, 250 mL (2) | |
volumetric flask, 100 mL | |
eye dropper or disposable pipette | |
thermometer | |
stirring rod | |
funnel | |
filter paper | |
ring stand | |
support ring | |
wire gauze | |
burette clamp | |
gas burner | |
sodium chloride (~50 g) | |
potassium hydrogen tartrate (~2 g) | |
sodium hydroxide, 0.1000 M, standardized (~50 mL) | |
phenolphthalein indicator solution (a few drops) |
which means that the concentration of either ion is the square root of the Ksp. That means that if we know the Ksp, we can determine the molar concentration, and vice versa. For example, if we know that the Ksp of silver chloride at a particular temperature and pressure is 1.8 · 10–10, we can rewrite the equilibrium expression as:
1.8 · 10–10 = x2
Solving for x (taking the square root of the Ksp) tells us that the molar concentrations of the silver ions and chloride ions are both 1.3 · 10–5, so a saturated solution of silver chloride at this temperature and pressure is 0.000013 M. The gram molecular mass of silver chloride is 143.32 g/mol, so a saturated silver chloride solution contains about 0.0019 g/L.
Conversely, if we know the molar solubility, we can calculate the Ksp. For example, if we know that a saturated solution of silver bromide at a particular temperature and pressure is 7.2 · 10–7 M, we can calculate the Ksp of silver bromide as follows:
Ksp = (7.2 · 10–7)2 = 5.2 · 10–13
Until now, we’ve considered only the simplest case, compounds in which a molecule dissociates into two ions. For ionic compounds that dissociate into three or more ions, the calculations are a bit more involved. Consider silver chromate, which dissociates as follows:
Ag2CrO4(s) ⇔ 2 Ag+(aq) + CrO42–(aq)
We can no longer use simple squares and square roots, because each molecule of silver chromate dissociates into three ions—two silver ions and one chromate ion. We can generalize the solubility product constants for a salt AxBy that dissociates into the ions Ay+ and Bx– in the form:
Ksp = [Ay+]x· [Bx–]y
so, for silver chromate, Ag2(CrO4)1, the solubility product constant can be expressed as:
Ksp = [Ag+]2·[Cr2O42–]1
If we assign the molar concentration of the chromate ion as x, and knowing that the concentration of the silver ion is twice that of the chromate ion, the expression becomes:
Ksp = [2x]2· [x]1 = 4x3
If we make up a saturated solution of silver chromate, we might determine by titration that the concentration of silver ion is 1.3 · 10–4 M, from which it follows that the concentration of the chromate ion must be 0.65 · 10–4 M, which can also be expressed as 6.5 · 10–5 M. Plugging these values into the equation yields:
Ksp = (1.3 · 10–4)2· (6.5 · 10–5)1 = 1.1 · 10–12
If the Ksp is known, we can determine the molar solubility of silver chromate as follows:
Ksp = 1.1 · 10–12 = 4x3
x3 = (1.1 · 10–12/4) = 2.8 · 10–13
x = 6.5 · 10–5
In principle, you can use solubility product constants to calculate molar solubilities of ionic salts. In practice, observed solubilities often differ significantly from such calculated values. Electrostatic interaction between ions can affect the results, particularly in any but very dilute solutions, and other equilibria occurring simultaneously in the solution can have a major impact. For example, the solubility of magnesium hydroxide varies greatly with the pH of the solution because of the common ion effect. Magnesium hydroxide is much more soluble in acid solutions, which have a very low [OH–] and correspondingly less soluble in basic solutions, which have a high [OH–]. Finally, the assumption that ionic substances fully dissociate in solution into simple solvated ions is not always true. For example, magnesium fluoride (MgF2) does not dissociate completely into Mg2+ and F– ions. Instead, some partially dissociated magnesium fluoride is present as MgF+ ions.
In this lab, we’ll determine solubility product constants for two compounds, using two different methods.
Caution
Glacial acetic acid is corrosive, toxic, and has a very strong, biting odor. Wear splash goggles, gloves, and protective clothing.
Procedure
This lab session is in two parts. In Part I, we’ll determine the solubility product constant for sodium chloride gravimetrically, by evaporating a known volume of a saturated sodium chloride solution and determining the mass of the dry solid remaining. In Part II, we’ll determine the solubility product constant for potassium hydrogen tartrate titrimetrically (volumetrically) by titrating a known volume of saturated potassium hydrogen tartrate solution with sodium hydroxide titrant of known molarity.
Part I: Determine the Ksp of Sodium Chloride Gravimetrically
If you have not already done so, put on your splash goggles, gloves, and protective clothing.
Weigh a clean, dry, empty 250 mL beaker and record its mass to 0.01 g on line A of Table 13-2.
Add about 125 mL of room temperature distilled or deionized water and about 50 g of sodium chloride to the 250 mL beaker. Allow the salt to dissolve for at least 15 minutes, stirring periodically, to ensure that the solution is saturated. Record the temperature of the solution on line B of Table 13-2.
Carefully decant 100.00 mL of the saturated sodium chloride solution into the 100 mL volumetric flask, making sure that none of the undissolved sodium chloride is transferred to the volumetric flask. Fill the flask to the index line, using a dropper or disposable pipette to add the last mL or so, and record the volume on line C of Table 13-2. (If you’ve calibrated the volumetric flask, as described in Chapter 5, record the actual volume that the flask contains rather than the nominal 100.00 mL volume.)
Discard the remaining sodium chloride solution and undissolved sodium chloride and rinse the beaker thoroughly, first with tap water and then with distilled water.
Pour the contents of the 100 mL volumetric flask into the 250 mL beaker. Do a quantitative transfer, rinsing the volumetric flask several times with a few mL of distilled water and adding the rinse water to the beaker to make sure all of the sodium chloride is transferred to the beaker.
Place a stirring rod in the beaker to prevent boil-over, and heat the beaker until the solution comes to a gentle boil. Continue heating the beaker gently until nearly all of the water has evaporated. When most of the water has vaporized, remove the stirring rod from the beaker, and use a few mL of distilled water to rinse any sodium chloride into the beaker.
Continue heating the beaker to drive off all remaining water. When you have evaporated all of the remaining water, remove the heat and allow the beaker to cool to room temperature.
Weigh the beaker and record the mass of the beaker plus sodium chloride to 0.01 g on line D of Table 13-2.
Subtract the empty mass of the beaker from the mass of the beaker plus sodium chloride, and record the mass of sodium chloride on line E of Table 13-2.
Calculate the solubility of sodium chloride in g/L and enter that value on line F of Table 13-2.
The gram molecular mass of sodium chloride is 58.442 g/mol. Use that value to calculate the molar solubility of sodium chloride and enter that value on line G of Table 13-2.
The Ksp of sodium chloride is the product of the activities (concentrations) of the sodium ion and the chloride ion: Ksp = [Na+] · [Cl–]. Those concentrations are the same, and are equal to the concentration of sodium chloride, so we can simplify the equation to: Ksp = [NaCl]2. Therefore, the solubility product constant of sodium chloride equals the square root of the molar solubility of sodium chloride. Calculate that value and enter it on line H of Table 13-2.
Item | Data |
A. Mass of empty beaker | _________._____ g |
B. Temperature of solution | _________._____ °C |
C. Volume of sodium saturated sodium chloride solution | _________._____ mL |
D. Mass of beaker + sodium chloride | _________._____ g |
E. Mass of sodium chloride (D – A) | _________._____ g |
F. Mass solubility of sodium chloride ([1000/C] · E) | _________._____ g/L |
G. Molar solubility of sodium chloride (F/58.442 g/mol) | _________._____ mol/L |
H. Ksp of sodium chloride (square root of G) | _________._____ |
Part II: Determine the Ksp of Potassium Hydrogen Tartrate Titrimetrically (volumetrically)
If you have not already done so, put on your splash goggles, gloves, and protective clothing.
To about 150 mL of distilled or deionized water in the 250 mL beaker, add about 2 g of potassium hydrogen tartrate. Allow the salt to dissolve for at least 15 minutes, stirring periodically, to ensure that the solution is saturated. Record the temperature of the solution on line A of Table 13-3.
Allow the solution to sit undisturbed for several minutes until most of the undissolved material has settled to the bottom of the beaker. Carefully filter the supernatant liquid through a dry filter paper into a second 250 mL beaker, trying to keep as much as possible of the undissolved solid in the first beaker.
Transfer 100.00 mL of the saturated potassium hydrogen tartrate solution into the 100 mL volumetric flask. Fill the flask to the index line, using a dropper or disposable pipette to add the last mL or so, and record the volume on line B of Table 13-3. (If you’ve calibrated the volumetric flask, as described in Chapter 5, record the actual volume that the flask contains rather than the nominal 100.00 mL volume.)
Discard the remaining potassium hydrogen tartrate solution and undissolved potassium hydrogen tartrate. Rinse the beaker thoroughly, first with tap water and then with distilled water.
Pour the contents of the 100 mL volumetric flask into the 250 mL beaker. Do a quantitative transfer, rinsing the volumetric flask several times with a few mL of distilled water and adding the rinse water to the beaker to make sure all of the potassium hydrogen tartrate is transferred to the beaker.
Rinse your 50 mL burette twice with 0.1000 M sodium hydroxide, running the rinse solution through the tip and into a waste container.
Clamp the 50 mL burette in a burette clamp and fill it to above the 0.00 mL index line with 0.1000 M sodium hydroxide solution.
Run solution through the burette until the level drops to or slightly below the 0.00 mL index line. Make sure that there are no bubbles in the body or tip of the burette. Record the initial burette reading as accurately as possible on line C of Table 13-3. Interpolate the reading to 0.05 mL or better.
Add a few drops of phenolphthalein indicator solution to the beaker of potassium hydrogen tartrate solution, and swirl slightly to mix.
Titrate the potassium hydrogen tartrate solution until you reach the endpoint, when the phenolphthalein gives the solution a distinct pink coloration that persists for at least 30 seconds. Record the final burette reading as accurately as possible on line D of Table 13-3. Interpolate the reading to 0.05 mL or better.
Subtract the initial burette reading from the final burette reading to determine the volume of 0.1000 M sodium hydroxide titrant that was required to neutralize the potassium hydrogen tartrate aliquot. Record that volume on line E of Table 13-3.
The gram molecular mass of sodium hydroxide is 39.9971 g/mol. Use that value and the actual molarity of the titrant solution to calculate the number of moles of sodium hydroxide required to neutralize the potassium hydrogen tartrate aliquot, and enter that value on line F of Table 13-3.
Using the number of moles of sodium hydroxide needed to neutralize the aliquot of potassium hydrogen tartrate and the volume of that aliquot, calculate the molar solubility of potassium hydrogen tartrate and enter that value on line G of Table 13-3.
Calculate the Ksp of potassium hydrogen tartrate, and enter that value on line H of Table 13-3.
Item | Data |
A. Temperature of solution | _________._____ °C |
B. Volume of sodium saturated potassium hydrogen tartrate solution | _________._____ mL |
C. Initial burette reading | _________._____ mL |
D. Final burette reading | _________._____ mL |
E. Volume of 0.1000 M sodium hydroxide titrant used (D – C) | _________._____ mL |
F. Moles of sodium hydroxide required | ______.________ mol |
G. Molar solubility of potassium hydrogen tartrate | ______.________ mol/L |
H. Ksp of potassium hydrogen tartrate | ___________________ |