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Lab 8 - Equilibrium and Le Châtelier's Principle

Purpose

To observe systems at equilibrium, and to determine what happens when stresses are applied to such systems.

Goals

Introduction

Many chemical systems are considered to be reversible. For example, drop the temperature of water to 0°C and it freezes; raise the temperature above 0°C and it melts. Many chemical reactions are also reversible. If one mixes ammonia and oxygen, the products form according to Equation 1:
( 1 )
4 NH3(g) + 3 O2(g) → 2 N2(g) + 6 H2O(g)
 
Conversely, a mixture of nitrogen and water, under the right conditions, can give ammonia and oxygen:
( 2 )
2 N2(g) + 6 H2O(g) → 4 NH3(g) + 3 O2(g)
 
Perhaps unsurprisingly, in either case one actually obtains a mixture of all four gases. A reaction in which the reactants are not completely consumed to form products because the reverse reaction also occurs (products form reactants) is a reversible reaction. Such reactions are indicated by the use of double arrows as shown in Equation 3:
( 3 )
4 NH3(g) + 3 O2(g) equilibrium arrow 2 N2(g) + 6 H2O(g)
 
In dealing with equilibrium reactions, several definitions are useful and are given below.
  • Products are the chemical species to the right of the equilibrium arrow, as the reaction equation is written.
  • Reagents are the chemical species to the left of the equilibrium arrow, as the reaction equation is written.
  • The forward reaction is the process as written from left to right in the reaction equation.
  • The reverse reaction is the process as written from right to left in the reaction equation.
In mixtures of the sort shown in Equation 3, products are constantly being transformed to reactants and vice versa. When the rate of the forward reaction is equal to the rate of the reverse reaction, the amounts of the chemical species remain constant, and the system is in a state of equilibrium. Anything that changes a variable associated with the equilibrium induces a stress on the system. If a stress is applied, the system will shift to accommodate and offset this stress, and a new equilibrium condition will be established. This principle was first articulated by a French chemist, Henri-Louis Le Châtelier, in 1884, and it still bears his name. His concise expression of the principle is:
  • If a stress is applied to a system at equilibrium, the system will respond by shifting in the direction that reduces the stress.
As an example of a stress, consider the addition of more ammonia to the equilibrium in Equation 3. To reduce this stress, some of the added ammonia reacts with oxygen to produce more products (nitrogen and water). A new equilibrium condition is established. Consequently, the amount of oxygen decreases, the amounts of nitrogen and water vapor increase, and the equilibrium shifts to the right or favors the products. If nitrogen were added to equilibrium in Equation 3, the result would be exactly the opposite and there would be a shift to the left to favor the reactants. Another way of inducing a stress on such a system is to remove a reactant or product; the system responds by replacing some of the substance that was removed. In this experiment, three equilibrium systems will be examined. All are easy to study because they involve color changes. The first is the reaction between the iron(III) ion (Fe3+) and the thiocyanate anion (SCN1-). These ions form the red complex cation* ferrithiocyanate (FeSCN2+) according to Equation 4.
( 4 )
Fe3+ + SCN-equilibrium arrowFeSCN2+
colorlessequilibrium arrowred
 
Another equilibrium involves two complex ions of cobalt:
( 5 )
CoCl42- + 6 H2Oequilibrium arrowCo(H2O)62+ + 4 Cl-
blueequilibrium arrowpink
 
The CoCl42- ion is an intense blue, the color of the patterns on Delft china. The Co(H2O)62+ ion is pale pink. You will be stressing these equilibria by adding products and reactants, and observing the color changes that result. In Part C of the experiment, students will use a spectrophotometer to calculate the equilibrium constant of bromothymol blue at three different hydronium ion (H3O+) concentrations.** According to Le Châtelier's principle, the amount of reactant and product present will adjust when a stress is applied, such that the same equilibrium constant is obtained. Bromothymol blue (HC27H27Br2O5S) is an indicator which is yellow in acidic solution and blue in basic solution. The equilibrium reaction with hydronium ion (H3O+) is shown below. For simplicity, the acid form will be represented as HBB and the base form will be represented by BB.
( 6 )
HBB(aq) + H2O(l)equilibrium arrowBB + H3O+(aq)
yellowequilibrium arrowblue
Amax ~470 nmAmax ~635 nm
 
Students will be provided with solutions at three different hydronium concentrations as buffers. A buffer solution is a solution that contains a consistent amount of hydronium ion (H3O+) and is resistant to pH changes. The hydronium ion concentration of the solution remains the same when other acids or bases are mixed with it. The relative amounts of the yellow-colored acidic form and the blue-colored basic form can be determined using a spectrophotometer. The yellow-colored acidic form absorbs maximum light (Amax) at a violet wavelength near 470 nm. The blue-colored basic form absorbs light at an orange wavelength near 635 nm. * Complex ions are formed when metals (usually transition metals) or their ions form covalent bonds with molecules or ions that have electron pairs to donate. The electron donors, species such as H2O, NH3, and halide anions, are called ligands. The electrons are shared between vacant d orbitals (or hybrid orbitals formed from them) on the metal and nonbonding pairs on the ligand. There are usually vacant d orbitals in complex ions. Visible light promotes electrons into these orbitals. Thus, complex ions absorb visible light, and have intense and beautiful colors. Chapter 14 of your CH 101 textbook has more information on the chemistry of transition metals. ** Klotz, E.; Doyle, R.; Gross, E.; Mattson, B. J. Chem. Educ. 2011, 88, 637-639.

Equipment

  • 1
    ceramic spot plate
  • 2
    glass stir rods
  • 1
    hot plate
  • 1
    250 mL beaker for waste collection
  • 1
    MicroLab spectrophotometer
  • 4
    MicroLab spectrophotometer vials
  • 1
    deionized water squirt bottle

Reagents

  • ~3 mL 0.010 M Fe(NO3)3
  • ~1 mL 0.10 M Fe(NO3)3
  • ~1 mL 0.05 M NaSCN
  • ~0.1 mL 0.10 M AgNO3
  • ~0.1 mL 1.0 M NaNO3
  • ~0.1 mL 0.10 M Co(NO3)2
  • ? mL 12 M HCl
  • ~1 mL bromothymol blue indicator solution
  • ~5 mL phosphate buffer of pH 6.3
  • ~5 mL phosphate buffer of pH 6.8
  • ~5 mL phosphate buffer of pH 7.3
  • deionized water

Safety

Concentrated hydrochloric acid (12 M HCl) is very corrosive, and its vapor is a respiratory irritant. Work with it under the fume hood at your lab bench, and avoid inhaling the vapor. Liquid hydrochloric acid can attack the skin and cause permanent damage to the eyes. If it splashes into your eyes, flush them in the eyewash station for at least 15 minutes; hold your eyes open or have someone assist you. If you spill the concentrated acid on your skin or clothing, flush the area immediately with water for at least 15 minutes. Have your lab partner notify your lab teaching assistant and the lab director about the spill. Silver solution will form dark spots on skin if spilled. The spots will not appear for about 24 hours, as the ions are slowly reduced to the metal. They are not hazardous, and will fade in a few days. Students will have access to and are encouraged to use gloves during the lab period owing to the use of 12M HCl.

Waste Disposal

Solutions from Part A and B of the experiment should be discarded in the waste container on the bench. You may wish to have a beaker in your work area to collect waste while you are doing the experiment. Make sure it is labeled. Use a squeeze bottle of deionized water to rinse the solutions into the beaker; use a minimum amount of water to avoid creating large volumes of waste solution. The plates and test tubes can then be washed in the normal manner. All of the solutions prepared in Part C of the experiment may be rinsed down the drain.

Prior to Class

Please complete WebAssign prelab assignment. Check your WebAssign Account for due dates. Students who do not complete the WebAssign prelab are required to bring and hand in the prelab worksheet.

PDF file

Lab Procedure

Please print the worksheet for this lab. You will need this sheet to record your data.

PDF file

Part A: Fe3+ + SCN- equilibrium arrow FeSCN2+ Equilibrium

Data Table A: Observations for the Equilibrium: Fe3+ + SCN-
equilibrium arrow 
FeSCN2+
Question 1: When Fe(NO3)3 was added to the system,
  • a
    Which ion in the equilibrium system caused the "stress"?
  • b
    Which way did the equilibrium shift?
  • c
    What happened to the concentration of SCN-?
  • d
    What happened to the concentration of FeSCN2+?
Question 2: When NaSCN was added to the system,
  • a
    Which ion in the equilibrium system caused the "stress"?
  • b
    Which way did the equilibrium shift?
  • c
    What happened to the concentration of Fe3+?
  • d
    What happened to the concentration of FeSCN2+?
Question 3: When AgNO3 was added to the system, it caused the precipitation of solid AgSCN.
  • a
    Which ion in the equilibrium had its concentration changed by addition of AgNO3?
  • b
    Did the concentration of that ion in solution increase or decrease?
  • c
    When AgNO3 was added, which way did the equilibrium shift?
Question 4: When you added NaNO3, did anything happen? Can you explain this result?

Part B: CoCl42- + 6 H2O equilibrium arrow Co(H2O)62+ + 4 Cl- Equilibrium

Data Table B: Observations for the Equilibrium: CoCl42- + 6 H2O equilibrium arrow Co(H2O)62+ + 4 Cl-
Question 5: Adding HCl has the effect of adding Cl- ions to the system. When Cl- was added to the system,
  • a
    Which way did the equilibrium shift?
  • b
    What happened to the concentration of CoCl42-?
  • c
    What happened to the concentration of Co(H2O)62+?
Question 6: When water was added to the system,
  • a
    Which way did the equilibrium shift?
  • b
    What happened to the concentration of CoCl42-?
  • c
    What happened to the concentration of Co(H2O)62+?
Question 7: When you added AgNO3, it caused the precipitation of solid AgCl.
  • a
    Which ion in the equilibrium had its concentration changed by addition of AgNO3?
  • b
    Did the concentration of that ion in solution increase or decrease?
  • c
    When AgNO3 was added, which way did the equilibrium shift?
Question 8: State a general rule concerning a system at equilibrium when more of one of the components is added.
Question 9: State a general rule concerning a system at equilibrium when one of the components is removed.
Question 10: For the CoCl42- + 6 H2O equilibrium arrow Co(H2O)62+ + 4 Cl- Equilibrium
  • a
    Which way did the equilibrium shift upon heating?
  • b
    Which way did the equilibrium shift upon cooling?
  • c
    A general rule concerning temperature changes to equilibrium systems is that the input of energy (raising the temperature) shifts the equilibrium to the higher energy side of the equilibrium. Based on your observations, which side of the equilibrium is the higher energy side?
  • d
    Is the reaction, CoCl42- + 6 H2O equilibrium arrow Co(H2O)62+ + 4 Cl- endothermic or exothermic?

Part C: Bromothymol Blue Equilibrium

Data Table C: Observations and Measurements for Bromothymol Blue Equilibrium
Question 11a: In the series from pH 6.30 to 6.80 to 7.30, the pH is increasing and the [H3O+] is decreasing. As the [H3O+] decreases, what happens to the concentration of BB represented by the absorbance at ~635 nm?
Question 11b: Explain how this observation agrees with Le Châtelier's principle.
Question 12a: As the [H3O+] decreases, what happens to the concentration of HBB represented by the absorbance at ~470 nm?
Question 12b: Explain how this observation agrees with Le Châtelier's principle.
Question 13: What is the equilibrium expression for the reaction under study? See Equation 6
HBB(aq) + H2O(l)equilibrium arrowBB + H3O+(aq)
yellowequilibrium arrowblue
Amax ~470 nmAmax ~635 nm
 
.