Purpose

To evaluate the effect of temperature and addition of reactants on the position of equilibrium in chemical reactions in equilibrium.

Expected Learning Outcomes

• Determine the equilibrium constant for a reaction.
• Observe the shift in position of equilibrium when concentrations or temperature are changed in a reacting mixture.
• Explain observations of the position of equilibrium using Le Châtelier’s Principle

Textbook Reference

The content in this experiment spans all of Tro, Chemistry: Structures and Properties, 2nd Ed., Ch. 15.

Introduction

Chemical Equilibrium: Reversible Reactions and Dynamic Equilibrium

In CHEM-C 105/125, the focus (in terms of chemical reactions) was on reactions where the reaction goes to completion. In this case, at the end of the reaction,^[1] we expect all of the limiting reactant to react to form product, with no limiting reactant left; the reverse reaction will not occur.

However, in many cases, the reaction can occur in both directions. Consider the hypothetical reaction:

[latex]\ce{A} \rightleftharpoons \ce{B}[/latex]

If I start with pure A, then clearly no reverse reaction occurs; only the forward reaction occurs, and A is converted to B. However, as this proceeds, as [A] decreases, the rate of the forward reaction decreases. Conversely, the reverse reaction ([latex]\ce{B}\to\ce{A}[/latex]) occurs at an increasing rate as [B] increases..  Eventually, as the rates converge, the forward reaction rate will equal the reverse reaction rate. At that point, the macroscopic concentrations of reactants and products will remain constant. However, it is a dynamic equilibrium where the forward and reverse rates are the same, but not non-zero.

Equilibrium Constants

In order to describe the position of a reaction equilibrium – the extent to which the forward reaction proceeds to completion – we express the relative concentrations at equilibrium in terms of the equilibrium constant [latex]K[/latex]. For this experiment, we will focus on the equilibrium constant in terms of concentrations (molarity) [latex]K_C[/latex].

In this case, we express the equilibrium constant for a reaction such that we multiply together the concentration terms for the products divided by those for the reactants.  If there is a coefficient associated with a given reactant/product, the stoichiometric coefficient should become the power to which the concentration term is raised.

Note that:

• The equilibrium constant depends on temperature but not concentrations of reactants and products.
• The equilibrium constant should be unitless.^[2]
• When writing down the expression for the equilibrium constant, pure liquids and solids (including solvents) do not appear in the equilibrium constant expression.

Le Châtelier’s Principle

We can use Le Châtelier’s Principle to predict the direction to which a system will shift when a change is made to a system at equilibrium:

That is to say, if you add or remove something from the system, the reacting mixture shifts its position to “undo” whatever change you have made.  In this lab manual, we will focus on how the equilibrium position is affected by the change of concentration and temperature; see your textbook for more examples of how Le Châtelier’s Principle can be applied.

Effect of Changing Concentration on Position of Equilibrium

When the concentration of a species (reactant or product) is decreased, the system shifts its position to increase the concentration of the species whose concentration was decreased. Conversely, when a species is added, the reaction position shifts to the other side as that specific species is reacted.

We can account for this by calculating the reaction quotient^[3] for this reaction and showing the direction to which the equilibrium will shift. The equilibrium constant will remain constant in this case. Using an ICE/equilibrium table, you can calculate the final, resultant concentrations at equilibrium.

Effect of Changing Temperature on the Position of equilibrium

To think about this, we can use the idea that increasing the temperature is, in a pictoral sense, “adding” heat; decreasing the temperature would be “removing” heat. In turn, depending on whether the forward reaction is endothermic or exothermic, we can determine whether heat is a reactant or product:

• In an endothermic reaction ([latex]\Delta_r H > 0[/latex]), we can treat it as requiring heat to drive a reaction forward and hence heat should be treated as a reactant in the equilibrium reaction.
• Conversely, in an exothermic reaction ([latex]\Delta_r H < 0[/latex]), heat would be evolved from the reacting mixture into the system.  Heat can therefore be treated as a product of this reaction.

The easiest way to look at this would be to write down heat explicitly into the equation as a reactant or product, and analyze this similarly to adding/removing a particular reactant/product from a reacting mixture.

Examples

The reaction

\begin{equation}
2\ce{NO2}(g) \rightleftharpoons \ce{N2O4}(g)
\end{equation}

is endothermic. Therefore, heat can be written in to the equation on the reactants side.

\begin{equation}
2\ce{NO2}(g) + \textrm{heat} \rightleftharpoons \ce{N2O4}(g)
\end{equation}

When the temperature increases, you are adding heat to the system.  Given the equation above, you are effectively adding reactant to the reacting mixture, which would shift the position of equilibrium to the right.  In this example, the equilibrium position shifts to the right – more [latex]\ce{N2O4}[/latex] is produced from [latex]\ce{NO2}[/latex] as illustrated below:

Temperature changes will alter the value of the equilibrium. In the experiment Thermodynamics of the Solvation of Calcium Hydroxide, you will study how the equilibrium constant changes as a function of temperature. This is explained in your textbook as well.

Reactions Studied

In this experiment, you will study the shift in equilibrium for two different chemical reactions:

Antimony Trichloride

Antimony trichloride reacts with water to form SbOCl, a white solid.

\begin{equation}
\ce{SbCl3} (aq) + \ce{H2O} (l) \rightleftharpoons \ce{SbOCl}(s) + 2\ce{H+} (aq) + 2\ce{Cl-}(aq)
\end{equation}

Note that antimony trichloride will normally not dissolve in water; it will be supplied to you in a solution containing 0.50 M antimony trichloride in 6.0 M HCl.

Tetrachlorocobaltate (II)

In the presence of chloride ions, cobalt (II) undergoes the following reaction to form tetrachlorocobaltate (II):

\begin{equation}
\underset{\mbox{pink}}{\ce{Co(H2O)}_6^{2+}} (aq)+ 4\ce{Cl-}(aq) \rightleftharpoons \underset{blue}{\ce{CoCl}_4^{2-}} (aq) + 6\ce{H2O}(l)
\end{equation}

In this experiment, you will study how the position of equilibrium shifts as the following are changed:

• Concentration of Cl^–
• Temperature

Procedures

Antimony Trichloride

As mentioned above, the antimony trichloride is dissolved in 6 M HCl. The initial solution contains 0.50 M antimony trichloride and 6 M HCl

Qualitative Observation of Equilibrium Shifts

For this part of the experiment, we will use microscale techniques with a disposable plastic dropper. We will assume that each drop has approximately the same volume.^[4].  The solution will be placed in a small test tube.

1. In a small test tube, place five drops of 0.50 M SbCl₃ in 6.0 M HCl.  Describe its appearance.
2. Add four drops of deionized water to the test tube. Using a clean pipet, ensure that the sample is mixed well together. Record your observations.
3. Repeat step 2 three more times.
4. Add two drops of 6 M hydrochloric acid to the test tube. Mix and record your observations.

Determining [latex]K_C[/latex] of this reaction

1. Pipet 5.00 mL of 0.50 M SbCl₃ in 6.0 M HCl into a clean, dry, small Erlenmeyer flask. Do not add additional water to the Erlenmeyer flask.
2. Fill the buret with deionized water and read the buret to the correct number of significant figures.
3. Titrate the contents of the Erlenmeyer flask with deionized water from the buret until there is a speck of undissolved white SbOCl that persists. Read the buret to the correct number of significant figures at this point.

Calculation

Considering the reaction

\begin{equation}
\ce{SbCl3} (aq) + \ce{H2O} (l) \rightleftharpoons \ce{SbOCl}(s) + 2\ce{H+} (aq) + 2\ce{Cl-}(aq)
\end{equation}

the reverse reaction cannot occur until there is some SbOCl accumulated in the solution.  However, when there is only an infinitesimally small amount of SbOCl present in the solution, an equilibrium is maintained.

We can, in this case, assume that minimal amounts of SbCl₃ has reacted to form SbOCl; therefore, the concentrations of all species in this reaction can be calculated by only considering the dilution of the initial solution:

\begin{equation}
M_1 V_1 = M_2 V_2
\end{equation}

Using the initial concentrations of Sb³⁺, H+ and Cl^– from the antimony trichloride solution (0.50 M for SbCl₃ and 6.0 M for both H⁺ and Cl^–), you can find the equilibrium concentrations of the species required to substitute into the equilibrium constant expression for this reaction.

Formation of Tetrachlorocobaltate (II)

Effect of Change of Concentration on the Equilibrium

1. You are provided with a 0.4 M aqueous Co(NO₃)₂ solution. Add five drops of this solution into a clean, dry small test tube. Write down the color of the solution immediately after addition and after mixing.
2. Add two drops of concentrated hydrochloric acid into the test tube and write down the color of the solution immediately after addition as well as after mixing. Repeat this seven times to give a final volume of 16 drops of concentrated (12 M) HCl.
3. Add four of deionized water into the test tube that you have used previously and write down the color of the solution immediately after addition as well as after mixing. Repeat this four times until 16 drops of deionized water has been added.

Effect of Temperature on Equilibrium

1. Boil a 400 mL beaker of water on a hot plate.
2. In a small beaker, mix together 5 mL of 0.4 M Co(NO₃)₂ with 6 mL concentrated hydrochloric acid. Split the contents of this solution into three approximately equal portions in three different, small test tubes.
3. Place one of the test tubes in an ice bath, and another test tube into the boiling water bath. Observe the colors change in each of the test tubes and record them.
4. Swap around the test tubes in the boiling and ice baths. Observe any changes that occur.

Waste Disposal

All chemicals in this experiment must be disposed of in the designated waste beaker.