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Chemical Equilibrium and Le Chatelier’s Principle
Objectives
The objective of this lab is to observe the effect of an applied stress on chemical systems at equilibrium.
Background
A reversible reaction is a reaction in which both the conversion of reactants to products (forward reaction) and the re-conversion of products to reactants (backward reaction) occur simultaneously:
forward reaction A + B → C + D Reactants Products
backward reaction C + D → A + B Products Reactants
reversible reaction A + B ' C + D
Consider the case of a reversible reaction in which a concentrated mixture of only A and B is supplied. Initially the forward reaction rate (A + B → C + D) is fast since the reactant concentration is high. However as the reaction proceeds, the concentrations of A and B will decrease. Thus over time the forward reaction slows down. On the other hand, as the reaction proceeds, the concentrations of C and D are increasing. Thus although initially slow, the backward reaction rate (C + D → A + B) will speed up over time. Eventually a point will be reached where the rate of the forward reaction will be equal to the rate of the backward reaction. When this occurs, a state of chemical equilibrium is said to exist. Chemical equilibrium is a dynamic state. At equilibrium both the forward and backward reactions are still occurring, but the concentrations of A, B C and D remain constant.
A reversible reaction at equilibrium can be disturbed if a stress is applied to it. Examples of stresses include increasing or decreasing chemical concentrations, or temperature changes. If such a stress is applied, the reversible reaction will undergo a shift in order to re-establish its equilibrium. This is known as Le Chatelier’s Principle.
Consider a hypothetical reversible reaction already at equilibrium: A + B ' C + D. If, for example, the concentration of A is increased, the system would no longer be at equilibrium. The rate of the forward reaction (A + B → C + D) would briefly increase in order to reduce the amount of A present and would cause the system to undergo a net shift to the right. Eventually the forward reaction would slow down and the forward and backward reaction rates become equal again as the system returns to a state of equilibrium. Using similar logic, the following changes in concentration are expected to cause the following shifts:
Increasing the concentration of A or B causes a shift to the right. Increasing the concentration of C or D causes a shift to the left. Decreasing the concentration of A or B causes a shift to the left. Decreasing the concentration of C or D causes a shift to the right.
In other words, if a chemical is added to a reversible reaction at equilibrium, a shift away from the added chemical occurs. When a chemical is removed from a reversible reaction at equilibrium, a shift towards the removed chemical occurs.
A change in temperature will also cause a reversible reaction at equilibrium to undergo a shift. The direction of the shift largely depends on whether the reaction is exothermic or endothermic. In exothermic reactions, heat energy is released and can thus be considered a product. In endothermic reactions, heat energy is absorbed and thus can be considered a reactant.
exothermic A + B ' C + D + heat
endothermic A + B + heat ' C + D
As a general rule, if the temperature is increased, a shift away from the side of the equation with “heat” occurs. If the temperature is decreased, a shift towards the side of the equation with “heat” occurs.
In this lab, the effect of applying stresses to a variety of chemical systems at equilibrium will be explored. The equilibrium systems to be studied are given below:
Saturated Sodium Chloride Solution NaCl ( s ) ' Na+1^ ( aq ) + Cl -1^ ( aq )
Acidified Chromate Solution 2 CrO 4 -2^ ( aq ) + 2 H+1^ ( aq ) ' Cr 2 O 7 -2^ ( aq ) + H 2 O ( l ) yellow orange
Aqueous Ammonia Solution (with phenolphthalein) NH 3 ( aq ) + H 2 O ( l ) ' NH 4 +1^ ( aq ) + OH-1^ ( aq ) clear pink
Cobalt(II) Chloride Solution Co(H 2 O) 6 +2^ ( aq ) + 4 Cl -1^ ( aq ) ' CoCl 4 -2^ ( aq ) + 6 H 2 O ( l ) pink blue
Iron(III) Thiocyanate Solution Fe +3^ ( aq ) + SCN-1^ ( aq ) ' Fe(SCN) +2^ ( aq ) pale yellow colorless deep red
By observing the changes that occur (color changes, precipitate formation, etc.) the direction of a particular shift may be determined. Such shifts may then be explained by carefully examining the effect of the applied stress as dictated by Le Chatelier’s Principle.
Part 4: Cobalt(II) Chloride Solution
a. Place 3-mL of 0.1M CoCl 2 ( aq ) into 3 small test tubes. Label these test tubes 1-3. b. The solution in test tube #1 remains untouched. It is a control for comparison with other tubes. c. To the solution in test tube #2, carefully add concentrated 12M HCl ( aq ) drop-wise until a distinct color change occurs. Record your observations. d. To the solution in test tube #3, first add a medium scoop of solid NH 4 Cl. Then heat this solution directly in your Bunsen burner flame (moderate temperature). Firmly hold test tube #3 with your test tube holder, and waft it back and forth through the flame (to prevent overheating and “bumping”) for about 30 seconds, or, until a distinct change occurs. Record your observations. Then cool the solution in test tube #3 back to room temperature by holding it under running tap water, and again record your observations.
Part 5: Iron(III) Thiocyanate Solution
Instructor Prep : At the beginning of lab prepare a stock solution of iron(III) thiocyanate. Add 1-mL of 0.1M FeCl 3 ( aq ) and 1-mL of 0.1M KSCN ( aq ) to a 150-mL (medium) beaker, top it up with 100-mL of distilled water, and mix with a stirring rod. Label the beaker and place it on the front desk. The entire class will then use this stock solution in Part 5.
a. Place 3-mL of the prepared stock solution into 4 small test tubes. Label these test tubes 1-4. b. The solution in test tube #1 remains untouched. It is a control for comparison with other tubes. c. To the solution in test tube #2, add 1-mL of 0.1M FeCl 3 ( aq ). Record your observations. d. To the solution in test tube #3, add 1-mL of 0.1M KSCN ( aq ). Record your observations. e. To the solution in test tube #4, add 0.1M AgNO 3 ( aq ) drop-wise until all the color disappears. A light precipitate may also appear. Record your observations. Here the added silver nitrate is effectively removing thiocyanate ions from the equilibrium system via a precipitation reaction: Ag +1^ ( aq ) + SCN-1^ ( aq ) → AgSCN ( s ).
