Chemical equilibrium and le chatelier's principle lab report answers

Equilibrium and Le Châtelier’s Principle Hands-On Labs, Inc. Version 42-0166-00-02

Review the safety materials and wear goggles when working with chemicals. Read the entire exercise before you begin. Take time to organize the materials you will need and set aside a safe work space in which to complete the exercise.

Experiment Summary:

In this experiment, you will describe the components of a reaction at chemical equilibrium and use Le Châtelier’s principle to predict the direction a chemical system will shift upon changes in concentration, temperature, and pressure. You will perform equilibrium reactions and test Le Châtelier’s principle by manipulating concentration and temperature of the reactions.

EXPERIMENT

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Learning Objectives Upon completion of this laboratory, you will be able to:

● Define chemical equilibrium and chemical system.

● Define equilibrium constant and reaction quotient.

● Explain how changes in temperature, pressure, and concentration affect the chemical system of a reaction.

● State Le Châtelier’s Principle.

● Perform chemical equilibrium reactions and manipulate chemical systems through concentration and temperature.

● Perform calculations to determine the equilibrium constant (K) and reaction quotient (Q) of reactions.

● Apply Le Châtelier’s principle to predict changes and explain observed changes in a chemical system.

Time Allocation: 2.5 hours

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Experiment Equilibrium and Le Châtelier’s Principle

Materials Student Supplied Materials

Quantity Item Description 1 Dish soap 1 Hot water bath (hot water, cup) 1 Ice bath (ice, water, cup) 1 Pair of scissors, 4 in 1 Roll of paper towels 1 Source of tap water 1 Tape: clear, duct, or masking

HOL Supplied Materials

Quantity Item Description 2 Pairs of gloves 1 Pair of safety goggles 1 Test tube cleaning brush 1 Well plate – 24 1 Experiment Bag: Equilibrium and Le Châtelier’s Principle:

1 - Potassium chromate (K2CrO4), 1 M, 2 mL in pipet 1 - Potassium ferrocyanide (K4Fe(CN)6), 0.2 M, 2 mL in pipet 1 - Iron(III) nitrate (Fe(NO3)3), 0.1 M, 2 mL in pipet 1 - Hydrochloric acid (HCl), 2 M, 10 mL in dropper bottle 1 - Sodium hydroxide (NaOH), 2 M, 10 mL in dropper bottle 3 - Short, thin-stem pipets

Note: To fully and accurately complete all lab exercises, you will need access to:

1. A computer to upload digital camera images.

2. Basic photo editing software, such as Microsoft Word® or PowerPoint®, to add labels, leader lines, or text to digital photos.

3. Subject-specific textbook or appropriate reference resources from lecture content or other suggested resources.

Note: The packaging and/or materials in this LabPaq kit may differ slightly from that which is listed above. For an exact listing of materials, refer to the Contents List included in your LabPaq kit.

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Experiment Equilibrium and Le Châtelier’s Principle

Background Equilibrium

Chemical reactions may go to completion, that is, the reactants produce products until the one of the reactants (the limiting reactant) is used up. However, most reactions are reversible and products react to reform the reactants. The reactants, products, and energy associated with a chemical reaction are referred to as a chemical system. The reaction reaches a state of chemical equilibrium when the rate of the forward reaction is the same as the rate of the reverse reaction, which means that the concentrations of the reactants and products are constant. The term dynamic equilibrium is used to emphasize that the reactions are still occurring, even though the reaction appears to have stopped changing. See Figure 1.

Figure 1. (Top) This reaction proceeds to completion, as noted by the pink arrow. (Bottom) This reaction is reversible, as noted by the blue arrows which move in two directions.

The situation at equilibrium varies greatly from reaction to reaction. Some chemical reactions reach equilibrium with mainly reactants present, other reactions reach equilibrium with appreciable amounts of both reactants and products present, while still others do not reach equilibrium until mostly products are present. In addition, it makes no difference if a reaction is started with 100% reactants or with 100% products, or whether a reaction is started with 1 mole of reactants or 10 moles of products, it will reach the same equilibrium point where free energy is equal to zero. See Figure 2.

Figure 2. Reaction at chemical equilibrium. In this reaction, the reactants (9 green balls and 9 pink balls) react to form product (9 purple balls). As the reaction is at chemical equilibrium, the reaction does not go to 100% completion. Rather, at chemical equilibrium, there is a mixture of

both products and reactants: 8 purple balls, 1 green ball, and 1 pink ball.

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Experiment Equilibrium and Le Châtelier’s Principle

Equilibrium Constant

When a reaction is at equilibrium (and at a specific temperature), the concentrations of reactants and product remain constant, which makes it possible to define an equilibrium constant, K. (See below) The equilibrium constant, K, is equal to the molar concentrations of the products, each raised to a power equal to the coefficient in the balanced chemical equation, divided by the molar concentrations of the reactants, each raised to a power equal to the coefficient in the balanced chemical equation.

If the products concentrations are much higher than the reactant concentration at equilibrium, then K will be a large number. K values close to 1 mean that both reactants and products are present in similar amounts at equilibrium. Small K values (less than 1) indicate that reactant concentrations are higher than product concentrations at equilibrium.

Generally, if the value of K is greater than 1, we say that the reaction favors the products and if the value of K is less than 1, the reaction favors the reactants. The larger the value of K, the more the reaction favors the products. For example, consider the sample reaction and equilibrium constant calculations below:

The concentrations of the reactants and products at equilibrium are: [N2] = 1.4 x 10 -3 M, [Cl2] =

4.3 x 10-4 M, and [NCl3] = 1.9 x 10 -1 M. The equilibrium constant is calculated as shown below:

As the value of K is much larger than 1, the reaction favors the formation of product. See Figure 3.

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Experiment Equilibrium and Le Châtelier’s Principle

Figure 3. Relationship between K and the formation of products vs. reactants.

Le Châtelier’s Principle

Once a reaction is at equilibrium, any change in concentration (or pressure of a gaseous reactant or product) will disrupt the chemical system. The reaction will spontaneously return to equilibrium, but at a slightly different position than the original one. For example, adding more reactant to a system at equilibrium will speed up the forward reaction and produce more products. The resulting increase in product concentration will cause the reverse reaction to speed up also until it equals the forward rate, re-establishing the equilibrium at a higher product concentration than before. We say that this reaction “shifted to the right” to again achieve equilibrium. Adding product to a reaction at equilibrium results in the opposite situation. Le Châtelier’s principle states that a chemical system will adjust in response to in order change to return to equilibrium. See Table 1.

Table 1. Chemical system responses to change.

Type of Change Chemical System Shift

(Right shifts toward product. Left shifts toward reactant.)

Increase concentration of reactant OR

Decrease concentration of product Right

Decrease concentration of reactant OR

Increase concentration of product Left

Increase temperature of an exothermic reaction Left Increase temperature of an endothermic reaction Right Decrease temperature of an exothermic reaction Right

Decrease Temperature of an endothermic reaction Left Decrease pressure (gases only) More gas molecules Increase Pressure (gases only) Less gas molecules

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Experiment Equilibrium and Le Châtelier’s Principle

Reaction Quotient

The value of the equilibrium constant for a reaction varies with temperature. In a chemical system at equilibrium, if the forward reaction is exothermic, then the reverse reaction must be endothermic, and vice versa. Increases in temperature cause both reactions to speed up but will have a greater effect on the endothermic reaction. Increases in temperature are said to “favor” the endothermic process while decreases in temperature “favor” the exothermic process. See example below:

If additional heat were applied to the exothermic reaction, the system would shift to the left, reducing the overall temperature of the reaction. Likewise, if additional AB3 was added to the reaction, the system would again shift to the left to create the appropriate balance of products and reactants.

The chemical system can also be shifted toward the right upon addition of A2 or B2, as both additions would cause an increased concentration of product. Likewise, as both A2 and B2 are gaseous substances, increasing or decreasing the pressure of these two gases would cause a shift of the chemical system to re-achieve the equilibrium position.

The above examples discussed predicting how a reaction at equilibrium will respond to changes. It is also possible to predict how a newly combined mixture of reactants and products will proceed to reach equilibrium. We do this by determining the reaction quotient, Q, and comparing it to the value of the equilibrium constant. Q is calculated the same way as K, but using initial concentrations rather than equilibrium concentrations. See Table 2.

Table 2. Relationship between reaction quotient (Q) and equilibrium constant (K).

Relationship Chemical System Response Q is equal to K System remains at equilibrium

Q is greater than K The reverse reaction will predominate as the system approaches equilibrium.

Q is less than K The forward reaction will predominate as the system approaches equilibrium.

Example:

For example, examine the reaction for the synthesis of ammonia, which has a K = 6.0 x 10-2. Note that the initial concentrations are differentiated from equilibrium concentrations by the subscripted “0.”