CHEMICAL EQUILIBRIUM LABORATORY 3

GEORGIA MILITARY COLLEGE

NATURAL SCIENCE DEPARTMENT

ONLINE CAMPUS

LABORATORY 3 - CHEMICAL EQUILIBRIUM

NAME

STUDENT NUMBER

CLASS

PROFESSOR’S TITLE AND NAME

Introduction

Generally, equilibrium is a state of balance between opposing forces. In chemical reactions, equilibrium is achieved when the concentrations of the products and reactants are in balance, so no further changes are observed in the system (Smith, 2024). Chemical equilibrium plays a fundamental role in industrial processes as well as in human life. Why is it important to understand the concept of equilibrium in industrial processes? This question is best answered using the Le Chatelier’s principle, which states that a change in one of the elements of a system in dynamic equilibrium will trigger a shift in the equilibrium position in an attempt to counter the change and reestablish equilibrium (Smith, 2024). Factors that can cause changes to a system in equilibrium include reactant concentrations, temperature, and pressure (Smith, 2024). According to Le Chatelier’s principle, an increase or decrease in any of these factors will trigger a shift in the equilibrium point in the opposite direction (Smith, 2024). Using this knowledge, industrial chemists can adequately manipulate chemical reactions to increase or decrease the production of certain products.

This laboratory uses the reaction between Cobalt (II) and chloride ions to observe how the equilibrium point changes due to changes in temperature and concentration of reagents. Figure 1 below presents the chemical equation for the above reaction:

[Co (H2O)6]+2 + 4Cl- ? CoCl4-2 +6H2O

Figure 1: [Co (H2O)6]+2 forms a pink complex, while Co Cl 4-2 is a blue complex.

Combined with six water molecules, Cobalt (II) forms a pink complex that turns into a blue complex, CoCl42 upon reacting with chloride ions. This laboratory seeks to realize three objectives:

i) To enhance the ability to apply Le Chatelier’s principle.

ii) To enhance understanding of the equilibrium constant concept.

iii) To analyze the effects of changes in temperature or concentration on the equilibrium constant.

From figure 1 above, the equilibrium constant (k) for the reaction is given by:

The general hypotheses established at the start of the lab were:

i) Addition of chloride ions will increase the reactants above the products in the equilibrium constant, thus shifting the equilibrium in figure 1 to the right (the products side), leading to more production of CoCl42, and causing the solution to turn blue.

ii) Addition of distilled water to the reaction in figure 1 will decrease the concentration on the products side of the constant. Thus, there will be more reactants than products, causing the equilibrium to shift to the left (the reactants side), leading to more production of [Co (H2O)6]+2, which causes the solution to turn pink

iii) Addition of silver Nitrate (AgNo3) will reduce chloride ions (Cl-) due to the formation of AgCl. This will imply that there will be more products than reactants and the equilibrium will shift to the right (towards the products) to produce more chloride ions. Consequently, the solution turns blue.

Materials and Methods

Preparing the Lab

i) Select ‘Virtual Lab’ on the home page of the course to load the lab environment.

ii) Wait for the lab environment to load, then select ‘File’, and subsequently, ‘Load an Assignment’.

iii) Choose the category labeled ‘Chemical Equilibrium’ and then the assignment titled ‘Cobalt Lab.’ At this point, the lab preparations are complete and one is ready to perform the experiment.

Performing the Experiment

iv) In the stockroom, choose ‘Glassware’ and then ‘Empty 1000mL Erlenmeyer flask.’

v) Return to the stockroom and in the ‘Solutions’ tab, select ‘Cobalt (II) Chloride Exp’ solutions. Move to the work area the flask containing 1M Cobalt (II) Chloride solution.

vi) Return to the ‘Solutions’ tab, select HCL, and move the flask containing 12M HCL to the work area.

vii) Return to the ‘Solutions’ tab, select AgNO3, and move the flask containing 6M AgNO3 to the work area.

viii) Return to the ‘Solutions’ tab, select distilled water and move the respective flask to the work area.

ix) Add 25mL of 1M cobalt (II) chloride into the empty Erlenmeyer flask by dropping the flask containing the solution into the empty one, and record the solution’s color as well as molarity in table 1.

x) Add 2mL 12M HCL into the flask containing Cobalt (II) Chloride, and keep adding HCL until there is an observable change in the mixture. Record the color of the resulting solution in table 1 when no further change can be seen.

xi) Add 2mL distilled water to the flask containing the mixture, and keep adding until the mixture changes color and no further changes in color occur. Record the observations in table 1.

xii) Repeat steps (iv) to (x), and then add 20Ml of 6M AgNO3 to the solution and record the observations in table 2.

xiii) Finally, clear the work area, but remember to keep the virtual lab open.

Data Analysis

i) Explain the observations in each step in line with Le Chatelier’s principle.

ii) Using the Le Chatelier’s principle, explain what is likely to happen when the pink equilibrium mixture is placed in hot water. Is the reaction exothermic or endothermic?

Data and Discussion

Results

Table 1. Recorded Observations with Addition of 12M HCL and Distilled Water

Observations

Step (ix): Addition of 25mL1M cobalt (II) chloride into empty Erlenmeyer flask (Observations at equilibrium)

Step (x): Addition of 2mL 12M HCL into the Solution

(Observations at new equilibrium)

Step (xi): Addition of 2mL distilled water into the solution (Observations at equilibrium)

H+ (Molarity)

9.843e-8

9.843e-8

OH- (Molarity)

9.843e-8

9.843e-8

Co (H2O)6 (Molarity)

CL- (Molarity)

CoCl4-2 (Molarity)

Temperature

25 degrees

24.12 degrees

25 degrees

Color

Pink

Blue

Pink

pH

Volume

25mL

31mL

75mL

Precipitate

None

None

None

Table 2. Observations with Addition of 6M AgNO3

Observations

Step (xii): Addition of 2mL of 6M AgNO3

H+ (Molarity)

OH- (Molarity)

Co (H2O)6 (Molarity)

CL- (Molarity)

CoCl4-2 (Molarity)

AG+

9.50908e-11

NO3-

Temperature

25 degrees

Color

Pink and cloudy

pH

Volume

436.27mL

Precipitate

White precipitate present

Discussion

The starting solution, cobalt (II) chloride exists as a pink complex at room temperature. Addition of concentrated (12M) HCL increases the concentration of chloride (Cl-) ions in the solution from 1.866 to 3.23219 as shown in table 1. According to Le Chatelier, the equilibrium position will shift to counter the change and reestablish equilibrium (Smith, 2024). In this case, the equilibrium position will shift to reduce the concentration of Cl- by reacting it with [Co (H2O)6]+2 to form more CoCl4-2 and H2O. Therefore, the equilibrium position shifts to the right (products side) as more CoCl4-2 is produced, causing the solution to turn blue. When distilled water is added, the reverse happens. Addition of distilled water reduces the concentration of CoCl4-2 from 0.175822 to 0.0334 as shown in table 1. In line with Le Chatelier, this destabilizes the equilibrium and the equilibrium point has to shift to increase the concentration of CoCl4-2 again (Smith, 2024). As a result, more [Co (H2O)6]+2 and Cl- will react to produce more CoCl42. The equilibrium point thus shifts to the left (products side), causing the solution to turn pink.

Addition of AgNO3 to the blue equilibrium solution removes Cl- ions from the solution as Ag+ reacts with Cl- to form Silver chloride (AgCl), which appears as a white precipitate at the bottom of the flask. The removal of Cl- is evidenced by the reduced concentration of Cl- in table 2 (molarity =1.95961) compared to the Cl- molarity in the equilibrium blue solution (molarity = 3.23219) in table 1. According to Le Chatelier, the removal of Cl- induces a shift in the equilibrium point as the system seeks to increase the concentration of Cl- by getting more CoCl42 and H2O to react to produce Cl-. Consequently, the equilibrium point shifts to the left, causing the solution to turn pink. The cloudiness in the reaction is due to the production of hydrogen chloride gas.

When the pink equilibrium mixture is placed in hot water, the solution turns blue. The temperature increase causes the equilibrium point to shift to the right, which means that the reaction is endothermic (absorbs heat) (Smith, 2024). The Le Chatelier principle holds that the temperature increase will trigger a shift in the equilibrium point as the system seeks to reduce the temperature to reestablish equilibrium. To do this, the reaction has to absorb the extra heat caused by the temperature increase (Smith, 2024). Thus, the equilibrium shifts to the right, producing more CoCl42, which leads the solution to turn blue. The accuracy of the procedure could be increased by ensuring all instruments are properly calibrated and conducting parallel experiments to help identify anomalies in the collected data.