Equilibrium Lab PDF

Preview

Summary

Equilibrium Lab Report...

Description

Equilibrium Lab Report Introduction: The purpose of this experiment is to explore LeChatelier’s Principle and the effects of the addition of ions on reactions. In addition, we will see how the addition and/or removal of heat affects reactions at equilibrium.

Experimental: No deviations from standard procedure were made.

Data Analysis: Chromate-Dichromate Equilibrium 2CrO42-(aq) + 2H+(aq) ⇌ Cr2O72-(aq)+ H2O(l) ●

Initially, the left side is favored, as there are more yellow chromate ions in the solution than orange dichromate ions.

1. a. Adding H2SO4 caused the yellow solution to turn red upon contact. After mixing, the solution is orange. b. After the addition of NaOH, the solution turns back to its original yellow. 2. The addition of H2SO4 introduces H+ ions to the solution, causing the reaction to shift to the right, due to the common ion effect. This shift causes the colour to shift to orange, as more orange dichromate ions are produced. 3. OH- exerts an effect even though it doesn’t appear in the equation for the chromate-dichromate equilibrium because the OH- ions are able to react with H+ ions, forming H2O. As a result, the addition of OH- increases the concentration of products and causes the reaction to shift left, creating more chromate ions and changing the colour back to yellow. OH-(g) + H+ (g) = H2O(l) A Weak Acid HOAc(aq) + H2O(l) ⇌ H+ (aq)+ Ac(aq) ●

Initially, the right side of the reaction was favored because the methyl orange indicator turned red, indicating a pH of below 3.1. Therefore, there had to be a greater concentration of H+ ions in solution because the solution was acidic.

●

At first, the solution was a dark red orange colour due to the methyl orange indicator. Upon adding a few drops of sodium acetate, the solution turned a very light yellow/peachy colour.

●

The sodium acetate introduced the solution to acetate ions, causing there to be an increase in concentration of the products and a consequent shift to the left of the equilibrium. As a result, the concentration of hydrogen ions decreased and the pH increased to above 4.4 (as evidenced by the change in colour of the indicator to yellow).

A Weak Base NH3(aq) + H2O(l) ⇌ NH4(aq) + OH- (aq) ●

The solution contained no smell originally and was clear. Upon adding the phenolphthalein indicator, the colour of the solution turned a bright pink fuschia colour, indicating that the solution was originally basic. The solution also had no noticeable odor change. Given that the solution was basic, we can assume that the concentration of OH- ions was initially high and that the right side of the reaction was favored.

1. a. When NH3 is treated with NH4Cl, there was a lightening of colour in the solution and it eventually turned pink; there was no odor change. b. When NH3 is treated with HCl, the solution turns clear, and the smell of ammonia was present. 2. NH4Cl shifts the solution to the left by introducing more NH4 ions to the solution, causing the concentration of products to increase and forcing the reaction to shift left to overcome this disturbance. This lowered the pH of the solution below 8.0 (due to an increase in H+ ions due to the shift), a range at which the phenolphthalein indicator turns colourless. Thus, the solution lightened. On the other hand, HCl introduces H+ ions to the solution, which increases the concentration of products and also caused the reaction to shift left, producing ammonia and water. The decrease in pH similarly caused the solution to become colourless, however, it was to a greater extent because the HCl ions completely dissociate (as it is a strong acid). The ammonia smell was also caused by an increase in ammonia on the the reactants side caused by the shift, 3. Net ionic equation: NH3(aq) + H + (aq) ⇌ NH4+ (aq) Saturated Sodium Chloride NaCl(s) ⇌ Na+ (aq) + Cl- (aq) ●

There was no precipitant visible in the solution initially, this means that the sodium chloride must have been completely dissociated in the water, so the products side was favored initially.

1. When 12 M HCl is added to the saturated 5.4M NaCl, the solution turns a cloudy-white, with a white precipitate at the bottom. Once settled, there was a white precipitate at the bottom of a clear solution. This means that the reaction was shifted left, as solid NaCl was formed. 2. The equilibrium can be shifted by adding compounds that share common ions with the products. For example, aqueous KCl would dissociate into K+ and Cl- ions and cause the concentration of the products side to increase and the reaction to shift left.

Saturated Barium Chromate K2CrO4(aq) + BaCl2(aq) ⇌BaCrO4(s) ●

The initial reaction was as follows: BaCl2(s) ⇌ Ba 2+  (aq) + Cl - (aq). Given that there was no solute in the clear solution, therefore, the reaction initially favored the reactants.

●

Upon adding K2CrO4 to the solution, a cloudy yellow solution forms and there is some precipitate formed. This indicates that solid barium chromate was formed.

●

Hydrochloric acid is a strong acid, meaning it dissociates in water, forming H+ and Cl- ions. The H+ ions react with the CrO42- ions in the solution, causing the equilibrium to shift left because there’s less chromate on the reactants side of the reaction and accounting for the colour to change back to golden yellow.

1. Upon adding HCl, the precipitate dissolves and the solution turns a clear yellow colour. 2. We could shift the reaction by adding a compound with the same ions as the reactants or products (common ion effect). For example, if we added aqueous BaO, the additional Ba2+  ions introduced to the solution would increase the concentration of the reactants and shift the reaction to the right. Iron (III) Thiocyanate Solution Fe3+ (aq) + SCN - (aq) ⇌ Fe(SCN)2+  (aq) 1. Changes in the tubes were observed as follows: Tube 1 = Control Clear peachy orange solution was the initial condition for all four tubes. Tube 2 = FeCl3 added; solution turned dark orange and becomes cloudy. The FeCl3 introduced Fe3+  ions on the reactant side increased, and in order to respond to the stress, the equilibrium shifted to the right to produce additional Fe(SCN)2+  . The darker orange colour can be attributed to the additional red Fe(SCN)2+  produced.

Tube 3 = KSCN added; solution turned blood red colour. The concentration of thiocyanate ions on the reactant side increased, and in order to respond to the stress, the equilibrium shifted to the right to produce additional iron(III) thiocyanate. The blood red colour can be attributed to the additional red Fe(SCN)2+  produced. Tube 4 = AgNO3 added; solution lost colour then turned slightly cloudy, indicating that a precipitate formed. The AgNO3 introduced additional Ag+ ions, which bonded with the thiocyanate ions in the solution to form the precipitate, silver thiocyanate. The concentration of thiocyanate ions on the reactant side decreased, and to respond to the stress, the equilibrium shifted to the left and produced additional thiocyanate and iron(III) ions. The colour loss can be attributed the the dissociation of the iron (III) thiocyanate, which caused the red colour. The cloudy colour can be attributed to the formation of silver thiocyanate.

Equilibrium of a Cobalt Complex 2+ - CoCl4 2+  (aq) + 6H2O(l) ⇌ Co(H2O)6  (aq) + 4Cl  (aq)

1. Temperature Effect Add Heat

The addition of heat caused the solution to turn a dark blue. Given that the reaction is exothermic, the addition of heat shifts the reaction to the left, creating more cobalt chloride, the source of additional blue color.

Remove Heat

The removal of heat caused the solution to turn light purple and then clearish-light pink. This is because when heat is removed from an exothermic reaction, the reaction shifts right, creating more colorless hydrous cobalt.

●

Cobalt chloride is a magenta salt in the solid state. Adding ethanol causes the cobalt chloride to turn into a dark blue aqueous solution.

2. Common Ion Reaction Observation Tube #

1

Added

Added

H2O

HCl

After 2 drops of water, the

The HCl caused the solution to

solution reaches mid-blue.

turn dark blue.

2

H2O

CaCl2

After 2 drops of water, the

CaCl2 dropped to the bottom and

solution reaches mid-blue.

bubbled. Solution was dark blue with a precipitate.

3

H2O

AgNO3

After 2 drops of water, the

The AgNO3 caused the solution

solution reaches mid-blue.

to turn pink, and a precipitate was formed.

4

Nothing

Nothing

Control; blue

Control; blue

The solution reaches mid-blue

The solution reaches mid-blue

after 2 drops.

after 2 drops.

3. LeChatelier’s Principle ●

HCl = The concentration of chloride ions on the product side increased and to relieve stress on that side of the reaction the equilibrium shifted to the left. The blue color change can be attributed to the increase in CoCl42- being produced which came from the equilibrium shift.

●

CaCl2 = The concentration of chloride ions on the product side increased, and by the common ion effect, equilibrium shifted to the left. The blue color change can be attributed to the increase in CoCl4 2- production, which came from the equilibrium shift.

●

AgNO3 = The silver and chloride ions already present in the solution formed silver chloride precipitate. The concentration of chloride ions on the product side decreased, and the system shifted to the right. The cloudy color is due to the silver chloride precipitate.

●

Add heat: The reaction is exothermic, so the addition of heat to the reaction at equilibrium will shift the reaction left. This shift created the blue color because more CoCl42- was formed.

●

Remove heat: The reaction is exothermic, so the removal of heat to the reaction at equilibrium will shift the reaction right. This shift created the light blue (and then pink) color because more Co(H2O)62+ was formed.

Conclusion: This experiment allowed us to explore the common ion effect and the effects of heat that are predicted for a reaction at equilibrium by LeChatelier’s Principle. It was observed that, through the common ion effect, if the concentration of products/reactants is increased or decreased, the reaction will shift accordingly to reduce the stress created by the change. The addition of heat shifts reactions to the right for endothermic reactions and to the left for exothermic reactions. The removal of heat causes the opposite effects. Our observations were consistent with LeChatelier’s Principles and, as a result, this experiment is successful. There were no calculations in this lab, thus, there are no error calculations....