Experiment 4 Rdr – Chemical Equilibrium Essay
In the first phase of the experiment, iron sulfate was mixed with silver nitrate, and the reaction produced solid silver and iron nitrate, which is formally written as,
Fe2+ (aq) + Ag+ (aq) ⇄ Ag (s) + Fe3+ (aq)
The mixture was then placed in a centrifuge in order for the solid silver to settle and separate from the supernate. The supernate was then tested for the presence of Fe2+, Fe3+ and Ag+ by placing K3Fe(CN)6, KSCN and HCl to 3 separate samples of the supernate. The result of the tests was as follows:
1.Addition of K3Fe(CN)6
After K3Fe(CN)6 was added to the supernate, a Prussian blue precipitate was formed, more formally written as,
Fe2+ (aq) + Fe(CN)63- (aq) + K+ (aq) → K∙Fe2(CN)6 (s)
This reaction thus proves that the supernate contains Fe2+.
2.Addition of KSCN
After KSCN was added to the supernate, a blood red complex was formed, more formally written as,
Fe3+ (aq) + SCN- (aq) → FeSCN2+ (aq)
This reaction thus proves that the supernate contains Fe3+.
3.Addition of HCl
The last test was the addition of HCl to the supernate. This produced a white precipitate, which we can formally write as,
Ag+ (aq) + Cl- (aq) → AgCl (s)
This reaction thus proves that the supernate conatins Ag+.
The 3 tests showed that all 3 ions were present in the supernate. This goes to show that the reaction between iron sulfate and silver nitrate was in a state of equilibrium since both the ions in the product and reactant side were present, meaning to say that the forward and reverse reactions were proceeding at the same rate.
Furthermore, the range of the equilibrium constant (Keq) for the reaction of iron sulfate and silver nitrate is from 10-10 to 1010 .
The second phase of the experiment dealt with the reaction between copper sulfate and ammonia, more formally written as,
CuSO4 (aq) + 2 NH4OH (aq) → Cu(OH)2 (aq) + (NH4)2SO4 (aq)
The pale blue precipitate formed at the beginning of the reaction of copper sulfate and ammonia was the Cu(OH)2.
From the pale blue color, it turned into a deep cerulean blue when 11 drops of ammonia was added. The solution then went back to the pale blue color after 4 drops of hydrochloric acid was added. The addition of hydrochloric acid added more H+ ions to the solution; therefore, drawing the equilibrium back to the reactant side. This equilibrium reaction is formally stated as,
[Cu(H2O)6]2+ (aq) + 4 NH3 (aq) ⇄ [Cu(NH3)6]2+ (aq) + H2O (l)
As observed, it took almost 3 times the number of drops of ammonia to change the pale blue color to a deep cerulean blue as compared to the number of drops of hydrochloric acid that changed the deep cerulean blue back to pale blue. This means that the reverse reaction was more spontaneous than the forward reaction.
In the third phase of the experiment, the chromate and dichromate solutions were observed. Chromate had a yellow color, while dichromate had an orange color. When sulfuric acid was added to a sample of chromate and dichromate solutions, the yellow chromate solution turned orange, while the dichromate solution remained orange. The equation for the chromate’s change in color is as follows,
2 CrO42- (aq) + 2 H+ (aq) → H2O (l) + Cr2O7- (aq)
When sodium hydroxide was added to a sample of chromate and dichromate, the chromate solution remained yellow, while the orange dichromate solution turned yellow, formally written as,
2 OH- (aq) + Cr2O7- (aq) → 2 CrO42- (aq) + H2O (l)
The change in color of chromate as hydrochloric acid was added and the change in color of dichromate as sodium hydroxide was added was due to the instability of the reactions, causing a shift in the equilibrium.
The acid H2SO4 (sulfuric acid) was used in the reaction since it’s a strong acid, and strong acids dissociate more. The added H+ ions increase the concentration, therefore, shifting the equilibrium.
Based on the observations made, it could be said that the dichromate solution is stable under acidic conditions, while the chromate solution is stable under basic conditions.
In the fourth phase of the experiment, iron trichloride was reacted to thiocyanate giving way to this reaction,
[Fe(H2O)6]2+ (aq) + SCN- (aq) ⇄ [Fe(SCN)(H2O)5]2+ (aq) + H2O (l)
The reaction produced a light orange solution, which was then tested to determine to which direction the equilibrium shifted when Fe3+, SCN- and NaCl were added. The result of the tests was as follows:
1.Addition of Fe3+
After FeCl3 was added to a sample of the solution, the light orange color of the solution became a darker shade of orange, which could be said to show a shift to the right.
2.Addition of SCN-
After KSCN- was added to a sample of the solution, the shade of orange of the solution was lighter than the solution when FeCl3 was added, but darker than the original light orange color. It could then be said that the equilibrium shifted to the right.
3.Addition of NaCl
Lastly, NaCl was added to a sample of the solution. It was observed that the resulting solution had a lighter shade of orange than that of the original. With this, it could be said that the equilibrium shifted to the left.
The shift to the left of the equilibrium was brought about by the reaction of Cl- (from NaCl) with Fe(SCN)3.
Cobalt-Cobalt Chloride Equilibrium
The last and final phase of the experiment was on the reaction of cobalt dichloride and hydrochloric acid which could be formally written as,
[Co(H2O)6]2+(aq) + 4Cl- (aq) ⇄ [Co(Cl)4]2-(aq) + 6H2O(l)
The first part of this phase reacted cobalt dichloride with hydrochloric acid. The pink color of the cobalt dichloride (Co2+) turned blue (CoCl42-) when hydrochloric acid was added. In the second part of this phase, the test tube containing cobalt dichloride was immersed in a boiling water bath. The pink color of the solution turned blue when the temperature increased. With this, it could be said that the equilibrium shifted to the right upon heating. When the temperature is increased at constant pressure, an exothermic reaction would proceed backward or shift to the left. In this case, the increase in temperature made the equilibrium shift to the right; therefore, the reaction is said to be endothermic.
The experiment basically could be summarized into the 5 different equilibrium states. First, the iron-silver equilibrium states that reactants and products could indeed co-exist in a system. Second, the copper-ammonia equilibrium states that certain chemicals could affect the state of equilibrium of a system. Third, the chromate-dichromate equilibrium states that increasing the acidity or basicity of the solution could either move the equilibrium forward or backward.
Fourth, the iron-thiocyanate equilibrium states that “An increase in the concentration of the reactant or a decrease in the concentration of the product shifts the direction of the reaction towards the production of more products to return to the equilibrium position. On the contrary, a decrease in the concentration of the reactants or an increase in the concentration of the product shifts the equilibrium position towards the production of more reactants.”  Lastly, the cobalt-cobalt chloride equilibrium states that an increase in temperature in an endothermic reaction favors product formation, therefore shifting the equilibrium to the right. On the contrary, increasing the temperature in an exothermic reaction favors reactant formation causing a shift to the left.
The methods and procedures done in the experiment are sufficient to obtain the data and results needed. No further recommendation is needed.
 Petrucci, Ralph et.al. General Chemistry: Principles and Modern Applications, 10th ed.; Pearson Canada: Toronto, Ontario, 2010.
 Padolina, Ma. Cristina et.al. Conceptual and Functional Chemistry: Modular Approach, Vibal Publishing House, Inc.: Araneta, Quezon City, 2004.
 Chemical Equilibrium. [Online]. http://www.scribd.com/doc/30015115/Expt-9-Chemical-Equilibrium (accessed last January 15, 2013)
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