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The purpose of this experiment was to study the effects of concentration and temperature changes on the position of equilibrium in a chemical system and to observe the common-ion effect on a dynamic equilibrium (Beran, 2009). LeChatelier’s Principle states that if and external stress is applied to a system in a state of dynamic equilibrium, the equilibrium shifts in the direction that minimizes the effect of that stress (Beran, 2009). Dynamic equilibrium can be defined as the condition in which a chemical system has reached a state where the reactants combine to form the products at a rate equal to that of the products re-forming the reactants (Tro, 2010). Most chemical reactions do not produce 100% yield of product because of the chemical characteristics of the reaction. After a certain period of time, the concentrations of the reactants and products stop changing (Beran, 2009). This indicates the chemical reaction is in equilibrium and when a reactant is added, and then it shifts the solution either left or right as an attempt to relieve stress and compensate for the change. In this experiment, the solution was brought out of equilibrium indicated by either a color change or precipitate formation.
In part A of the experiment, concentrations of the reactants and products were changed to indicate the change of colors when drops of concentrated NH3 were added to 0.1 M of CuSO4, shifting the solution to the right, forming ammonia-complex ions. When the strong acid HCl is added, this removes the ammonia from the equilibria and the reactions shift left to relieve the stress. The net ionic equation can be represented as
[Cu(H2O)4]2+ (aq) + 4NH3 (aq) ïƒŸïƒ [Cu(NH3)4]2+ (aq) + 4H2O (l)
In part B of the experiment, silver ions such as carbonate, chloride, iodide, and sulfide are used in this experiment as they form precipitates, dissolve precipitates, and form gaseous substances. Nitric acid is added to silver carbonate in order to shift the equilibrium to the right. The addition of NH3 removes the silver ion, shifting the equilibrium to the left, causing AgCl to dissolve. Next, adding H+ reforms solid silver chloride, shifting the equilibrium right again. Then, adding the iodide ion to the equilibrium results in the formation of the precipitate of solid silver iodide, shifting the equilibrium to the left again. The net ionic equation for the reaction is represented as:
Ag2CO3 (s) ïƒŸïƒ 2Ag+ + CO32- (aq)
The changes in concentrations affect the equilibrium, significantly. If the concentration of a reactant was increased, to reduce the concentration, the system shifts to the left or towards the reactant, to make it as the product. This increase in concentration can occur to any species of the equilibrium reaction and the system would shift towards the increased concentration (Beran, 2009). The concentration can be chemically increased with additions of aqueous substances that can react with the equilibrium reaction. However, the same is not true for acids and bases. When an acid concentration is increased, the system tends to the opposite and thus flows towards the base, to increase its concentration (Beran, 2009). To prevent any changes in the acid-base pH, buffers are mainly used to sustain the equilibrium. Buffers have to consist of a weak or strong acid or base and its conjugate species (Tro, 2010).
As for parts D and E, the common ion effect and the temperature effect are tested. By adding drops of HCl to 1.0 mL of CoCl2 and by comparing the color change to that of 1.0 mL of CoCl2 placed in a hot water bath. The ionic equation for the reactions can be represented by:
4Cl- + Co(H2O)62+ + Heat ïƒŸïƒ CoCl42- + 6H2O
Materials and Procedures:
Please refer to Experiment 16 on pages 201-212 of Laboratory Manual for Principles of General Chemistry by J.A. Beran. Note that part C of the experiment was not performed.
Data and Results:
Table 1: Metal-Ammonia Ions
Table 1 shows the colors that the solution changed to as reactants were added.
Table 2: Multiple Equilibria with the Silver Ion
Observation of Na2CO3 combine with
AgNO3 and net ionic equation
Ag2CO3(g) â†” 2Ag+(aq) + CO32- (aq)
HNO3 was add to solution
The solution turned colorless and shifted to the right with the formation of water and carbon dioxide
HCl addition and net ionic equation
Ag +(aq) + Cl-(aq) â†” AgCl(g)
Cloudy Solution and shifts right
Ag +(aq) + Cl-(aq) â†” AgCl(g)
Clear Solution and shifts left
Cloudy and shifts to the right
Excess NH3 addition again
Clear again and shifts to the left
Turned to off white and shifts left
Na2S addition and net ionic equation
Ag2S(s) â†” S2-(aq) + Ag(aq)
Turned dirty brown
Table 2 shows the changes in the solution as reactants were added.
Table 3: A Buffer System
Bronstead acid equation
CH3COOH(aq) + H2O(l) â†” H3O+(aq) +
Color of universal indicator with CH3COOH
Red with pH~4
Color of universal indicator with NaCH3CO2
Light Red with pH~2
Shift to the left
Color of universal indicator in H2O
Orange with pH~6
After HCl was added
Well A1 with buffer, pH~4 (red), Î”pH~2
Well B1 with water, pH~3 (light red) Î”pH~3
After NaOH was added
Well A2 with buffer, pH=5 (red), Î”pH~3
Well B2 with water, pH=12 (blue), Î”pH~6
Table 3 shows the effect on equilibrium with the change in pH and the containment of a buffer system.
Table 4: Equilibrium (Common-Ion Effect)
Color of CoCl2(aq)
HCl addition and net ionic equation
4Cl-(aq) + [Co(H2O)6]2+(aq) â†” [CoCl4]2-(aq) + 6H2O(l)
Turns purple. Shifts to the left.
Water added to solution
Turns light red again
Table 4 shows the change in equilibrium due to the common-ion effect.
Table 5: Equilibrium (Temperature Effect)
Color at room temperature
Color when heated
Table 5 shows color change due to the temperature effect on equilibrium.
In Table 1, the solution CuSO4 was in equilibrium in a light blue color until the equilibrium shifted to the right as drops of NH3 were added, forming the product. When the acid HCl was added, the equilibrium shifted left again, turning the color of the solution light blue again. This was because the H+ ions equalized the Cu(NH3)42+ that formed. Because this chemical system was in a state of dynamic equilibrium, the color of the solution was able to turn light blue again as the products can reform the reactants.
In Table 2, 1/2 mL 0.01 M AgNO3 + 1/2 mL 0.1M Na2CO3 formed a precipitate. When HNO3 was added, this showed a change in equilibrium by dissolving the precipitate, forming a clear solution. The action was reversed again as HCl was added and a precipitate was formed again, supporting the dynamic equilibrium. Also supporting the dynamic equilibrium, by showing the reverse chemical reaction, the addition of concentrated NH3 then dissolved the precipitate that was formed. Once the second of addition of HNO3 was added, a white gas formed because there was too much concentration of HNO3. When NH3 was added after that, more white gas formed because there was too much concentration of HNO3. When NH3 was added after that, more white gas formed because the concentration of the product was now also in excess. The addition of KI formed a white foggy gas on top of the solution, which was probably a result of the iodide ion and excess concentration of the reactant. Not only did the equilibrium change, but a physical change occurred as well. The gas was less dense than the solution. The addition of Na2S turned the solution partially brown. This was because of an excess concentration of the product. In the test tube, there was a white has from the potassium iodide, a brown later between the white has and the clear liquid from the sulfide, and the clear liquid from the normal concentrations of the solution of silver, chloride, and nitric acid.
In part C with the buffer systems. The Bronstead acid equation was used to obtain the initial pH values. Table 3 shows the initial pH values with the color indicators. As HCl was added the reaction shifted to the left to balance the acid and increased the base. The effect of buffers on the equilibrium can be seen. As the strong acid HCl was added, the change in pH of the water system was higher than the change in the buffer system. Most importantly, as the strong base NaOH was added, the change in the water system increased significantly to 12. However, the buffer system stayed within its range. Thus, the useful property of buffers is accepted as they contain the high changes in the pH.
Tables 4 and 5 show the color changes using the common ion effect and temperature change. Both the addition of the concentration of HCl into CoCl2 and the placement of CoCl2 into a hot water bath both changed the equilibrium of the solution by turning it purple. The equilibrium shifted back to red when water was added again or the temperature was back to normal room temperature.
The systems under states of dynamic equilibrium shifted in directions to minimize the effect of the stress that was placed upon the system. In this experiment, the effects of stress were caused by changes in concentration and temperature. LeChatlier’s Principle was supported based on the experiments conducted where colors of the solutions were reversed and precipitates could be dissolved and formed again without the concentrations being in excess. Based on the experiment, the hypothesis stated if the equilibrium of a system in a state of dynamic equilibrium shifted left or right, then LeChatlier’s Principle would be supported. The hypothesis and LeChatlier’s Principle were both supported.
There are many improvements that can be done to this experiment. One of them is that there can be more tests done to see how the equilibrium is affected in the metal ammonia ions. Another improvement that can be made to the experiment is comparing to the pH chart to infer and analyze the color changes. A third improvement is further studying the effect of temperature in the equilibrium changes. This would give a better idea on its effects.
Post Lab Questions:
- Predict what would appear in the solution if NaOH were been added to
CuSO4 solution instead of NH3?
More of a bluer color would be observed with increase in formation of Cu and Ni ions.
- HNO3, a strong acid is added to shift the Ag2CO3 equilibrium to the right. Why does this shift occur?
The removal of the CO32- ions means that have to be replace so a shift to the right is necessary.
- Predict what would happen if .1M NaBr had been added to solution in part
B.3 instead of the Na2S solution. Explain?
It would probably combine with Br ion to form silver bromide due that the NaBr is more soluble, and thus is more able to remove the iodide ion from the AgI formed from the last equation when KI was added.
- Write an equation that shows the pH dependence on the chromate, dichromate equilibrium system.
CrO42- â†” Cr2- + 4O
Cr2O72- â†” 2Cr2- + 7O
- When 5 drops of .10 M HCl is added to 20 drops of a buffer solution that is .10 M CH3COOH and .10 M CH3CO2- only a very small change in pH occurs. Explain?
The CH3CO2- equaling the amount of moles in the acid and ends up absorbing all the acid and cause of this there is a decrease in CH3CO2- and an increase of CH3COOH which is where the small change in pH occurs.
Beran, J.A. (2009). Pages 201-212. Laboratory Manual for Principles of General Chemistry. Hoboken: John Wiley & Sons.
Tro, N. J. (2010). Pages 516-517. Principles of Chemistry: A Molecular Approach. Upper Saddle River: Pearson Education.
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