Experiment no. 9Chemical Equilibrium

Agustin, Victoria T.Crisostomo, Jan Christine R.Morales, Jessica Christine C

Group # 9, Chem. 14.1, WEG1, Mr. Ralph Julius L. MendozaMarch 18, 2009

Chemical equilibrium is an extremely important process in nature particularly in many industrial (e.g. production of ammonia) and biological processes (production of hemoglobin in relation to altitude). Experiment 9, chemical equilibrium, will determine how various stresses, according to Le Chatelier’s Priciple, being introduced in a system at equilibrium can alter the system’s equilibrium position by shifting in direction to counteract the effect of the stress. These stresses include increase or decrease in concentration, temperature, and pressure. To know the effect of the change in concentration, different reactants are added in the initial mixture and for the determination of the effect of temperature change, two mixtures with the same components are used: the temperature in one of the mixtures is increased while the temperature in the other mixture is decreased. For both set-ups, an undisturbed mixture is used as reference for the comparison in the color of the disturbed mixtures. A darker disturbed mixture will mean a forward reaction, conversely, a lighter disturbed mixture will mean a backward reaction. Again, the significance of this color change together with its corresponding shift in direction is to tell that the reaction is going towards its natural state which is to be in equilibrium.

Keywords: reversible reactions, equilibrium, chemical equilibrium, Le Chatelier’s principle, stress

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Introduction: Equilibrium is a state where there are no

observable changes as time goes by, that is, it is constant in time and space. Chemical reactions, particularly reversible reactions, have the tendency to alter its conditions to achieve equilibrium. At this chemical equilibrium, the rates of the forward and reverse reactions are equal. Furthermore, the concentrations of the products and reactants remain constant. Chemical equilibrium is important in understanding biological and industrial process such as production of hemoglobin in relation to higher altitudes and production of ammonia.

Le Chatelier’s principle is used to predict the direction or shift of the equilibrium position when stress such as change in concentration, pressure, volume, or temperature occurs in the reaction.

This experiment will focus more on the evaluation of the effects of changes in concentration and temperature on the equilibrium, explain these effects on the equilibrium system, and interpret these consequences through the Le Chatelier’s principle.

Experimental:20 drops of 1 M Fe(NO3)3 solution and 20

drops of 1 M KCNS solution were mixed in a 10-ml test tube, and 7 ml of water was added. The mixture was then shook and observed. After the solution was prepared, eight test tubes with labels were filled with ten drops of the solution. 10 drops of each of the following: distilled water, 0.1 M Fe(NO3)3, 0.1 M KCNS, 0.1 M KCl, 0.1 M AgNO3, and a pinch of NaF, was added to test tubes 1,2,3,4,5, and 6 respectively. Observations were made on test tubes 2 to 6 using test tube 1 as reference. The changes in the intensity of color or appearance of the mixture were noted. The observations recorded after each addition of reagent and the reactions involved were analyzed. On the other hand, 10 drops of distilled water was added to test tubes 7 and 8. Test tube 7 was placed in ice water and was then compared to test tube 1. The solution in test tube 8 was heated over a low Bunsen flame.

Results:Given the following reaction: Fe3+ + CNS- FeCNS2+

orange colorless blood-red

Table 1: Results gathered from the reactions

Discussion:Le Chatelier’s Principle states that if an

external stress is applied to a system at equilibrium, the system adjusts in such a way that the stress is partially offset as the system reaches a new equilibrium position. Change in concentration and

change in temperature are the most common forms of stress that affect a system in a state of equilibrium. It should be noted that a system has a tendency to go back to its natural state which is the state of equilibrium.

In the experiment, the initial mixture was composed of Fe(NO3)3 (orange) and KCNS-

(colorless) dissolved in 7 ml of water. After the solute had completely dissolved, a dark-colored mixture was obtained. Each of the eight labelled test tubes was filled with 10 drops of the initial mixture. In test tube # 1, 10 drops of distilled water was added. Upon adding, no change occurred because water as a solvent did not alter the concentration of the mixture. Test tube # 1 was used, too, as reference for comparison, in terms of color of other disturbed mixtures.

In test tube # 2, 10 drops of 0.1M Fe(NO3)3

was added in the mixture, experimentally producing a lighter colored product as compared to test tube # 1, but theoretically the product should have been darker. Keeping in mind that moles are additive, the addition of 0.1 M of Fe(NO3)3 to the initial mixture having 1 M of Fe(NO3)3 means that the reagent increased the concentration of the reactant side causing a shift to the right. Meanwhile, 10 drops of 0.1 M KCNS was added in test tube # 3, resulting to a darker-colored product. The same concept is observed for test tube # 3.

In test tube # 4, 0.1 M KCl was added in the mixture which yielded a lighter-colored product than that of the test tube # 1. This means that there was a shift to the left. Theoretically, since KCl is a salt, adding it to the initial mixture will cause it dissociate into K+ and Cl-. These ions did not interact with Fe3+

and CNS-. This non-interaction means that the dissociation of KCl in this case will yield spectator ions, or ions that are not involved in the overall reaction. Therefore, same color should have been observed which means that no shift occurred.

In test tube # 5, 0.1 M of Ag3NO3 was added in the mixture producing a lighter mixture with a precipitate causing it to shift to the left. This reaction can be described by the following equation: Ag3NO3

+ Fe3+ + SCN- AgSCN + Fe(NO3)3. Applying the solubility rules, the nitrate salt [in this case] Fe(NO3)3

is soluble in the mixture while salts with a transition metal [in this case] AgSCN is insoluble in the mixture, making AgSCN the precipitate.

Chem. 14.1, Chemical Equilibrium Page 2 of 4

Reagent /Treatment Observation (compared to test tube #1)

Direction of Shift

0.1 M Fe(NO3)3 Became lighter to the left0.1 M KCNS Became darker to the right0.1 M KCl Became lighter to the left0.1 M AgNO3 Became lighter to the leftPinch of NaF Became lighter to the leftIncrease in temp. Became lighter to the leftDecrease in temp. Became darker to the right

In test tube #6, a pinch of NaF was added producing a lighter colored solution that caused a shift to the left. In addition to this, white substances were formed but these are not what we call precipitate. Instead, these substances are complex ions? bat xa ndi nagdissolve? ang complex ions ba insoluble? The reaction is described by the equation NaF + Fe3+ + SCN- FeF3+ + NaSCN. NaSCN, a soluble salt and FeF3+, a complex ion (an ion containing a central metal cation bonded to one or more molecules or ions), were formed. The formation of FeF3+ increased the solubility of NaF.

In test tube # 7, 10 drops of distilled water was added, in addition to that, heat was also introduced in the system. After increasing the temperature to 67o C, the product yielded was lighter as compared to test tube # 1 and this means that the shift was to the left. This implies that heat was added in the product side making the reaction exothermic. It should be noted that adding heat is similar to the concept of adding concentration. Meanwhile in test tube # 8, 10 drops of distilled water was added in the mixture and was subjected to a significant decrease in temperature. Since the reaction is exothermic, “removing” that is, decreasing the temperature [in the product side] would cause a shift to the right accounting for the observed darker solution. It formed a darker solution as compared to test tube # 1 indicating that the shift was to the right. *ndi na ata endothermic kasi ang tanong ay kng exo or endo since pinili na natin ung exo, ndi na pedeng endo.

Conclusion and Recommendations:Chemical equilibrium is achieved when the

reaction in the reactant side, as well as in the product side proceeds at the same rate and the concentration of products and reactants ceases to change. The experiment showed how changes in concentration and temperature affect the direction of equilibrium. The La Chatelier’s Principle which states that if a chemical system is at equilibrium and we add a substance (either a reactant or a product), the reaction will shift so as to re-establish equilibrium by consuming part of the added substance. Conversely, removing a substance will cause the reaction to move in the direction that forms more of that substance described the shifting taking place in a system at equilibrium.

Changing the concentration of a substance in equilibrium will cause a shift in the system in the direction which is opposite to the side where the stress is applied to counter act the effect of the said stress. Furthermore, same principle is applied when there is a change in the temperature since heat is treated as either an agent which may be added in the product or in the reactant side.

One must be very careful in observing the changes in color that occurred during the experiment because these changes crucial in determining the shift in the experiment.

References:- Chemistry: The Central Science (Ninth Edition) by Brown, Lemay, and Bursten- Chemistry (Ninth Edition) by Chang- http://pages.towson.edu/ladon/solprod.html-http://www.chem1.com/acad/webtext/chemeq/Eq01.html- The Advanced Placement Examination in

Chemistry by the College Entrance Examination Board, Princeton, NJ

- http://www.public.asu.edu/~jpbirk/CHM-115_BLB/Chpt17/

I hereby certify that I have given substantial contribution to this report.

________________________Agustin, Victoria T.

________________________Crisostomo, Jan Christine R.

________________________Morales, Jessica Christine C.

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Fe (NO3)3

KCNS H2O

Initial mixture: shake, color?

Into 8 test tubes

TT #1: + distilled H2O [reference test tube]

TT#8: + distilled H2O+ increase in

temperature

TT #7: +distilled H2O+ decrease in temperature

TT#6: + a pinch of Na F

TT#5: + AgNO3

TT #4: + KCl

TT #3: + KCNS

TT #2: + Fe(NO3)3

Observe changes in intensities of colors and

compare with TT#1