EXPERIMENT NO. 9CHEMICAL EQUILIBRIUM

Apollo Gueco FCD4, Group No. 9, Sir BiadomangPaolo Ocampo March 2013

I. ABSTRACTChemical equilibrium can be characterized by the inability of the involved chemical

reaction to proceed to completion and by the reversible nature of the reaction. This reversibility is marked by the amount of reactants and products being constant and the rate of consumption of the reactants being equal to the rate of formation of the products (Engle & Ilao, 2008).

The significance of chemical equilibrium can be displayed by its role in maintaining numerous biological and environmental processes (Chemical equilibrium, n.d.).

In the experiment, the shifts in the position of the equilibrium were made possible by allowing the test tubes to change either in concentration by adding a unique reagent into the mixture or in temperature by either heating with an alcohol lamp or cooling with an iced water bath. After treating each test tube, the results were marked by a change in the color of each solution, which can be either to a lighter color or to a darker color. The change in color was essential because it indicated that a shift in the equilibrium truly occurred.

II. KEYWORDS: La Chatelier’s principle, reversible reactions, equilibrium, shift, concentration, temperature

III. INTRODUCTION A reaction in chemical equilibrium can by affected by changes in concentration, temperature and pressure, which can disrupt its equilibrium state. The reaction responds to these changes by shifting either backward to the reactant side or forward to the product side so that its equilibrium state can be achieved again. The preceding statements are with accordance to Le Chatelier’s principle (Engle & Ilao, 2008).

For experimenters to understand chemical equilibrium in an actual approach was an importance set by this experiment which was achieved by establishing its objectives. These were “to evaluate” and “to explain the effect of these changes” and “to interpret the results based on Le Chatelier’s Principle” (Committee on General Chemistry, 2012, p. 76).

IV. METHODOLOGY To initiate the experiment, a solution was prepared by placing 20 drops of 1 M iron (III) nitrate [Fe(NO3)3], 20 drops of 1 M potassium thiocyanate [KCNS] and 7 mL of distilled water into a test tube and gently shaking that tube. The color of the mixture after shaking was taken note of.

A portion of the mixture was then obtained by the use of a dropper. 7 test tubes were filled by

placing precisely 10 drops of the obtained mixture into each of them.

Each of the test tubes was treated by either adding a unique reagent or performing a common lab technique. For this experiment, only 2 common lab techniques were involved which were heating with an alcohol lamp and cooling with an iced water bath. The amount of added reagents was the same for all cases, which was 10 drops. The 1st test tube was treated by adding distilled water, the 2nd by adding 0.1 M iron (III) nitrate [Fe(NO3)3], the 3rd by adding 0.1 M potassium thiocyanate [KCNS], the 4th

by adding 0.1 M potassium chloride [KCl], the 5th by adding a pinch of sodium fluoride [NaF], the 6th by performing the heating technique and the 7th by performing the cooling technique. There was supposed to be another case where a test tube will be treated by adding 0.1 M silver nitrate [AgNO3] but the said reagent was unavailable during the time the experiment was performed.

After treating the test tubes, the color of the mixture of each tube was compared to that of the 1st

tube, which was considered as the reference by obtaining a portion of each mixture and placing it in a spot plate. The comparisons were taken note of and from there, the direction of the shift in the equilibrium was determined.

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V. RESULTS The results of the experiment were briefly presented by the succeeding table.

Table of Results

For the purpose of completion, the supposed results for the treatment of adding silver nitrate [AgNO3] were included in the preceding table and were obtained by the use of reference materials. These results will also be analyzed as if it were performed in the succeeding section.

VI. DISCUSSION Each general step in the procedure was represented correspondingly by each paragraph in the Methodology section.

The step of preparing the mixture was represented by the 1st paragraph. Preparing the mixture was important because this mixture would be used throughout the whole experiment. Thus, without this mixture, the experiment cannot be performed.

The step of distributing the prepared mixture to the test tubes was represented by the 2nd paragraph. Even a minimal portion of the mixture is sufficient for results to occur was emphasized by this step.

The step of treating each test tube was represented by the 3rd paragraph. The qualitative

comparison of the now different mixtures was made possible by this step.

The step of evaluating each test tube after the treatment was represented by the 4th paragraph. By this step, the qualitative comparison was continued and the application of the chemical equilibrium concepts to explain the results was made possible. The succeeding explanations of the results will be supported by the following reaction:

Fe3+ + CNS- FeCNS2+

The increase in concentration of the reactants brought by the common ion effect was displayed by the mixtures that were treated separately with 0.1 M Fe(NO3)3 and 0.1 M KCNS. Before interacting with the mixture, the 2 reagents being ionic compounds were turned to ions by the process of dissociation (Wicks, 2012). The effect of containing substances that share the same type of ion on a system of equilibrium is shown by the common ion effect (UCDavis, n.d.). The concentration of the Fe3+ ion from the mixture was increased by the same ion that was present in the added Fe(NO3)3. In a similar way, the concentration of the CNS- ion was increased by the same ion that was present in the added KCNS. No participation in the equilibrium was done by the NO3

- and K+ ions because they were considered as

spectator ions. The increase in concentration of the 2 cases was clearly observed by the darkening of the mixture’s color. Applying Le Chatelier’s Principle, the forward shift of the reaction’s equilibrium was caused by the increase in concentration in the reactant’s side.

The decrease in concentration of the reactants brought by the uncommon ion effect was displayed by the mixtures that were treated separately with 0.1 M KCl, 0.1 M AgNO3, and a pinch of NaF. As stated in the preceding paragraph, the 3 reagents being ionic compounds were turned first to ions by the process of dissociation before interacting with the mixture (Wicks, 2012). The uncommon ion effect was caused by the dissimilarity of the ions of the added reagents and the ions of the mixture. New substances were formed by the complex formation of these dissimilar ions (The uncommon ion effect, 2008). The concentration of the Fe3+ ion was decreased by its complex formation with the Cl- ion and the F- ion separately. FeCl3 and FeF3 were formed respectively by this process. Similarly, the concentration of the CNS- ion was decreased by its complex formation with the Ag+ ion. AgCNS was formed by this process and by its essence existed as a precipitate thus indicating that precipitation also

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Reagent/ Treatment

Observation Reaction Direction of Shift

0.1 M Fe(NO3)3

Darker color Dissociation Forward

0.1 M KCNS

Darker color Dissociation Forward

0.1 M KCl Lighter color Dissociation & complex formation

Backward

0.1 M AgNO3 Lighter color Dissociation, precipitation & complex formation

Backward

Pinch of NaF

Lighter color Dissociation & complex formation

Backward

Increase in temperature

Lighter color Release of heat in the reactant

side

Backward

Decrease in temperature

Darker color Absorption of heat in

the reactant side

Forward

occurred (Solubility product of silver sulfate, 2003). No participation in the equilibrium was done by the K+, NO3- and Na+ ions because they were considered as spectator ions. The decrease in concentration of 3 cases was clearly observed by the lightening of the mixture’s color. Applying Le Chatelier’s Principle, the backward shift of the reaction’s equilibrium was caused by the decrease in concentration in the reactant’s side.

The analysis of the results involving change in temperature was done by treating heat as part of the reactants. The color of the mixture lightened by heating it. It can be said then that the equilibrium performed a backward shift caused by the release or “decrease” of heat in the reactant side. In contrast, the color of the mixture darkened by cooling it. It can be said then that the equilibrium performed a forward shift caused by the absorption or “increase” of heat in the reactant side.

VII. GUIDE QUESTIONS AND ANSWERS1. Explain your observations on the basis of

Le Chatelier’s Principle. On the basis of the Le Chatelier's principle, the solution being lighter in color indicates that there is an addition on the product side and to relieve this condition, the equilibrium will shift towards the iron (III) and thiocyanate ions or the reactant side. The solution being darker in color indicates that there is an addition on the reactant side and to relieve this condition the equilibrium will shift towards the iron (III) thiocyanate ions or the product side.

2. Which species (ions) in the added reagents are effective in altering the state of the system?

Fe3+, CNS-, Cl-, Ag3+, and F- are the effective added ions because they show very observable change in color.

3. Is the reaction endothermic or exothermic? Why?

The reaction is exothermic because when the temperature was increased, the mixture became lighter in color. This observation implies that heat has been released in the reactant side causing the equilibrium to shift towards to that side.

VIII. CONCLUSIONS AND RECOMMENDATIONS The change in the equilibrium state of the given reaction was brought by change in concentration and in temperature and was physically shown by the change in color after a treatment was done on the mixture of a test tube. Explanations for these changes were supported by Le Chatelier’s Principle.

Future experimenters can efficiently produce true results by heeding the following recommendations:

1. Being more attentive to achieving consistency when preparing the mixture, distributing it to test tubes and adding the reagents

2. Being more assure in the reagents by double checking them in terms of content, purity and concentration to the instructor or appropriate personnel

IX. REFERENCESChemical equilibrium. (n.d.). Retrieved March 28, 2013 from http://www.pearsonhighered.com/ mcmurryfay6einfo/downloads/MCMUC13.pdfClark, J. (2002). Le Chatelier’s principle. Retrieved March 19, 2013 from http://www.chemguide.co. uk/physical/equilibria/lechatelier.html Committee on General Chemistry. (2012). Laboratory manual in general chemistry 1. Philippines: University of the Philippines Manila.Engle, H.L. & Ilao, L.V. (2008). Learning modules in general chemistry. Philippines: University of the Philippines Manila.Le Chatelier’s principle. (n.d.) Retrieved March 19, 2013 from http://www.adichemistry.com/physical /equilibrium/le-chatelier/le-chatelierprinciple.htmlLe Chatelier’s principle lab. (2013). Retrieved March 19, 2013 from http://chemistrylabsguy.weebly. com/le-chateliers-principle-lab.htmlOlomodosi, A. (2010). Demonstrating Le Chatelier’s principle. Retrieved March 19, 2013 from http:// prezi.com/txirlcieyh-m/demonstrating-le- chateliers-principle/Solubility product of silver sulfate. (2003). Retrieved March 28, 2013 from http://dhs.dist113.org/ faculty/HintonM/website/apchem/lab-02-2.pdfUCDavis. (n.d.). Common ion effect. Retrieved March 28, 2013 from http://chemwiki.ucdavis .edu/ Physical_Chemistry/Physical_ Properties_ of_Matter/Solubilty/Common_Ion_EffectWicks, K.M. (2012). Dissociation of ionic compounds and acids. Retrieved March 28, 2013 from http:// www.chemistrylecturenotes.com/html/ dissociation_of_ionic_compound.html

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

_______________ ________________Apollo Gueco Paolo Ocampo

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