Chemical equilibrium and the factors that affects the position of it is a topic of interest in industrial chemistry. I was first introduced to this concept when I studied about how the ideal value of temperature and pressure is decided in Haber process and Contact process. The fact that intrigued me was that if there are any factors apart from temperature, pressure, concentration and catalyst which may impact the position of equilibrium or the magnitude of equilibrium constant. After going through Topic-8 and getting a clearer idea about the concepts of pH, I wondered if the variation of pH at which an equilibrium is established would in any means affects the position of equilibrium and the magnitude of equilibrium constant. Thus, I arrived at the research question stated above.

A reaction is considered to be reversible in nature if the products can be transformed back into the reactants and thus the reaction can proceed both forward and backward. A reversible reaction attains equilibrium only when both the forward and backward reaction occurs at the same rate and the concentration of reactant or product remains constant with time.

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The equilibrium considered in this investigation is:

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Fe3+ (aq) + SCN- (aq) &larr;------&rarr; [Fe (SCN)]2+

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The expression of equilibrium constant would be:

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\(Kc=\frac{[[Fe(SCN)]^{2+}]}{[Fe^{3+}][SCN^-]}\)

Let us consider that the equilibrium concentration of the [Fe(SCN)]2+ be x mol dm-3 and the initial concentration of Fe3+ and SCN- be 0.10 mol dm-3.

\(\text{Equilibrium constant} (Kc =)\frac{[[Fe(SCN)]^{2+}}{[Fe^{3+}][SCN^-]}=\frac{x}{(0.10-x)(0.10-x)}=\frac{x}{(0.10-x)^2}\)&nbsp; &nbsp; &nbsp; &nbsp; &nbsp; &nbsp;(equation - 1)

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Thus, if the value of equilibrium concentration of the product is determined, the value of the equilibrium constant can be easily calculated.

Although the position of equilibrium is affected by several factors like concentration of the reactants or products, pressure, presence of catalyst and temperature, equilibrium constant is affected only by temperature. For an exothermic reaction, as temperature increases, the equilibrium moves more towards the reactant and thus the magnitude of equilibrium constant decreases while for an endothermic reaction, as temperature increases, the equilibrium moves more towards the product and thus the value of equilibrium constant increases. Thus, for exothermic reaction the temperature and equilibrium constant are indirectly related while for an endothermic reaction, they are indirectly related.

Fe3+ (aq) + SCN- (aq) &larr;----&rarr; [Fe(SCN)]2+

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Iron (III) reacts with thiocyanate (SCN-) and results in the formation of the complex Iron (III) thiocyanate ion. Here, Iron(III) acts as a Lewis acid and accepts a pair of electron from the lone pair on the N atom of the ligand SCN-. The complex formed is a red color complex.

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The equilibrium constant for this equilibrium can be expressed as:

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\(Kc =\frac{[Fe(SCN)]^{2+}}{[Fe^{3+}][SCN^{-}]} mol-1 dm3\)

As per, Beer Lambert law.

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A = &isin;&times; c &times; l

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Here,

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A = absorbance in abs

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&isin; = molar absorptivity contact in abs mol-1 dm2

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C = molar concentration in mol dm-3

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l = path length in dm

The purpose of this investigation is to elucidate the effect of pH on the magnitude of equilibrium constant. To do this, the equilibrium will be established at various pH levels. Dilute solutions of HCl or NaOH may be used to vary the pH of the medium. After that, the absorbance of the equilibrium solution at a wavelength at which the product displays maximum absorbance will be recorded using a photo-colorimeter. A calibration curve from literature will be considered and the equation from that would be used to calculate values of concentration from the equation of trend line. Using the values of concentration and the expression of the equilibrium constant (equation - 1), the magnitude of equilibrium constant will be computed.

An alternate procedure to measure the equilibrium constant would be to measure the amount of Iron-III left instead of measuring the concentration of the complex. Fe-III can be quantitively measured using iodometry and starch as an indicator.

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Fe3+ + 2 I- ----&rarr; I2 + 2Fe2+

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This reaction is quantitative in nature and requires an acidic medium to be carried out. Thus, this method may not be effective enough when the reaction is performed in a basic pH. Moreover, the results from a colorimeter are always more accurate and reliable than that from a redox titration as it is a digital device and has less instrumental error.

To delineate the relationship between pKa of various acidic functional groups in an enzyme and the kinetics of a enzyme substrate complex, a theoretical study was conducted. This study was done based on the &ldquo;rapid equilibrium model &ldquo; assumption. This study has revealed that there is a correlation of the velocity constant of a reaction and the pka of the acidic functional groups as well as the concentration of H+ ions in the medium. A mathematical relationship was established between the magnitude of equilibrium constant and the pH of the medium : Kc = 10 npH where n is a constant that depends on the type of the enzyme.

pH of the medium. The pH of the medium was varied in the basic region from 7.00, 8.00,9.00,10.00,11.00, 12.00 and 13.00. 0.10 molar NaOH was made and diluted to create different values of pH. For example, if 0.10 molar NaOH is diluted 10 times, the concentration would become 0.01 molar and the pH would become 2.00. As the purpose of the investigation was to understand the effect of basicity of the medium on the value of equilibrium constant, the pH was varied in this basic region.

The magnitude of equilibrium constant is the dependent variable. It will be measured in mol-1 dm3. The absorbance of the solution will be measured using a colorimeter and a calibration curve will be used to compute the values of concentration. Using the value of equilibrium concentration of the Iron (III) thiocyanate complex, the value of equilibrium constant will be calculated.

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list of chemicals required.svg

<ul>
	<li>NaOH is corrosive in nature. If exposed to skin, it may cause burns and irritations. Use a safety mask and gloves.</li>
	<li>Inhalation of Iron (III) nitrate nona hydrate may cause convulsions, breathing issues.</li>
	<li>Continuous exposure to Potassium thiocyanate may cause dizziness, chronic skin diseases and nausea.</li>
</ul>

<ul>
	<li>Use safety masks to prevent inhaling any of the chemicals.</li>
	<li>Use a safety gloves and a laboratory coat to prevent exposure of the chemicals to skin.</li>
	<li>Use tongs carefully to avoid spillage especially when weighing the chemicals and transferring them to beakers.</li>
	<li>Handle the concentrated HCl with utmost care and preferably in presence of a laboratory technician.</li>
</ul>

Use of toxic chemicals was prohibited.

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Minimum amount of chemicals was used.

The waste liquids were disposed in the waste bin

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The unused solutions were preserved for re-use.

Does The Value Of Equilibrium Constant For The Reaction Between Iron (Iii) & Thiocyanate In Basic Medium (Using Naoh) Depend On The Basicity Of The Medium (Expressed In Terms Of Ph Ranging From 8.00 To 13.00), Determined Using Colorimetric Estimation Of Iron-iii?

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