Does the value of equilibrium constant for the reaction between Iron (III) and thiocyanate in basic medium (using NaOH) depends 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?

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.

The equilibrium considered in this investigation is:

Fe³⁺ (aq) + SCN^- (aq) ←------→ [Fe (SCN)]²⁺

The expression of equilibrium constant would be:

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

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.

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

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) ←----→ [Fe(SCN)]²⁺

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.

The equilibrium constant for this equilibrium can be expressed as:

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

As per, Beer Lambert law.

A = ∈× c × l

Here, A = absorbance in abs

∈ = molar absorptivity contact in abs mol^-1 dm²

C = molar concentration in mol dm^-3

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.

Fe³⁺ + 2 I^- ----→ I₂ + 2Fe²⁺

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 pK_a 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 “rapid equilibrium model “ 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.

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

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

Use of toxic chemicals was prohibited.

Minimum amount of chemicals was used.

The waste liquids were disposed in the waste bin

The unused solutions were preserved for re-use.

• Preparation of 100 cm³ of 0.10 moldm^-3 Iron (III) nitrate solution:

Molar mass of Fe(NO₃)₃.9H₂0 = 404 g mol^-1

(This value will change if you are using the anhydrous Iron-III nitrate or any other Ferric salt and thus the calculations will change too. However, this does not change the procedure).

Concentration = 0.10 mol dm^-3

Volume = 100 cm³ = 0.10 dm³

Moles = concentration × Volume = 0.10 × 0.10 = 0.01

Mass = moles × molar mass = 0.01 × 404 = 4.04 g

A top-pan digital mass balance was switched on.

A watch glass was placed on the top pan and the reading was tared to zero.

The solid Iron (III) nitrate was transferred from the reagent bottle to the watch glass using a spatula until the balance reads 4.04 ± 0.01 g The weighed solid was then transferred to a 100 cm³ volumetric flask.

Distilled water was added to the flask till the mark.

The lid of the flask was closed and the flask was shaken to homogenize the solution.

• Preparation of 100 cm³ of 0.10 moldm^-3 potassium thiocyanate solution:

Molar mass of KSCN = 97.18

(This value will change if you are using the thiocyanate salt and thus the calculations will change too. However, this does not change the procedure).

Concentration = 0.10 mol dm^-3

Volume = 100 cm³ = 0.10 dm³

Moles = concentration × Volume = 0.10 × 0.10 = 0.01

Mass = moles × molar mass = 0.01 × 404 = 4.04 g

A top-pan digital mass balance was switched on.

A watch glass was placed on the top pan and the reading was tared to zero.

The solid Iron (III) nitrate was transferred from the reagent bottle to the watch glass using a spatula until the balance reads 4.04 ± 0.01 g The weighed solid was then transferred to a 100 cm³ volumetric flask.

Distilled water was added to the flask till the mark.

The lid of the flask was closed and the flask was shaken to homogenize the solution.

Determining absorbance of Iron-thiocyanate complex at 447 nm at pH=13.00:

• A 100 cm³ glass beaker was taken.
• A 10 cm³ graduated pipette was used to transfer 10.00 ± 0.05 cm³ of Iron-III solution to the beaker.
• Another 10 cm³ graduated pipette was used to transfer 10.00 ± 0.05 cm³ of thiocyanate solution.
• 0.40 g of NaOH was weighed using a watch glass and a digital mass balance and added to the beaker.
• Distilled water was added to make the total volume to the mark of 100 cm3.
• The stop-watch was started.
• The colorimeter was switched on.
• The cuvette was taken and filled up with distilled water.
• The wavelength of the colorimeter was set at 447 nm or the value closest to this.
• The cuvette was inserted in the colorimeter and the reading of the absorbance was adjusted to 0.000 abs using the ‘Set Zero’ button (or the button used to calibrate or adjust the absorbance readings).
• The cuvette was taken out and the water inside it was disposed off.
• It was wiped with a soft tissue and dried using a hair-dryer.
• As soon as the stop-watch to read 60.00 ± 0.01 minutes, 1.00 ± 0.05 cm³ of the reaction mixture was taken out from the reaction mixture and kept in the cuvette (the test tube or glass tube used in the colorimeter).
• The reading of the absorbance of the sample was taken keeping the wavelength at 447 nm.
• Repeat steps 15 & 16 for 4 more times.

The same process was repeated for other pH values – 12.00, 11.00, 10.00, 9.00, 8.00 and 7.00. Serial dilution method was followed for all other pH values and distilled water was used for pH value of 7.00. For example, a 0.10 moldm^-3 NaOH solution was diluted 10 times for the pH value of 12.00. Refer to appendix for more details.

Average absorbance of Iron (III) thiocyanate complex at 447 nm at pH = 13.00 \(=\frac{0.497+0.498+0.492+0.497+0.495}{5}\) = 0.496 ± 0.001 abs

Standard deviation (SD) \(=\frac{(0.497-0.496)^2+(0.498-0.496)^2+(0.492-0.496)^2+(0.497-0.496)^2+(0.495-0.496)^2}{5}\) = 0.002

Deriving an equation between absorbance and concentration using a literature calibration curve:

The equation: y = 4312x + 0.0075

Absorbance (y) = 4312 × molar concentration (x) + 0.0075

Molar concentration (x) = \(\frac{absorbance\ (y)-0.0075}{4312}\) mol dm^-3

Fe³⁺ (aq) + SCN^- (aq) -----→[Fe (SCN)]²⁺

Equilibrium constant (Kc) = \(\frac{[Fe(SCN)]^{2+}}{[Fe^{3+}][SCN^-]}=\frac{x}{(0.10-x)^2}\)

At pH = 7.00,

Mean absorbance = 0.496 ± 0.001 AU

Mean concentration (x) = \(\frac{0.496-0.0075}{4312}\) = 1.13 × 10^-4 ± 0.001 mol dm^-3

Equilibrium constant (K_c) = \(\frac{1.13×10^{-4}}{(0.10-1.13×10^{-4})^2}\) = 1.14 × 10^-2 mol^-1 dm^-3

At pH = 7.00

Mean absorbance (A) = 0.496 ± 0.001 abs

Mean concentration (x) = \(\frac{0.496-0.0075}{4312}\) = 1.13 × 10^-4 ± 0.001 mol dm^-3

Percentage error in equilibrium constant = \(\frac{∆k_c}{k_c}×100=\frac{±0.001}{1.74×10^{-2}}×100=±5.74\)

The graph above clearly shows that the magnitude of the equilibrium constant is decreasing from 1.74 × 10^-2 mol^-1 dm³. to 1.14 × 10^-2 mol^-1 dm³ as the pH of the medium increases from 7.00 to 13.00. This indicates that as the medium becomes more basic in nature, the magnitude of equilibrium constant is decreasing that is the equilibrium is shifting more towards the reactants.

The correlation between the equilibrium constant and pH of the medium is represented using an equation of trend line: y = - 0.1046 x + 2.4969 where y indicates the magnitude of equilibrium constant and x represents the pH of the medium. As the difference between the consecutive data points are almost the same, the decrease of values of equilibrium constant with pH can be considered to be gradual and uniform. As the values of percentage errors are not uniform, it has not been displayed on the graph. Standard error bars have been displayed in the graph that shows some random errors in the data collected.

Fe³⁺ (aq) + SCN^- (aq) <---→ [Fe(SCN)]²⁺

As the pH of the medium increases above 7.00, it becomes a basic medium and thus contains OH- in excess of H₃0⁺ ions. Fe³⁺ can react with OH- according to the equation below:

Fe³⁺ (aq) + 3 OH^- (aq) <---→ [Fe(OH)₃]²⁺ (S)

This reaction results in the formation of a orange solid Iron (III) hydroxide. As the medium is basic, the compound is insoluble in water and thus separates out as an orangish red precipitate. As the pH of the medium increases, the concentration of OH^- in the medium increases and thus as a result, more Iron(III) ions reacts with OH-. This eventually reduces the concentration of Iron (III).

Thus, according to the Le-Chateleir’s principle, the equilibrium shifts towards the left. This ultimately increases the amount of reactants and reduces the amount of products. Thus, the magnitude of the equilibrium constant decreases.

• The magnitude of the equilibrium constant is decreasing from 1.74 × 10^-2 mol^-1 dm³ to 1.14 × 10^-2 mol^-1 dm³ as the pH of the medium increases from 7.00 to 13.00. Thus, it can be concluded that as the medium becomes more basic, the concentration of OH^- ions increases and the magnitude of equilibrium constant decreases.
• The decrease in the value of equilibrium constant indicates that the equilibrium is shifting more towards the reactants side.
• Thus, as the medium becomes more basic, the equilibrium moves more towards the reactant and this in turn causes the amount of product to decrease. This is well supported by the fact that as the pH value increases, the values of absorbance of the Iron-III complex formed decreases.
• In Graph-1, the value of gradient reported in the equation of trend line is -0.1046 x. The value of gradient is negative which also supports the fact that there is a negative correlation between the values of equilibrium constant and the pH of the medium. An equation of trend line has also been plotted in the graph displayed and the equation is: y = -0.1046x + 2.4969
• The value of the correlation coefficient (R² ) as obtained using MS-Excel is also shown in the graph and it is 0.9959. This value also indicates that there is a strong negative correlation between the magnitude of equilibrium constant and the pH of the medium.
• As the graphical analysis confirms a strong negative correlation between the values of equilibrium constant and the pH of the medium, the null hypotheses has been rejected and the alternate hypotheses has been accepted.

• A wide range of independent variable has been chosen.
• Low values of percentage error indicate minimum systematic error in the investigation and thus the result can be considered to be accurate.
• The methodology is simple and does not involve the use of any complicated apparatus or chemicals.
• The raw data has been collected in trials. This process creates reproducible and replicable data increasing the accuracy of the processed data.

A calibration curve has been used in the data analysis. The equation between absorbance and concentration as indicated in the calibration curve has been utilized to calculate the values of the concentration of the complex. But the reaction conditions in the investigation and that in the calibration curve do differ. To be more specific, the pH of the reaction in the investigation has been varied while that based on which the calibration curve was made was in a particular pH. Thus, in one way, the equation obtained for a reaction at one particular value of pH has been used to analyze the data for reaction happening at various pH values. This limits the reliability of the data processing. This could have been optimized by preparing a calibration curve at various pH which was not possible as the complex required to make the curve was not available. This is a methodological limitation and thus introduces an inherent processing error.

To extend the investigation further, I would like to investigate the effect of temperature on the value of equilibrium constant. To do this, I will perform the investigation in a water bath and vary the temperature, measure the absorbance after a certain period using the colorimeter and calculate the values of equilibrium constant. This will allow me to have values of equilibrium constant at various values of temperature. A scatter graph can be plotted with the temperature in the x axes and the magnitude of the equilibrium constant along the y axes. The relationship or correlation between these two would allow us to interpret if the reaction is exothermic or endothermic in nature. A positive correlation between the two indicates that the reaction is endothermic while a negative correlation would indicate that the reaction is exothermic.

• Alberty, Robert A. “Effects of PH in Rapid-Equilibrium Enzyme Kinetics.” The Journal of Physical Chemistry B, vol. 111, no. 50, Dec. 2007, pp. 14064–68. ACS Publications, doi:10.1021/jp076742x.

• Chegg.Com. http://https://www.chegg.com/homework-help/questions-and-answers/week-1-chemistry- equilibrium-project-create-calibration-curve-absorbance-vs-concentration--q41730024. Accessed 17 May 2021.

• Missouri University . DETERMINATION OF AN EQUILIBRIUM CONSTANT. Missouri University, https://chemistry.missouri.edu/sites/default/files/class-files/use_det_eq_const_1.pdf. Accessed 26 Mar. 2021.

Part-B: Determining absorbance of Iron-thiocyanate complex at 447 nm at pH = 12.00:

• A 100 cm³ glass beaker (labelled as 0.1 NaOH) was taken.
• Weigh 0.40 g of NaOH and add to the beaker.
• Distilled water was added till the mark of 100 cc. (this solution will be used partly in this part and partly in other steps).
• Take another 100 cc beaker.
• A 10 cm³ graduated pipette was used to transfer 10.00 ± 0.05 cm³ of Iron-III solution to the beaker.
• Another 10 cm³ graduated pipette was used to transfer 10.00 ± 0.05 cm³ of thiocyanate solution.
• A 10.00 cm³ graduated pipette was used to transfer 10.00 cm³ of 0.10 moldm^-3 NaOH to the same beaker.
• Distilled water was added to make the total volume to the mark of 100 cm³.
• The stop-watch was started.
• Switch on the colorimeter.
• Take the cuvette and fill it up with distilled water.
• Set the wavelength of the colorimeter at 447 nm or the value closest to this.
• Insert the cuvette in the colorimeter and set the reading of the absorbance as 0.000 abs using the ‘Set Zero’ button (or the button used to calibrate or adjust the absorbance readings).
• Take out the cuvette.
• Dispose of the water.
• Wipe it with a soft tissue and dry it using a hair-dryer.
• Wait for the stop-watch to read 60.00 ± 0.01 minutes,
• 1.00 ± 0.05 cm³ of the reaction mixture was taken out from the reaction mixture and kept in the cuvette (the test tube or glass tube used in the colorimeter).
• Take the reading of the absorbance of the sample keeping the wavelength at 447 nm.
• Repeat steps 17 & 18 for 4 more times.

Part-C: Determining absorbance of Iron-thiocyanate complex at 447 nm at pH = 11.00:

• A 100 cm³ glass beaker (labelled as 0.01 NaOH) was taken.
• A 10.00 cm³ graduated pipette was used to transfer 10.00 cm³ of 0.10 moldm^-3 NaOH to the same beaker. (from part-B).
• Distilled water was added till the mark of 100 cc (this solution will be used partly in this part and partly in next steps).
• Take another 100 cc beaker.
• A 10 cm³ graduated pipette was used to transfer 10.00 ± 0.05 cm³ of Iron-III solution to the beaker.
• Another 10 cm³ graduated pipette was used to transfer 10.00 ± 0.05 cm³ of thiocyanate solution.
• A 10.00 cm³ graduated pipette was used to transfer 10.00 cm³ of 0.01 moldm^-3 NaOH to the same beaker.
• Distilled water was added to make the total volume to the mark of 100 cm³.
• The stop-watch was started.
• Switch on the colorimeter.
• Take the cuvette and fill it up with distilled water.
• Set the wavelength of the colorimeter at 447 nm or the value closest to this.
• Insert the cuvette in the colorimeter and set the reading of the absorbance as 0.000 abs using the ‘Set Zero’ button (or the button used to calibrate or adjust the absorbance readings).
• Take out the cuvette.
• Dispose of the water.
• Wipe it with a soft tissue and dry it using a hair-dryer.
• Wait for the stop-watch to read 60.00 ± 0.01 minutes,
• 1.00 ± 0.05 cm³ of the reaction mixture was taken out from the reaction mixture and kept in the cuvette (the test tube or glass tube used in the colorimeter).
• Take the reading of the absorbance of the sample keeping the wavelength at 447 nm.
• Repeat steps 17 & 18 for 4 more times.

Part-D: Determining absorbance of Iron-thiocyanate complex at 447 nm at pH = 10.00:

• A 100 cm³ glass beaker (labelled as 0.001 NaOH) was taken.
• A 10.00 cm³ graduated pipette was used to transfer 10.00 cm³ of 0.01 moldm^-3 NaOH to the same beaker. (from part-C).
• Distilled water was added till the mark of 100 cc (this solution will be used partly in this part and partly in next step).
• Take another 100 cc beaker.
• A 10 cm³ graduated pipette was used to transfer 10.00 ± 0.05 cm³ of Iron-III solution to the beaker.
• Another 10 cm³ graduated pipette was used to transfer 10.00 ± 0.05 cm³ of thiocyanate solution.
• A 10.00 cm³ graduated pipette was used to transfer 10.00 cm³ of 0.001 moldm^-3 NaOH to the same beaker.
• Distilled water was added to make the total volume to the mark of 100 cm³.
• The stop-watch was started.
• Switch on the colorimeter.
• Take the cuvette and fill it up with distilled water.
• Set the wavelength of the colorimeter at 447 nm or the value closest to this.
• Insert the cuvette in the colorimeter and set the reading of the absorbance as 0.000 abs using the ‘Set Zero’ button (or the button used to calibrate or adjust the absorbance readings).
• Take out the cuvette.
• Dispose of the water.
• Wipe it with a soft tissue and dry it using a hair-dryer.
• Wait for the stop-watch to read 60.00 ± 0.01 minutes,
• 1.00 ± 0.05 cm³ of the reaction mixture was taken out from the reaction mixture and kept in the cuvette (the test tube or glass tube used in the colorimeter).
• Take the reading of the absorbance of the sample keeping the wavelength at 447 nm.
• Repeat steps 17 & 18 for 4 more times.

Part-E: Determining absorbance of Iron-thiocyanate complex at 447 nm at pH = 9.00:

• A 100 cm³ glass beaker (labelled as 0.0001 NaOH) was taken.
• A 10.00 cm³ graduated pipette was used to transfer 10.00 cm³ of 0.001 moldm^-3 NaOH to the same beaker. (from part-D).
• Distilled water was added till the mark of 100 cc (this solution will be used partly in this part and partly in next step).
• Take another 100 cc beaker.
• A 10 cm³ graduated pipette was used to transfer 10.00 ± 0.05 cm³ of Iron-III solution to the beaker.
• Another 10 cm³ graduated pipette was used to transfer 10.00 ± 0.05 cm³ of thiocyanate solution.
• A 10.00 cm³ graduated pipette was used to transfer 10.00 cm³ of 0.0001 moldm^-3 HCl to the same beaker.
• Distilled water was added to make the total volume to the mark of 100 cm³.
• The stop-watch was started.
• Switch on the colorimeter.
• Take the cuvette and fill it up with distilled water.
• Set the wavelength of the colorimeter at 447 nm or the value closest to this.
• Insert the cuvette in the colorimeter and set the reading of the absorbance as 0.000 abs using the ‘Set Zero’ button (or the button used to calibrate or adjust the absorbance readings).
• Take out the cuvette.
• Dispose of the water.
• Wipe it with a soft tissue and dry it using a hair-dryer.
• Wait for the stop-watch to read 60.00 ± 0.01 minutes,
• 1.00 ± 0.05 cm³ of the reaction mixture was taken out from the reaction mixture and kept in the cuvette (the test tube or glass tube used in the colorimeter).
• Take the reading of the absorbance of the sample keeping the wavelength at 447 nm.
• Repeat steps 17 & 18 for 4 more times.

Part-F: Determining absorbance of Iron-thiocyanate complex at 447 nm at pH = 8.00:

• A 100 cm³ glass beaker (labelled as 0.00001 NaOH) was taken.
• A 10.00 cm³ graduated pipette was used to transfer 10.00 cm³ of 0.0001 moldm^-3 NaOH to the same beaker. (from part-E).
• Distilled water was added till the mark of 100 cc (this solution will be used partly in this part and partly in next step).
• Take another 100 cc beaker.
• A 10 cm³ graduated pipette was used to transfer 10.00 ± 0.05 cm³ of Iron-III solution to the beaker.
• Another 10 cm³ graduated pipette was used to transfer 10.00 ± 0.05 cm³ of thiocyanate solution.
• A 10.00 cm³ graduated pipette was used to transfer 10.00 cm³ of 0.00001 moldm^-3 NaOH to the same beaker.
• Distilled water was added to make the total volume to the mark of 100 cm³.
• The stop-watch was started.
• Switch on the colorimeter.
• Take the cuvette and fill it up with distilled water.
• Set the wavelength of the colorimeter at 447 nm or the value closest to this.
• Insert the cuvette in the colorimeter and set the reading of the absorbance as 0.000 abs using the ‘Set Zero’ button (or the button used to calibrate or adjust the absorbance readings).
• Take out the cuvette.
• Dispose of the water.
• Wipe it with a soft tissue and dry it using a hair-dryer.
• Wait for the stop-watch to read 60.00 ± 0.01 minutes,
• 1.00 ± 0.05 cm³ of the reaction mixture was taken out from the reaction mixture and kept in the cuvette (the test tube or glass tube used in the colorimeter).
• Take the reading of the absorbance of the sample keeping the wavelength at 447 nm.
• Repeat steps 17 & 18 for 4 more times.

Part-G: Determining absorbance of Iron-thiocyanate complex at 447 nm at pH = 7.00:

• A 100 cm³ glass beaker was taken.
• A 10 cm³ graduated pipette was used to transfer 10.00 ± 0.05 cm³ of Iron-III solution to the beaker.
• Another 10 cm³ graduated pipette was used to transfer 10.00 ± 0.05 cm³ of thiocyanate solution.
• Distilled water was added to make the total volume to the mark of 100 cm³.
• The stop-watch was started.
• Switch on the colorimeter.
• Take the cuvette and fill it up with distilled water.
• Set the wavelength of the colorimeter at 447 nm or the value closest to this.
• Insert the cuvette in the colorimeter and set the reading of the absorbance as 0.000 abs using the ‘Set Zero’ button (or the button used to calibrate or adjust the absorbance readings).
• Take out the cuvette.
• Dispose of the water.
• Wipe it with a soft tissue and dry it using a hair-dryer.
• Wait for the stop-watch to read 60.00 ± 0.01 minutes,
• 1.00 ± 0.05 cm³ of the reaction mixture was taken out from the reaction mixture and kept in the cuvette (the test tube or glass tube used in the colorimeter).
• Take the reading of the absorbance of the sample keeping the wavelength at 447 nm.
• Repeat steps 14 & 15 for 4 more times.