Reading assignment:  Chang, Chemistry 10^th edition, Chapter 15:  Acids and Bases, sections 1-5.

 

Goals

To determine the acid dissociation constant (K_a) for bromocresol green (BCG), an acid-base indicator.

 

Discussion Acid-base indicators are often used to demonstrate the end-point of an acid-base reaction.  Examples include phenolphthalein and the mixture of indicators used in universal indicator solution.  Acid-base indicators are weak acids that dissociate into a hydronium ion (H₃O⁺) and a conjugate base anion (In^–).   This dissociation can be represented through the following equation and equilibrium expression:

HIn (aq)  +  H₂O (l)   ?  H₃ O⁺ (aq)  +  In^- (aq)        K_a= [H₃ O⁺ ][In^- ] weakacid hydronium conjugatebase [ HIn ] indicator                  ion indicator

In order for a compound to be a useful indicator, the acidic form (HIn) and the basic form (In^–) of the indicator should differ in color.  Since equilibrium in acidic solution favors the formation of HIn, this species is called the acidic form of the indicator.  Likewise, the In^– form is called the basic form since it is favored in basic solutions.  An equilibrium mixture of the indicator will be colored according to the relative concentration of each form of the indicator.  The position of the equilibrium and, therefore, the relative concentration of the two forms of the indicator will depend on the H₃O⁺ concentration, [H₃O⁺] or in shorthand notation [H⁺].

The absorption curve of an indicator at different pH values can be studied to determine the equilibrium constant of the indicator.  In this experiment, we will determine the equilibrium constant of bromocresol green (BCG).  BCG is an indicator that is yellow in acidic solutions blue in basic solutions.  When dissolved in water the conjugate pair (acidic and basic forms) display different absorption spectra since they possess

 

Department of Physical Sciences Kingsborough Community College The City University of New York Spring 2012

Experiment 5:  Determination of the Equilibrium Constant for Bromocresol Green

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For simplicity we'll use the following symbols to represent bromocresol green:

 

H₂B      The protonated form of the indicator

HB^–      The deprotonated form of the indicator

NaHB The sodium salt of the deprotonated form of the indicator

B^2–       The fully deprotonated form of the indicator

 

The sodium salt of bromocresol green ionizes completely in water:

NaHB (aq)   _→   Na⁺ (aq)  +  HB^– (aq)

 

The HB^– form, which is a monoprotic acid, then partially dissociates to give B^–2: HB^– (aq)  +  H₂O (l)  ?  H₃O⁺ (aq)   +  B^2– (aq) 

Writing the dissociation without water (shorthand notation), we have

HB– (aq) ? H⁺(aq) + B2-(aq)        K=

 

 

acidicform(yellow)       basicform(blue)

HB^– is the acidic form and is yellow in solution.  B^2– is the basic form and is blue in solution.  

Taking logarithms of the above equation gives

logK=log [H⁺   ]+log

 

 

 

where we have rearranged and noted that pH = –log [H⁺].

This is a linear equation.  A plot of log

 versus pH should yield a straight line with a slope of 1 (m

 

= 1) and an intercept equal to log K, where K is a concentration equilibrium constant.  So our strategy will be to measure the ratio log

 as a function of pH and use this data to determine the equilibrium constant for the dissociation of bromocresol green.

 

 

 

Absorbance and Spectrophotometry

Solutions that possess colors absorb visible light energy of specific wavelengths.  Recall that a red solution appears red because it absorbs much of the blue-green part of the spectrum (complementary colors).  Measurements of the amount of light absorbed by a substance at each wavelength (color) can be graphed giving an “absorption curve.”  The shape of this curve depends almost entirely on the electronic structure of the substance and is almost unique for each substance.  Thus the curve serves as an aid to identification and, with the aid of modern theory, a clue to the structure of a substance.

At a given wavelength the amount of light absorbed by a solute is proportional to its molar concentration, thus providing a widely used method of concentration analysis.  The Beer-Lambert Law states that A = εlc, where A = absorbance, _ε  = a constant characteristic of the absorbing molecule, l = path length, c = concentration.  In our case, _ε and l are each constant (known absorbing substance and a path length determined by the width of the cuvette).  Thus the absorbance is proportional to the concentration in this experiment.

A spectrophotometer is an instrument used for measuring the amount of absorption at different wavelengths.  Many such instruments are now commercially available.  Some are designed to operate in the region of visible light, some in infra-red regions, some in ultra-violet and still others in several of these energy bands.  You will use either a Spectronic 301 Spectrophotometer or a Genesys 20

Spectrophotometer to measured absorbance in this experiment.  Both instruments are designed for use in the visible portion of the spectrum (400 to 700 nanometers, nm).  

Determining the Equilibrium Constant

If the acidic form (HB^–) and basic form (B^2–) of an indicator both absorb light, then the ratio

 

 

can be determined by measuring the absorption of light at the wavelength at which one of the forms absorbs a lot and the other form absorbs a little.

Plots of the absorption of light of various wavelengths (λ) versus wavelength for bromocresol green in its basic form and in its acidic form are shown in the chart to the right.  Such graphs are called absorption spectra. 

 

The graph shows that the basic form has maximum absorption (A_II) at λ_max.  The absorbance of the acidic form (A_I) is small at λ_max (almost zero). 

If we start with a solution of the pure basic form (B^2–) and add an acid to the solution (for example, acetic acid), then some of the basic form will be converted to the acid form (HB^–).  Then the absorbance at λ_max will drop because the acidic form absorbs almost no light at λ_max.  This change may be visible in that the solution will change from blue to yellow.

of Physical Sciences Kingsborough Community College The City University of New York

Experiment 6:  Determination of the Equilibrium Constant for Bromocresol Green

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We will call an absorbance measurement with a mixture of the acidic and basic forms of the indicator A_x.  In fact, all absorbance values along the vertical line drawn at λ_max in the graph are A_x values.  These points all lie between A_II and A_I.

Referring back to the equation, log

=pH+logK , we see that in order to calculate the equilibrium

 

constant, K, different ratios of

 must be found for different pH values.  When these are plotted, a straight line with slope of one and y-intercept of log K is obtained.  The relationship between the absorbance data at λ_max and the ratio

 is found from the following equation:

 

[B2--]]= AAIIx- -A Ax I  where

[HB

A_I  = absorbance of the solution when all indicator is in the acidic form (pure HIn).

A_II = absorbance of the solution when all indicator is in the basic form (pure In^–).

A_x  = absorbance of the solution when both some acidic and some basic form are present

 

Procedure

Equipment:   UV-vis spectrophotometer, two cuvettes, pipette pump, three 250-mL beakers, 100-mL graduated cylinder, 2-mL pipet, 10-mL graduated cylinder, Kimwipes, USB flash drive.  Chemicals: 1.00 M acetic acid (HAc), 3.00 x 10^-4 M bromocresol green (BCG), 0.200 M sodium acetate (NaAc), cuvette containing acidic form of BCG obtained from your instructor.

Use eye protection for this experiment.

Students will work in groups as assigned by the instructor.  

Part 1:  Absorbance Spectrum of Bromocresol Green in Basic and Acidic Solutions

1.  We will need four solutions to perform this experiment.  The instructions for making the first three are given in the text below.  The fourth solution has already been prepared and will be distributed by your instructor in a cuvette.  All solutions should be prepared in a graduated cylinder since volume markings on beakers are good to only 5%.

1. Bromocresol green in sodium acetate solution:  Add 5.00 ml of 3 x 10^–4 M bromocresol green solution and 5.00 mL of 0.200 M sodium acetate (NaAc) solution to a 100-mL graduated cylinder.  Dilute to 100.0 mL with deionized water and pour quantitatively (this means all of the solution) into a clean, dry 250-mL beaker.  This solution should be blue and will be referred to as the basic solution.  The sodium acetate concentration in this solution is 0.0100 M.
2. Acetic acid solution:  Rinse the graduated cylinder with distilled water, add 25.0 mL of the 1.00 M acetic acid (HAc) solution and dilute to 100.0 mL with deionized water.  Pour this solution quantitatively into a clean, dry 250-mL beaker.  The concentration of the acetic acid in this solution is 0.250 M.
3. Blank Solution:  Add 5 mL of 1.00 M acetic acid and 5 mL of 0.200 M sodium acetate to a 100 mL graduated cylinder and dilute with deionize water to 100 mL.
4. Bromocresol green solution in hydrochloric acid:  The instructor will supply a cuvette containing a BCG solution that has excess acid (HCl) added to it.  This solution should be yellow.

1. Measure the absorption spectrum of the solution containing BCG in sodium acetate (basic BCG solution) and the BCG solution containing HCl (acidic BCG solution) beginning at 380 nm and ending at 680 nm at 20 nm intervals.  Record these measurements in the Data Sheet.

of Physical Sciences Kingsborough Community College The City University of New York Experiment 5:  Determination of the Equilibrium Constant for Bromocresol Green

 

2. Review the absorbance data and locate the greatest value of the absorbance for the basic solution.  Note what wavelength this absorbance corresponds to.  Now measure the absorbance of the basic BCG solution and the acidic BCG solution at 5 nm intervals from 20 nm below to 20 nm above this wavelength.  For example, if the wavelength of maximum absorbance occurs at 520 nm, make measurements at 5 nm intervals from 500 nm to 540 nm.  However, don’t duplicate readings—in this case don’t take readings at 500 nm, 520 nm, and 540 nm.  
3. Pour the BCG in sodium acetate solution in the cuvette back into the 250-mL beaker, taking care not to lose any of the solution. 
4. Return the cuvette containing the acidic form of BCG to your instructor.
5. From the data taken for the basic and acidic forms, decide which wavelength is suitable for measuring A_x values and set the spectrophotometer at this wavelength.  The suitable wavelength, λ_max, is the wavelength in which the basic form has a maximum absorbance and the acidic form has a very small absorbance. 

Part 2:  Effect of pH on Absorbance of Bromocresol Green

1. Add precisely 2.00 mL of the 0.250 M acetic acid solution to the 250-mL beaker of BCG solution using a 2 mL pipet and a pipette pump.  Mix the solution with a glass stirring rod and measure the absorbance at λ_max.  Pour the solution in the cuvette back into the 250-mL beaker.  Don’t throw the solution in the cuvette out.
2. Add another 00 mL of the 0.250 M acetic acid solution, mix well, and measure the absorbance at λ_max.  Again, pour the solution in the cuvette back into the beaker.  

 

3. Repeat this procedure with a third and fourth 2.00 mL increment of the 0.250 M acetic acid solution.  There should now be a total of 8.00 mL of 0.250 M acetic acid solution added to the BCG solution.

 

4. Rinse all solutions down the drain and clean any glassware used in the experiment before performing any calculations.