Experiment 1: Acids and Bases
Anne Versluijs, CHU11B01 23/24, 02/11/23
Introduction
This experiment consists of two parts. The aim of the first part is to determine the pH of different
aqueous solutions using multiple indicators. The objective of part two of the experiment is to
determine the dissociation constant of three different acid-base reactions: weak acid with strong base,
strong acid with strong base, and strong base with weak acid.
Part 1: pH of aqueous solutions
A solution is said to be acidic (pH<7) when the concentration of H+ is bigger than the OH-
concentration, whereas in a basic solution (pH>7) the OH- concentration is larger than the H+
concentration. An acid-base indicator helps determine where an aqueous solution falls on the pH scale,
quantifying the acidity of the solution. The indicator, a weak acid, has two different colours: one in its
acidic form and one in its basic form [1].
𝐻(𝐼𝑛𝑑𝑖𝑐𝑎𝑡𝑜𝑟) + 𝐻2 𝑂 ⇌ 𝐻3 𝑂+ + 𝐼𝑛𝑑𝑖𝑐𝑎𝑡𝑜𝑟 −
In an acidic solution [H(Indicator)] > [Indicator -], and the solution will show the colour of the
indicator’s acidic form.
Bromothymol blue is used to measure the pH of solutions with a relatively neutral pH. It has a yellow
colour below a pH of 6,0 and is blue above pH 7,6 [2].
Figure 1: solutions of bromothymol blue at different pH.[3]
Methyl orange is usually used in the titration of strong acids with weak bases, because, around the
stoichiometric point, the colour change will start to appear. Methyl orange has a red colour below pH
3,1 and a yellow colour above pH 4,4 [2].
Phenolphthalein is mostly used in the titration of weak acids with strong bases because of the colour
change from pink to colourless at around pH 8,4 [2].
Figure 2: colour of methyl orange and phenolphthalein at different pH.[2]
, Part 2: dissociation constant of a weak acid / weak base
A strong acid (or base) fully dissociates in water. A common strong acid is HCl:
𝐻𝐶𝑙 + 𝐻2 𝑂 ⟶ 𝐶𝑙 − + 𝐻3 𝑂+
A weak acid (or base) is only partially dissociated, resulting in an equilibrium reaction. A common
weak acid is CH3COOH:
𝐶𝐻3 𝐶𝑂𝑂𝐻 + 𝐻2 𝑂 ⇌ 𝐶𝐻3 𝐶𝑂𝑂− + 𝐻3 𝑂+
During a strong base – strong acid titration, all H+ will have reacted with the base at equilibrium to
form H2O (pH = 7). During a weak acid – strong base titration (where the titrant is the strong base),
the dissociation equilibrium will have settled in the analyte, resulting in both the acid and its conjugate
base. Because protons will transfer only from the strongest acid to the strongest base, the OH- ions will
react with the formed H3O+ first. The system will restore the equilibrium by dissociating more weak
acid. At equilibrium, all of the weak acid will have dissociated, and all H3O+ will have reacted with
OH-, leaving only the conjugate base of the weak acid and the conjugate acid of the strong base in the
solution. Meaning the pH at equilibrium is larger than 7. The pH at equilibrium in a strong acid–weak
base titration will be lower than 7, following the same logic [1]. At the half-stoichiometric point, the
concentration of the conjugate base is equal to the concentration of the weak acid.
Experimental procedure
Both methods are based on the method in the ‘Junior Fresh Practical Chemistry for Life and Health
Sciences’ lab manual [1].
Part 1: pH of aqueous solutions
Using a Pasteur pipette approximately 1 cm3 of dilute HCl (2M) was added to a clean test tube. pH
paper was inserted into the tilted test tube until the bottom of the paper was just in the solution. The
paper was removed from the solution after one second and the colours on the strip were compared to
the colour chart.
This method was repeated for the six other solutions: neutral buffer solution (pH 7), dilute NaOH
(2M), AlCl3 (aq), KNO3 (aq), CH3NH3Cl (aq), and NaHCO3 (aq).
1 drop of bromothymol blue indicator was added to all seven test tubes. Any colour change that
occurred after shaking the solutions was observed.
The test tubes were washed and refilled with the seven aforementioned solutions. The experiment was
repeated using methyl orange and phenolphthalein.
Part 2: dissociation constant of a weak acid / weak base
0,1 M HCl was diluted 1:2 with deionized water using a pipette, creating a 50 mL 0,05 M HCl
solution. The pre-calibrated electrode was washed with distilled water. The electrode and a magnetic
stirrer were added to a 150 mL beaker containing the HCl solution. The burette was filled with 0,1 M
NaOH. After each addition of 1 mL NaOH, the pH of the solution was read off the display of the
electrode.
This method was repeated with 0,1 M CH3COOH instead of HCl while titrating with the same NaOH
solution in the burette.
Then the method was once more repeated using a 0,1 M NH3 solution instead of the HCl solution and
switching the solution in the burette to 0,1 M HCl.
Anne Versluijs, CHU11B01 23/24, 02/11/23
Introduction
This experiment consists of two parts. The aim of the first part is to determine the pH of different
aqueous solutions using multiple indicators. The objective of part two of the experiment is to
determine the dissociation constant of three different acid-base reactions: weak acid with strong base,
strong acid with strong base, and strong base with weak acid.
Part 1: pH of aqueous solutions
A solution is said to be acidic (pH<7) when the concentration of H+ is bigger than the OH-
concentration, whereas in a basic solution (pH>7) the OH- concentration is larger than the H+
concentration. An acid-base indicator helps determine where an aqueous solution falls on the pH scale,
quantifying the acidity of the solution. The indicator, a weak acid, has two different colours: one in its
acidic form and one in its basic form [1].
𝐻(𝐼𝑛𝑑𝑖𝑐𝑎𝑡𝑜𝑟) + 𝐻2 𝑂 ⇌ 𝐻3 𝑂+ + 𝐼𝑛𝑑𝑖𝑐𝑎𝑡𝑜𝑟 −
In an acidic solution [H(Indicator)] > [Indicator -], and the solution will show the colour of the
indicator’s acidic form.
Bromothymol blue is used to measure the pH of solutions with a relatively neutral pH. It has a yellow
colour below a pH of 6,0 and is blue above pH 7,6 [2].
Figure 1: solutions of bromothymol blue at different pH.[3]
Methyl orange is usually used in the titration of strong acids with weak bases, because, around the
stoichiometric point, the colour change will start to appear. Methyl orange has a red colour below pH
3,1 and a yellow colour above pH 4,4 [2].
Phenolphthalein is mostly used in the titration of weak acids with strong bases because of the colour
change from pink to colourless at around pH 8,4 [2].
Figure 2: colour of methyl orange and phenolphthalein at different pH.[2]
, Part 2: dissociation constant of a weak acid / weak base
A strong acid (or base) fully dissociates in water. A common strong acid is HCl:
𝐻𝐶𝑙 + 𝐻2 𝑂 ⟶ 𝐶𝑙 − + 𝐻3 𝑂+
A weak acid (or base) is only partially dissociated, resulting in an equilibrium reaction. A common
weak acid is CH3COOH:
𝐶𝐻3 𝐶𝑂𝑂𝐻 + 𝐻2 𝑂 ⇌ 𝐶𝐻3 𝐶𝑂𝑂− + 𝐻3 𝑂+
During a strong base – strong acid titration, all H+ will have reacted with the base at equilibrium to
form H2O (pH = 7). During a weak acid – strong base titration (where the titrant is the strong base),
the dissociation equilibrium will have settled in the analyte, resulting in both the acid and its conjugate
base. Because protons will transfer only from the strongest acid to the strongest base, the OH- ions will
react with the formed H3O+ first. The system will restore the equilibrium by dissociating more weak
acid. At equilibrium, all of the weak acid will have dissociated, and all H3O+ will have reacted with
OH-, leaving only the conjugate base of the weak acid and the conjugate acid of the strong base in the
solution. Meaning the pH at equilibrium is larger than 7. The pH at equilibrium in a strong acid–weak
base titration will be lower than 7, following the same logic [1]. At the half-stoichiometric point, the
concentration of the conjugate base is equal to the concentration of the weak acid.
Experimental procedure
Both methods are based on the method in the ‘Junior Fresh Practical Chemistry for Life and Health
Sciences’ lab manual [1].
Part 1: pH of aqueous solutions
Using a Pasteur pipette approximately 1 cm3 of dilute HCl (2M) was added to a clean test tube. pH
paper was inserted into the tilted test tube until the bottom of the paper was just in the solution. The
paper was removed from the solution after one second and the colours on the strip were compared to
the colour chart.
This method was repeated for the six other solutions: neutral buffer solution (pH 7), dilute NaOH
(2M), AlCl3 (aq), KNO3 (aq), CH3NH3Cl (aq), and NaHCO3 (aq).
1 drop of bromothymol blue indicator was added to all seven test tubes. Any colour change that
occurred after shaking the solutions was observed.
The test tubes were washed and refilled with the seven aforementioned solutions. The experiment was
repeated using methyl orange and phenolphthalein.
Part 2: dissociation constant of a weak acid / weak base
0,1 M HCl was diluted 1:2 with deionized water using a pipette, creating a 50 mL 0,05 M HCl
solution. The pre-calibrated electrode was washed with distilled water. The electrode and a magnetic
stirrer were added to a 150 mL beaker containing the HCl solution. The burette was filled with 0,1 M
NaOH. After each addition of 1 mL NaOH, the pH of the solution was read off the display of the
electrode.
This method was repeated with 0,1 M CH3COOH instead of HCl while titrating with the same NaOH
solution in the burette.
Then the method was once more repeated using a 0,1 M NH3 solution instead of the HCl solution and
switching the solution in the burette to 0,1 M HCl.
