Unit 13a: Applications of Inorganic Chemistry
Introduction:
The purpose of this assignment is to create a report that demonstrates the process of evaluating the
accuracy of acid-base titrations using different methods, including an indicator, a pH meter, and an
auto titrator. Additionally, the report will cover the use of complex calculations to rearrange
equations and determine the pH of solutions. The report will also explain the process of selecting
appropriate indicators or titrators and demonstrate the use of a pH meter to measure the Ka value
of a weak acid and the action of a buffer solution. Finally, the report will demonstrate accurate use
of a pH meter in order to choose suitable indicators.
What is a pH meter?
A pH meter is a scientific instrument that measures the acidity or basicity (alkalinity) of a solution. It
consists of a combination of a probe (typically a glass electrode) and a meter that displays the pH
value.
To use a pH meter, first make sure the probe is properly calibrated. This is typically done by
immersing the probe in two or three buffer solutions of known pH and adjusting the meter to match
the known values.
Next, immerse the probe in the solution you want to test. Stir the solution gently to ensure that the
electrode is in contact with a representative sample of the solution. Allow the reading to stabilize
before taking note of the pH value.
In order to select a suitable indicator for a given acid-base titration, you must know the pH range of
the indicator and the expected endpoint pH of the titration.
Indicators are chemical compounds that change colour at certain pH ranges to signal the endpoint of
a titration. For example, phenolphthalein changes colour at around pH 8.2 (pink in acidic solutions,
colourless in basic solutions), while methyl orange changes colour at around pH 3.1 (red in acidic
solutions, yellow in basic solutions). Therefore, when titrating an acidic solution with a strong base,
phenolphthalein is a better indicator than methyl orange, because the endpoint will be close to the
pH at which phenolphthalein changes colour.
It's important to note that even though pH meters are precise instruments, indicators are commonly
used because most chemical reactions in a lab are not fast or accurate enough to measure pH
changes for an accurate endpoint detection. Also, indicators are much cheaper than pH meters and
can be used for applications where pH meters can't be used such as presence of interference (e.g
solids, turbidity)
Calculations to find the pH of H+ ions and acid dissociation constants (Ka)
Strong acid:
1. A sample of HCl has a concentration of 1.22x10^-3 mol/dm^3, What is the pH?
, HCl(aq) -> H+(aq) + Cl(aq) - Dissociates fully
[H+] = [HCl] = 1.22x10^-3 mol/dm^-3
pH = -log[H+]
= -log(1.22x10^-3)mol/dm^-3
= 2.91
2. What is the concentration of HNO3, if it has a pH of 5.63
HNO3(aq) -> H+(aq) + NO3- - Dissociates fully
[H+] = [HNO3]
[H+] = 10^-pH
pH = 5.63
= 10^-5.63
= 2.34x10^-6 mol/dm^-3
Concentration of HNO3 = 2.34x10^-6 mol/dm^-3
Strong base:
The Kw is 1x10^-4. What is the pH of 0.1 mol/dm^-3 of KOH?
KOH(aq) -> K+(aq) + OH-(aq)
[H+] [OH-] = 1x10^-14
[H+] = 1x10^-.1 = 1x10^-13
= -log(1x10^-13) = 13
Weak acid:
What is the pH when the concentration of nitrous acid (HNO2) is 0.055mol/dm^-3 at 25°C?
HNO2(aq) <--> H+(aq) + NO2(aq) - Partial dissociation
Ka = [H+]^2 / [HNO2]
Ka = 4.70x10^-4 mol/dm^-3
[H+]^2 = Ka[HNO2]
[H+] = √(Ka[HNO2])
[H+] = √(4.70x10^-4 x 0.055)
[H+] = 5.08x10^-3 mol/dm^-3
Introduction:
The purpose of this assignment is to create a report that demonstrates the process of evaluating the
accuracy of acid-base titrations using different methods, including an indicator, a pH meter, and an
auto titrator. Additionally, the report will cover the use of complex calculations to rearrange
equations and determine the pH of solutions. The report will also explain the process of selecting
appropriate indicators or titrators and demonstrate the use of a pH meter to measure the Ka value
of a weak acid and the action of a buffer solution. Finally, the report will demonstrate accurate use
of a pH meter in order to choose suitable indicators.
What is a pH meter?
A pH meter is a scientific instrument that measures the acidity or basicity (alkalinity) of a solution. It
consists of a combination of a probe (typically a glass electrode) and a meter that displays the pH
value.
To use a pH meter, first make sure the probe is properly calibrated. This is typically done by
immersing the probe in two or three buffer solutions of known pH and adjusting the meter to match
the known values.
Next, immerse the probe in the solution you want to test. Stir the solution gently to ensure that the
electrode is in contact with a representative sample of the solution. Allow the reading to stabilize
before taking note of the pH value.
In order to select a suitable indicator for a given acid-base titration, you must know the pH range of
the indicator and the expected endpoint pH of the titration.
Indicators are chemical compounds that change colour at certain pH ranges to signal the endpoint of
a titration. For example, phenolphthalein changes colour at around pH 8.2 (pink in acidic solutions,
colourless in basic solutions), while methyl orange changes colour at around pH 3.1 (red in acidic
solutions, yellow in basic solutions). Therefore, when titrating an acidic solution with a strong base,
phenolphthalein is a better indicator than methyl orange, because the endpoint will be close to the
pH at which phenolphthalein changes colour.
It's important to note that even though pH meters are precise instruments, indicators are commonly
used because most chemical reactions in a lab are not fast or accurate enough to measure pH
changes for an accurate endpoint detection. Also, indicators are much cheaper than pH meters and
can be used for applications where pH meters can't be used such as presence of interference (e.g
solids, turbidity)
Calculations to find the pH of H+ ions and acid dissociation constants (Ka)
Strong acid:
1. A sample of HCl has a concentration of 1.22x10^-3 mol/dm^3, What is the pH?
, HCl(aq) -> H+(aq) + Cl(aq) - Dissociates fully
[H+] = [HCl] = 1.22x10^-3 mol/dm^-3
pH = -log[H+]
= -log(1.22x10^-3)mol/dm^-3
= 2.91
2. What is the concentration of HNO3, if it has a pH of 5.63
HNO3(aq) -> H+(aq) + NO3- - Dissociates fully
[H+] = [HNO3]
[H+] = 10^-pH
pH = 5.63
= 10^-5.63
= 2.34x10^-6 mol/dm^-3
Concentration of HNO3 = 2.34x10^-6 mol/dm^-3
Strong base:
The Kw is 1x10^-4. What is the pH of 0.1 mol/dm^-3 of KOH?
KOH(aq) -> K+(aq) + OH-(aq)
[H+] [OH-] = 1x10^-14
[H+] = 1x10^-.1 = 1x10^-13
= -log(1x10^-13) = 13
Weak acid:
What is the pH when the concentration of nitrous acid (HNO2) is 0.055mol/dm^-3 at 25°C?
HNO2(aq) <--> H+(aq) + NO2(aq) - Partial dissociation
Ka = [H+]^2 / [HNO2]
Ka = 4.70x10^-4 mol/dm^-3
[H+]^2 = Ka[HNO2]
[H+] = √(Ka[HNO2])
[H+] = √(4.70x10^-4 x 0.055)
[H+] = 5.08x10^-3 mol/dm^-3
