UNIT 13B: Investigate oxidation-reduction reactions in order to understand their
many applications in analysis.
Introduction
This report will demonstrate how to determine the concerntration of analytes using
analytical procedures involving oxidation and reduction also demonstrate how to
determine accurate oxdation numbers for species in equations to identify reactions
involving oxidation and reduction and a comparison measured cell voltages for
electrochemical cells inolving metal/metal ion half cells with voltages calculated using
oxidation, reduction and redox equations. Furthermore, explain the redox reactions
involved in analytical procedures in terms of the oxidation numbers for the species
involved. It will also express oxidation, reduction and redox reactions
Index
• Introduction
• Oxidation and reduction
• Half equations and redox reactions
• Standard cell voltages and comparison
• Electrochemical cell half equations and redox reactions
• Balanced redox equations in terms of numbers of electrons
• Standardize solution by titration
• Redox equations involved in each titrations
• Evaluating techniques used in industry
Oxidation and reduction
Using analytes through oxidation-reduction (redox) reactions is a common analytical
technique that is used in chemistry. Oxidatiion is the process of a molecule, ion or itom
losing electrons or increasing its oxidaion state. During this reaction process, a
molecule, ion or atom will either receive or lose an electron, depending on its oxidation
status, Reduction and oxidation, are two chemical processes that involve transferring
electrons between chemical species. Reduction Is the process of gaining electrons
meanwhile oxidation is the process of losing electrons. Redox reactions are important in
a wide range of chemical and industrial processes such as energy production,
corrosion and metabolising living organisms.
Half equations and redox reactions
Task 1: Aldehyde (Clear) and dichromate (Orange) Reaction
What was observed, what oxidized and what reduced?
In the experiment, the substance’s colour changed significantly, from orange to a
greenish blue. This is because when an aldehyde is reacted with dichromate (Cr2O7^2-)
, in the presence of sulfuric acid (H2SO4), an oxidation reaction takes place. The
dichromate is reduced to Chromium (II) ion (Cr^3+) and the aldehyde is oxidized to
carboxylic acid.
Electron half-equation for the reduction of dichromate (VI)
Cr2O7+ 14H + 6e 2Cr + 7HO
Half-equation for the oxidation of an aldehyde
RCHO + HO RCOOH + 2H+ 2e
The overall equation
3RCHO + CrO7 + 8H 3RCOOH + 2Cr + 4HO
Task 2: Electrochemical cells and practical
For the following four electrochemical cells, write the half equations and redox equations
and calculate the standard cell voltage:
2𝐴𝑔(𝑠) → 𝐴𝑔+∥ 𝐶𝑢2(𝑎𝑞) → 𝐶𝑢(𝑠) = 0.46
𝑍𝑛(𝑠) → 𝑍𝑛2 + (𝑎𝑞) ∥ 𝐻 + (𝑎𝑞) → 𝐻2(𝑔) = 0.76
𝑁𝑖(𝑠) → 𝑁𝑖2(𝑎𝑞) ∥ 2𝐹𝑒3 + (𝑎𝑞) → 𝐹𝑒2 + (𝑎𝑞) = 1.04
𝐹𝑒2 + (𝑎𝑞) → 𝐹𝑒3 + (𝑎𝑞) ∥ 𝐴𝑔 + (𝑎𝑞) → 𝐴𝑔(𝑠) = 0.03
Task 3: Balanced redox equations in terms of numbers of electrons
Write out the half equations and redox equations for the following cells:
We used silver nitrate rather than zinc. I changed the equations for this from zinc to
silver nitrate. Which is shown below:
2Aq (s) | 2Aq (aq) Pb (aq) | Pb (s) = - 0.87
2Aq (s) | 2Aq (aq) Cu | Cu (s) = - 0.44
Pb (s) | Pb (aq) | Cu2+ (aq) | Cu (s) = +0.45
Standard results: 0.46, 0.93,0.47
The results were negative due to the polarity being switched.
Task 4: Standardize solution by titration
Standardizing a solution of Fe 2+2+ ions using potassium (vii) manganate in a redox
titration Here is the balanced ionic equation
MnO- 44 + 8H + 5Fe 2+2+ → → Mn 2+2+ + 5Fe 3+3+ + 4H22 O
many applications in analysis.
Introduction
This report will demonstrate how to determine the concerntration of analytes using
analytical procedures involving oxidation and reduction also demonstrate how to
determine accurate oxdation numbers for species in equations to identify reactions
involving oxidation and reduction and a comparison measured cell voltages for
electrochemical cells inolving metal/metal ion half cells with voltages calculated using
oxidation, reduction and redox equations. Furthermore, explain the redox reactions
involved in analytical procedures in terms of the oxidation numbers for the species
involved. It will also express oxidation, reduction and redox reactions
Index
• Introduction
• Oxidation and reduction
• Half equations and redox reactions
• Standard cell voltages and comparison
• Electrochemical cell half equations and redox reactions
• Balanced redox equations in terms of numbers of electrons
• Standardize solution by titration
• Redox equations involved in each titrations
• Evaluating techniques used in industry
Oxidation and reduction
Using analytes through oxidation-reduction (redox) reactions is a common analytical
technique that is used in chemistry. Oxidatiion is the process of a molecule, ion or itom
losing electrons or increasing its oxidaion state. During this reaction process, a
molecule, ion or atom will either receive or lose an electron, depending on its oxidation
status, Reduction and oxidation, are two chemical processes that involve transferring
electrons between chemical species. Reduction Is the process of gaining electrons
meanwhile oxidation is the process of losing electrons. Redox reactions are important in
a wide range of chemical and industrial processes such as energy production,
corrosion and metabolising living organisms.
Half equations and redox reactions
Task 1: Aldehyde (Clear) and dichromate (Orange) Reaction
What was observed, what oxidized and what reduced?
In the experiment, the substance’s colour changed significantly, from orange to a
greenish blue. This is because when an aldehyde is reacted with dichromate (Cr2O7^2-)
, in the presence of sulfuric acid (H2SO4), an oxidation reaction takes place. The
dichromate is reduced to Chromium (II) ion (Cr^3+) and the aldehyde is oxidized to
carboxylic acid.
Electron half-equation for the reduction of dichromate (VI)
Cr2O7+ 14H + 6e 2Cr + 7HO
Half-equation for the oxidation of an aldehyde
RCHO + HO RCOOH + 2H+ 2e
The overall equation
3RCHO + CrO7 + 8H 3RCOOH + 2Cr + 4HO
Task 2: Electrochemical cells and practical
For the following four electrochemical cells, write the half equations and redox equations
and calculate the standard cell voltage:
2𝐴𝑔(𝑠) → 𝐴𝑔+∥ 𝐶𝑢2(𝑎𝑞) → 𝐶𝑢(𝑠) = 0.46
𝑍𝑛(𝑠) → 𝑍𝑛2 + (𝑎𝑞) ∥ 𝐻 + (𝑎𝑞) → 𝐻2(𝑔) = 0.76
𝑁𝑖(𝑠) → 𝑁𝑖2(𝑎𝑞) ∥ 2𝐹𝑒3 + (𝑎𝑞) → 𝐹𝑒2 + (𝑎𝑞) = 1.04
𝐹𝑒2 + (𝑎𝑞) → 𝐹𝑒3 + (𝑎𝑞) ∥ 𝐴𝑔 + (𝑎𝑞) → 𝐴𝑔(𝑠) = 0.03
Task 3: Balanced redox equations in terms of numbers of electrons
Write out the half equations and redox equations for the following cells:
We used silver nitrate rather than zinc. I changed the equations for this from zinc to
silver nitrate. Which is shown below:
2Aq (s) | 2Aq (aq) Pb (aq) | Pb (s) = - 0.87
2Aq (s) | 2Aq (aq) Cu | Cu (s) = - 0.44
Pb (s) | Pb (aq) | Cu2+ (aq) | Cu (s) = +0.45
Standard results: 0.46, 0.93,0.47
The results were negative due to the polarity being switched.
Task 4: Standardize solution by titration
Standardizing a solution of Fe 2+2+ ions using potassium (vii) manganate in a redox
titration Here is the balanced ionic equation
MnO- 44 + 8H + 5Fe 2+2+ → → Mn 2+2+ + 5Fe 3+3+ + 4H22 O
