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Deciding If a Reaction is a Redox Reaction
Consider the following reaction:
Na + F → Na+ + F-
How do we know that this is a redox reaction?
If this is a redox reaction then:
- One species in the reaction must lose electrons (be oxidised)
The oxidation state (oxidation number) of this element will increase.
- One species in the reaction must gain electrons (be reduced)
The oxidation state (oxidation number) of this element will decrease.
Let's look at what has happend to the sodium atom and fluorine atom in this reaction:
- Sodium atom (Na) has lost an electron to form a positively charged sodium ion (Na+):
Na → Na+ e-
⚛ Na has been oxidized to Na+ since oxidation refers to the loss of electrons.
⚛ The oxidation number of sodium has increased from 0 for Na to +1 for Na+.
⚛ Na is the reducing agent (reductant), it has caused F to be reduced.
- Fluorine atom (F) has gained an electron to form a negatively charged fluoride ion (F-):
F + e- → F-
⚛ F has been reduced to F- since reduction refers to the gain of electrons.
⚛ The oxidation number of fluorine has decreased from 0 for F to -1 for F-.
⚛ F is the oxidizing agent (oxidant), it has caused Na to be oxidized.
We therefore know that this is a redox reaction because an electron has been transferred from the sodium atom to the flourine atom.
Worked Examples: Redox Reaction Concepts
Question 1:
Identify the reductant in the redox reaction given below:
2Ag+(aq) + Sn(s) → Sn2+(aq) + 2Ag(s)
Justify your answer.
Solution:
(Based on the StoPGoPS approach to problem solving.)
- What is the question asking you to do?
(i) Identify the reductant.
(ii) Justify your answer.
- What data (information) have you been given in the question?
Extract the data from the question:
redox reaction equation:
2Ag+(aq) + Sn(s) → Sn2+(aq) + 2Ag(s)
- What is the relationship between what you know and what you need to find out?
Definition: a reductant is a species that causes something else to the reduced.
A reductant is itself oxidised (oxidized).
The oxidation number (oxidation state) of a reductant will increase.
Step 1: Determine the oxidation number (state) for each element in the redox reaction:
| redox reaction: |
2Ag+(aq) |
+ |
Sn(s) |
→ |
Sn2+(aq) |
+ |
2Ag(s) |
|---|
oxidation states
(oxidation numbers) |
+1 |
|
0 |
|
+2 |
|
0 |
|---|
Note: oxidation state of an element in its standard state is 0 (oxidation state of Sn in Sn(s) is 0, and of Ag in Ag(s) is 0)
Oxidation state of metal cation is equal to the charge on the cation.
Step 2: Compare the oxidation state (oxidation number) of each element on the left hand side of the equation to the oxidation state of the same element on the right hand side of the equation:
Oxidation state (oxidation number) for silver has decreased from +1 to 0.
Oxidation state (oxidation number) for tin has increased from 0 to +2.
- Identify the reductant
The oxidation number of the reductant will have increased.
Oxidation number (state) for tin has increased from 0 to +1.
Sn(s) has been oxidised (oxidized) to Sn2+(aq)
The species that is oxidised (Sn(s)) will cause another species (Ag+(aq)) to be reduced, so the species that is oxidised is referred to as a reductant (or a reducing agent).
Reductant is therefore Sn(s).
- Is your answer plausible?
A redox reaction is a reaction in which electrons are transferred from one species (reductant or reducing agent) to another species (oxidant or oxidising agent).
An atom of Sn(s) has lost 2 electrons to form the tin(II) ion (Sn2+(aq))
Sn(s) → Sn2+(aq) + 2e-
An ion of silver (Ag+(aq)) gains an electron to produce an atom of silver (Ag(s)):
Ag+(aq) + e- → Ag(s)
Electrons have been transferred from Sn(s) to Ag+(aq).
Sn(s) has caused Ag+(aq) to be reduced, so Sn(s) must be the reductant (reducing agent).
- State your solution to the problem "identify the reductant and justify your answer":
(i) Identify the reductant: Reductant is Sn(s)
(ii) Justification: Sn(s) is the reductant because it has caused the silver ions, Ag+(aq), to be reduced.
Sn(s) has been oxidised to Sn2+(aq) by losing 2 electrons.
The species that is oxidised is the reductant.
Silver ions, Ag+(aq), have been reduced to silver atoms, Ag(s), by gaining an electron.
The species that is reduced is the oxidant.
Question 2:
Is the reaction below a redox reaction?
Hg2+(aq) + S2-(aq) → HgS(s)
Justify your answer.
Solution:
(Based on the StoPGoPS approach to problem solving.)
- What is the question asking you to do?
(i) Decide if the reaction is a redox reaction.
(ii) Justify your answer.
- What data (information) have you been given in the question?
Extract the data from the question:
Balanced chemical equation for the reaction:
Hg2+(aq) + S2-(aq) → HgS(s)
- What is the relationship between what you know and what you need to find out?
In a redox reaction one species is oxidised (the reductant or reducing agent) causing another species to be reduced (the oxidant or oxidising agent).
Oxidation number (oxidation state) of the reductant increases.
Oxidation number (oxidation state) of the oxidant decreases.
Step 1: Determine the oxidation number (oxidation state) of each element in the equation:
| balanced equation: |
Hg2+(aq) |
+ |
S2-(aq) |
→ |
HgS(s) |
|---|
oxidation states
(oxidation numbers) |
+2 |
|
-2 |
|
+2 (Hg2+)
-2 (S2-) |
|---|
Step 2: Compare the oxidation state (oxidation number) of each element on the left hand side of the equation to the oxidation state of the same element on the right hand side of the equation:
Oxidation state (oxidation number) for mercury (Hg) has not changed, it has remained +2.
Oxidation state (oxidation number) for sulfur (S) has not changed, it has remained -2.
- Determine whether the reaction is a redox reaction and why.
If the reaction is a redox reaction there will be a transfer of electrons from one species to another which will be apparent as a change in the oxidation state (oxidation number) of the species taking part in the reaction.
The oxidation state (oxidation number) of mercury has not changed, neither has the oxidation state (oxidation number) of sulfur changed, so this is NOT a redox reaction.
- Is your answer plausible?
The reaction looks like a simple precipitation reaction in which an aqueous solution containing mercury(II) ions is mixed with an aqueous solution containing sulfide ions to produce a precipitate of mercury(II) sulfide made up of mercury(II) ions and sulfide ions.
If we refer to the solubility rules we find that all sulfides are insoluble except those of Group 1, Group 2 and the ammonium ion.
Mercury is a transition metal, therefore we expect the sulfide of mercury to be insoluble and form a precipitate.
Note that there is NO transfer of electrons from one species to another in a precipitation reaction, so it is not a redox reaction, and we are reasonably confident that our answer is correct.
- State your solution to the problem "is this a redox reaction and justify your answer":
(i) The reaction is NOT a redox reaction.
(ii) Justification:
The oxidation number (oxidation state) for mercury and sulfur have not changed.
Neither mercury nor sulfur have been oxidised.
Neither mercury nor sulfur have been reduced.
There has been no transfer of electrons from mercury to sulfur, nor from sulfur to mercury, so this is NOT a redox reaction.
Question 3:
Is the reaction below a redox reaction?
2H2O2(aq) → 2H2O(l) + O2(g)
Justify your answer.
Solution:
(Based on the StoPGoPS approach to problem solving.)
- What is the question asking you to do?
(i) Decide if the reaction is a redox reaction.
(ii) Justify your answer.
- What data (information) have you been given in the question?
Extract the data from the question:
balanced chemical reaction:
2H2O2(aq) → 2H2O(l) + O2(g)
- What is the relationship between what you know and what you need to find out?
In a redox reaction one species is oxidised (the reductant or reducing agent) causing another species to be reduced (the oxidant or oxidising agent).
Oxidation number (oxidation state) of the reductant increases.
Oxidation number (oxidation state) of the oxidant decreases.
Step 1: Determine the oxidation number (oxidation state) of each element in the equation:
| balanced equation: |
2H2O2(aq) |
→ |
2H2O(l) |
+ |
O2(g) |
|---|
oxidation states
(oxidation numbers) |
+1 (H)
-1 (O) |
|
+1 (H)
-2 (O) |
|
0 (O) |
|---|
Note: oxidation state of an element in its standard state is 0 (oxidation state of O in O2(g) is 0)
Oxidation state of oxygen in compounds is always -2 EXCEPT for peroxides (-1), superoxides (-½) and oxygen fluorides (+2)
Step 2: Compare the oxidation state (oxidation number) of each element on the left hand side of the equation to the oxidation state of the same element on the right hand side of the equation:
Oxidation state (oxidation number) for hydrogen does NOT change.
Oxidation state (oxidation number) for oxygen DOES change:
One oxygen atom from the hydrogen peroxide molecule (H2O2) is oxidised to O2(g).
The oxidation number for this oxygen increases from -1 (in peroxide) to 0 (in oxygen gas)
The electrons released by the oxidation reaction above are consumed by another of the hydrogen peroxide's oxygen atoms which is then reduced when the water molecule is produced.
The oxidation number for this oxygen decreases from -1 (in peroxide) to -2 (in water)
- Decide if this is a redox reaction and why
This is a redox reaction because there is a transfer of electrons from one oxygen atom to another, one oxygen is oxidised and the other is reduced.
- Is your answer plausible?
Hydrogen peroxide, H2O2(aq) is commonly used as an oxidising agent, so we expect it to be reduced in the presence of a reductant (reducing agent).
The decomposition of hydrogen peroxide shows the product of this reduction reaction, water.
In the presence of a stronger oxidising agent such as permanganate ion, hydrogen peroxide undergoes oxidation and oxygen gas is formed.
The decomposition of hydrogen peroxide also shows the product of this oxidation reaction, O2(g).
Therefore, the decomposition of hydrogen peroxide is a redox reaction because it is undergoing both reduction and oxidation.
- State your solution to the problem "is this a redox reaction and justify your answer":
(i) Yes, the decomposition of hydrogen peroxide is a redox reaction
(ii) The decomposition of hydrogen peroxide is a redox reaction because:
(a) Some of the oxygen atoms undergo oxidation, producing oxygen molecules and electrons.
(b) These electrons are then consumed by another oxygen atom which undergoes reduction producing water molecules.
