Redox Reactions

1. Fundamentals of Redox Reactions

• Redox stands for reduction-oxidation.

• A redox reaction is a chemical reaction in which electrons are transferred from a reducing agent (reductant) to an oxidizing agent (oxidant).

• Oxidation and reduction always occur simultaneously; one process cannot take place without the other.

2. Key Definitions

• Oxidation:

  • Loss of electrons (mnemonic: OIL - Oxidation Is Loss).

  • Increase in oxidation state or oxidation number.

  • Addition of oxygen or removal of hydrogen.

• Reduction:

  • Gain of electrons (mnemonic: RIG - Reduction Is Gain).

  • Decrease in oxidation state or oxidation number.

  • Removal of oxygen or addition of hydrogen.

3. Oxidants and Reductants

• Oxidizing Agent (Oxidant):

  • The species that accepts electrons from another substance.

  • Undergoes reduction while causing another species to be oxidized.

• Reducing Agent (Reductant):

  • The species that donates electrons to another substance.

  • Undergoes oxidation while causing another species to be reduced.

4. Rules for Assigning Oxidation Numbers

Oxidation numbers (states) track electron transfer within reactions:

1. Uncombined or free elements always have an oxidation state of $0$ (e.g., $Mg(s)$, $O_2(g)$, $Cl_2(g)$).

2. Monatomic ions have an oxidation state equal to their charge (e.g., $Na^+$ is $+1$, $Fe^{3+}$ is $+3$, $Cl^-$ is $-1$).

3. Hydrogen in compounds is usually $+1$ (except in metal hydrides such as $NaH$, where it is $-1$).

4. Oxygen in compounds is usually $-2$ (except in peroxides such as $H_2O_2$, where it is $-1$, or when bonded to fluorine).

5. Neutral molecules have a sum of oxidation states equal to $0$.

6. Polyatomic ions have a sum of oxidation states equal to the charge of the ion (e.g., $SO_4^{2-}$ sums to $-2$, $MnO_4^-$ sums to $-1$).

5. Balancing Half-Equations in Acidic Conditions (KOHES Method)

To write balanced redox half-equations in acidic solutions, apply the KOHES sequence:

1. K (Key atoms): Balance all atoms except oxygen and hydrogen.

2. O (Oxygen): Balance oxygen atoms by adding water ($H_2O$) molecules.

3. H (Hydrogen): Balance hydrogen atoms by adding hydrogen ions ($H^+$).

4. E (Electrons): Balance the total charge by adding electrons ($e^-$) to the side with the higher positive charge.

5. S (State symbols): Include physical states ($(s)$, $(l)$, $(g)$, $(aq)$).

6. NCEA Level 2 Key Oxidants, Reductants, and Observations

6.1 Common Oxidants (Species Reduced)

• Acidified Permanganate Ion ($MnO_4^- / H^+$):

  • Half-equation: $MnO_4^-(aq) + 8H^+(aq) + 5e^- \rightarrow Mn^{2+}(aq) + 4H_2O(l)$

  • Oxidation state change: Mn changes from $+7$ to $+2$.

  • Observation: Purple solution turns colourless.

• Acidified Dichromate Ion ($Cr_2O_7^{2-} / H^+$):

  • Half-equation: $Cr_2O_7^{2-}(aq) + 14H^+(aq) + 6e^- \rightarrow 2Cr^{3+}(aq) + 7H_2O(l)$

  • Oxidation state change: Cr changes from $+6$ to $+3$.

  • Observation: Orange solution turns green.

• Halogens (e.g., Chlorine $Cl_2$, Iodine $I_2$):

  • Half-equation for $Cl_2$: $Cl_2(g) + 2e^- \rightarrow 2Cl^-(aq)$

  • Oxidation state change: Cl changes from $0$ to $-1$.

  • Observation: Pale green gas/solution turns colourless solution.

  • Half-equation for $I_2$: $I_2(aq) + 2e^- \rightarrow 2I^-(aq)$

  • Oxidation state change: I changes from $0$ to $-1$.

  • Observation: Brown solution turns colourless.

• Iron(III) Ions ($Fe^{3+}$):

  • Half-equation: $Fe^{3+}(aq) + e^- \rightarrow Fe^{2+}(aq)$

  • Oxidation state change: Fe changes from $+3$ to $+2$.

  • Observation: Orange-brown solution turns pale green.

• Hydrogen Peroxide ($H_2O_2$ as oxidant):

  • Half-equation: $H_2O_2(aq) + 2H^+(aq) + 2e^- \rightarrow 2H_2O(l)$

  • Oxidation state change: O changes from $-1$ to $-2$.

  • Observation: Colourless solution remains colourless.

6.2 Common Reductants (Species Oxidized)

• Reactive Metals (e.g., $Zn$, $Mg$, $Fe$):

  • Half-equation for $Zn$: $Zn(s) \rightarrow Zn^{2+}(aq) + 2e^-$

  • Oxidation state change: Zn changes from $0$ to $+2$.

  • Observation: Grey metal solid dissolves to form a colourless solution.

• Iodide Ions ($I^-$):

  • Half-equation: $2I^-(aq) \rightarrow I_2(aq) + 2e^-$

  • Oxidation state change: I changes from $-1$ to $0$.

  • Observation: Colourless solution turns brown.

• Iron(II) Ions ($Fe^{2+}$):

  • Half-equation: $Fe^{2+}(aq) \rightarrow Fe^{3+}(aq) + e^-$

  • Oxidation state change: Fe changes from $+2$ to $+3$.

  • Observation: Pale green solution turns orange-brown.

• Sulfur Dioxide ($SO_2$) / Sulfite Ions ($SO_3^{2-}$):

  • Half-equation for $SO_2$: $SO_2(g) + 2H_2O(l) \rightarrow SO_4^{2-}(aq) + 4H^+(aq) + 2e^-$

  • Oxidation state change: S changes from $+4$ to $+6$.

  • Observation: Colourless gas/solution forms a colourless solution.

• Hydrogen Peroxide ($H_2O_2$ as reductant):

  • Half-equation: $H_2O_2(aq) \rightarrow O_2(g) + 2H^+(aq) + 2e^-$

  • Oxidation state change: O changes from $-1$ to $0$.

  • Observation: Bubbles of colourless gas ($O_2$) form in a colourless solution.

7. Combining Half-Equations to Form Overall Balanced Equations

1. Multiply half-equations by integers so that the number of electrons lost equals the number of electrons gained.

2. Add the two equations together.

3. Cancel common species (electrons, $H^+$ ions, and $H_2O$ molecules) appearing on both sides.

4. Ensure total charge and numbers of atoms are balanced.

• Example: Combination of $MnO_4^-$ and $Fe^{2+}$ in acidic conditions:

  • Oxidation: $5 \times (Fe^{2+}(aq) \rightarrow Fe^{3+}(aq) + e^-)$

  • Reduction: $MnO_4^-(aq) + 8H^+(aq) + 5e^- \rightarrow Mn^{2+}(aq) + 4H_2O(l)$

  • Overall Equation: $MnO_4^-(aq) + 8H^+(aq) + 5Fe^{2+}(aq) \rightarrow Mn^{2+}(aq) + 4H_2O(l) + 5Fe^{3+}(aq)$

To determine whether a substance is oxidised or reduced in a reaction, you can use three key methods:

1. Track Oxidation Numbers (Most Reliable Method)

Assign oxidation numbers to elements in both reactants and products and observe the change:

• Oxidised: The oxidation number increases (becomes more positive) because the substance loses electrons. For example, in $Fe^{2+}(aq) \rightarrow Fe^{3+}(aq) + e^-$, the oxidation state of iron increases from $+2$ to $+3$.

• Reduced: The oxidation number decreases (becomes more negative) because the substance gains electrons. For example, in $Cl_2(g) + 2e^- \rightarrow 2Cl^-(aq)$, the oxidation state of chlorine decreases from $0$ to $-1$.

2. Look at Electron Transfer (OIL RIG)

Remember the mnemonic OIL RIG:

• OIL (Oxidation Is Loss): If electrons ($e^-$) are written on the product (right) side of a half-equation, electrons were lost, so the reactant is oxidised (e.g., $Zn(s) \rightarrow Zn^{2+}(aq) + 2e^-$).

• RIG (Reduction Is Gain): If electrons ($e^-$) are written on the reactant (left) side of a half-equation, electrons were gained, so the reactant is reduced (e.g., $MnO_4^-(aq) + 8H^+(aq) + 5e^- \rightarrow Mn^{2+}(aq) + 4H_2O(l)$).

3. Oxygen and Hydrogen Transfer

• Oxidised: The species gains oxygen atoms or loses hydrogen atoms.

• Reduced: The species loses oxygen atoms or gains hydrogen atoms.

Why This Happens

Redox reactions rely on the conservation of charge and mass through electron transfer. Electrons cannot exist freely in aqueous solution, so oxidation and reduction must always occur simultaneously:

• The substance that gets oxidised loses electrons and donates them to another species; therefore, it acts as the reducing agent (reductant).

• The substance that gets reduced gains electrons from another species; therefore, it acts as the **oxidising

1. Fundamentals of Redox Reactions

• Redox stands for reduction-oxidation.

• A redox reaction is a chemical reaction in which electrons are transferred from a reducing agent (reductant) to an oxidizing agent (oxidant).

• Oxidation and reduction always occur simultaneously; one process cannot take place without the other.

2. Key Definitions

• Oxidation:

  • Loss of electrons (mnemonic: OIL - Oxidation Is Loss).

  • Increase in oxidation state or oxidation number.

  • Addition of oxygen or removal of hydrogen.

• Reduction:

  • Gain of electrons (mnemonic: RIG - Reduction Is Gain).

  • Decrease in oxidation state or oxidation number.

  • Removal of oxygen or addition of hydrogen.

3. Oxidants and Reductants

• Oxidizing Agent (Oxidant):

  • The species that accepts electrons from another substance.

  • Undergoes reduction while causing another species to be oxidized.

• Reducing Agent (Reductant):

  • The species that donates electrons to another substance.

  • Undergoes oxidation while causing another species to be reduced.

4. Rules for Assigning Oxidation Numbers

Oxidation numbers (states) track electron transfer within reactions:

1. Uncombined or free elements always have an oxidation state of $0$ (e.g., $Mg(s)$, $O_2(g)$, $Cl_2(g)$).

2. Monatomic ions have an oxidation state equal to their charge (e.g., $Na^+$ is $+1$, $Fe^{3+}$ is $+3$, $Cl^-$ is $-1$).

3. Hydrogen in compounds is usually $+1$ (except in metal hydrides such as $NaH$, where it is $-1$).

4. Oxygen in compounds is usually $-2$ (except in peroxides such as $H_2O_2$, where it is $-1$, or when bonded to fluorine).

5. Neutral molecules have a sum of oxidation states equal to $0$.

6. Polyatomic ions have a sum of oxidation states equal to the charge of the ion (e.g., $SO_4^{2-}$ sums to $-2$, $MnO_4^-$ sums to $-1$).

5. Balancing Half-Equations in Acidic Conditions (KOHES Method)

To write balanced redox half-equations in acidic solutions, apply the KOHES sequence:

1. K (Key atoms): Balance all atoms except oxygen and hydrogen.

2. O (Oxygen): Balance oxygen atoms by adding water ($H_2O$) molecules.

3. H (Hydrogen): Balance hydrogen atoms by adding hydrogen ions ($H^+$).

4. E (Electrons): Balance the total charge by adding electrons ($e^-$) to the side with the higher positive charge.

5. S (State symbols): Include physical states ($(s)$, $(l)$, $(g)$, $(aq)$).