oxidation

Oxidation−Reduction Reactions

Learning Goal

• Define the terms oxidation and reduction; identify the reactants oxidized and reduced.

• Assign oxidation states to elements within compounds and polyatomic ions.

• Example: Rust forms when oxygen reacts with iron through electron transfer.

Definition of Oxidation−Reduction Reactions

• Oxidation–Reduction Reaction (Redox Reaction):
    - Electrons are transferred from one substance to another.
  

Mnemonic for Redox Reactions: OIL RIG

• Oxidation Is Loss of electrons

• Reduction Is Gain of electrons

Relationship Between Oxidation and Reduction

• In every oxidation reaction, there must be a corresponding reduction reaction.

• Reducing Agent: The substance that is oxidized (loses electrons), which causes the reduction of the other substance.

• Oxidizing Agent: The substance that is reduced (gains electrons), which causes the oxidation of the other substance.

Examples of Oxidation and Reduction

• Oxidation of Copper (Patina on the Statue of Liberty):
    - Reaction:
      - $2Cu(s) + O_2(g) \rightarrow 2CuO(s)$
    - Oxidation Half-Reaction:
      - $2Cu(s) \rightarrow 2Cu^{2+}(s) + 4e^-$ (Copper loses electrons)
    - Reduction Half-Reaction:
      - $O_2(g) + 4e^- \rightarrow 4O^{2-}(s)$ (Oxygen gains electrons)

Learning Check Example

• Reaction in UV Light:
    - Reaction: $2Ag^+ + 2Cl^- \rightarrow 2Ag + Cl_2$
    - Question A: Which reactant is oxidized?
    - Answer: $2Cl^- \rightarrow Cl_2 + 2e^-$ (Chloride ions are oxidized)
    - Question B: Which reactant is reduced?
    - Answer: $2Ag^+ + 2e^- \rightarrow 2Ag$ (Silver ions are reduced)

Oxidation States

Definition of Oxidation States

• Oxidation States (also known as oxidation numbers):
    - Numbers assigned to atoms in compounds to help keep track of the electron transfer in redox reactions.
    - Important Note: Do not confuse oxidation state with ionic charge. Oxidation states can be assigned to compounds that are not ionic.

Rules for Assigning Oxidation States

1. A neutral element that is not part of a compound has an oxidation state of zero.

2. Monoatomic ions have oxidation states equal to their ionic charges.

3. The sum of the oxidation states in any formula equals the overall charge of the formula.

4. Oxygen tends to have an oxidation state of −2 in compounds.

5. Hydrogen tends to have an oxidation state of +1 in compounds.

Learning Check - Determine Oxidation States

• Determine the oxidation state of:
    - a. Mg: 0
    - b. P^{3-}: -3
    - c. Fe in FeCl3: +3

Learning Check - Oxidation State of Nitrogen in Nitrite Ion

• To find the oxidation state of nitrogen in the nitrite ion ($NO_2^-$), consider:
    - Oxygen typically has an oxidation state of -2 (Rule 4). Thus:

• The sum of oxidation states must equal the charge of the ion (-1).

• Let the oxidation state of nitrogen be $x$:
    - $x + 2(-2) = -1$
    - Solve for $x$:
      - $x - 4 = -1$
      - $x = +3$

• Therefore, the oxidation state of nitrogen is +3.

Learning Check - Oxidation State of Phosphorus in K3PO4

• Determine the oxidation state of phosphorus in potassium phosphate ($K_3PO_4$):
    - The oxidation state of phosphorus is +5.

Identifying Redox Reactions Using Oxidation States

• Oxidation: Occurs when an oxidation state of an element increases.

• Reduction: Occurs when an oxidation state of an element decreases.

Synthesis of Aluminum Chloride Example

• Aluminum: Oxidation state increases from 0 to +3; Aluminum is oxidized and acts as the reducing agent.

• Chlorine: Oxidation state decreases from 0 to −1; Chlorine is reduced and acts as the oxidizing agent.

Learning Check - Comparing Reaction Elements

• Determine oxidation and reduction in a given reaction:
    - Calcium: Oxidation state increases from 0 to +2; Calcium is oxidized and is the reducing agent.
    - Nitrogen: Oxidation state decreases from 0 to −3; Nitrogen is reduced and is the oxidizing agent.

Redox Reaction Examples

Identifying Redox Reactions from Given Examples

1. Reaction: $2 H_2(g) + O_2(g) \rightarrow 2 H_2O(g)$
      - Identify oxidation states:
        - H: 0 (in H2) to +1 (in H2O)
        - O: 0 (in O2) to -2 (in H2O)
      - Conclusion: Reduction and oxidation occur - Redox Reaction.

2. Reaction: $MgI_2(aq) + 2 AgNO_3(aq) \rightarrow 2 AgI(s) + Mg(NO_3)_2(aq)$
      - Analyze oxidation states: No change for Mg and Ag; therefore, not a redox reaction.

3. Reaction: $CaCO_3(s) \rightarrow CaO(s) + CO_2(g)$
      - Analyze oxidation states: No change for Ca, therefore, not a redox reaction.

4. Reaction: $2 C_8H_{18}(g) + 25 O_2(g) \rightarrow 16 CO_2(g) + 18 H_2O(l)$
      - Identify oxidation states and conclude: Redox Reaction.

5. Reaction: $Zn(s) + 2 HCl(aq) \rightarrow ZnCl_2(aq) + H_2(g)$
      - Identify oxidation: zinc oxidized, H+ ions reduced - Redox Reaction.

Single-Replacement Reactions

Characteristics of Single-Replacement Reactions

• Single-replacement reactions are categorized as redox reactions.

• Active metals can oxidize by transferring electrons to ions of less active metals.
   - Example: Zinc (more active) transfers electrons to copper ions forming zinc ions and copper atoms.

• An activity series lists metals in order of their relative reactivity.

Redox Reaction Example Between Zn Metal and Cu2+

• Zinc reacts with copper ions as it is more active, leading to zinc being oxidized.

Learning Check - Predicting Reactions

• Analyze pairs of reactants:
    1. $Cu(s) + Li^+(aq) \rightarrow$ No Reaction
    2. $Li(s) + Cu^{2+}(aq) \rightarrow 2 Li^+(aq) + Cu(s)$
    3. $Zn^{2+}(aq) + Mg^{2+}(aq) \rightarrow$ No Reaction
    4. $2 Al(s) + 3 SnCl_2(aq) \rightarrow 2 AlCl_3(aq) + 3 Sn(s)$
    5. $Li(s) + Cr(s) \rightarrow$ No Reaction

Balanced Chemical Equations for Single-Replacement Reactions

• Li(s) + CuCl_2(aq):
    - Balanced reaction: $2 Li(s) + CuCl_2(aq) \rightarrow Cu(s) + 2 LiCl(aq)$

• Al(s) + Sn(NO3)2(aq):
    - Balanced reaction: $2 Al(s) + 3 Sn(NO_3)_2(aq) \rightarrow 3 Sn(s) + 2 Al(NO_3)_3(aq)$