Redox

Redox Review

• Redox Reactions

  • Definition: Electron transfer processes or reactions

  • Characterized by two half-reactions

  • Oxidation Reaction:

  • Half-reaction that involves the loss of 1 or more electrons

  • Reduction Reaction:

  • Half-reaction that involves the gain of 1 or more electrons

  • Oxidation Numbers:

  • Imaginary charges assigned to atoms in a molecule, indicating the distribution of electrons.

Applications of Redox Reactions

• Common Examples:

  • Batteries:

  • Used in devices like cars, flashlights, cell phones, and computers

  • Metabolism of Food:

  • Fundamental reactions in biological systems

  • Combustion:

  • Chemical reactions involving oxygen

  • Chlorine Bleach (NaOCl):

  • A dilute solution that cleans through redox reactions; acts as an oxidizing agent that destroys stains by oxidizing them.

Redox Reaction Examples

• Example: Fireworks Displays

  • Net Reaction:
    $2Mg + O_2 <br>ightarrow 2MgO$

  • Oxidation Process:

  • $Mg ightarrow Mg^{2+} + 2e^{-}$

    • Loses electrons = oxidized

    • Acts as a reducing agent

  • Reduction Process:

  • $O_2 + 4e^{-} ightarrow 2O^{2-}$

    • Gains electrons = reduced

    • Acts as an oxidizing agent

Fundamental Concepts of Oxidation-Reduction Reactions

• Oxidation-Reduction Reactions

  • Definition: Electron transfer reactions where electrons are transferred from one substance to another.

  • Characteristics: Originally identified in combustion of fuels or reactions of metals with oxygen

  • Emphasis: Redox reactions because reduction and oxidation must always occur together.

• Processes Involved:

  • Oxidation:

  • Loss of electrons; Example:

    • $Na <br>ightarrow Na^{+} + e^{-}$

  • Oxidation Half-Reaction

  • Reduction:

  • Gain of electrons; Example:

    • $Cl_2 + 2e^{-} <br>ightarrow 2Cl^{-}$

  • Reduction Half-Reaction

  • Net Reaction Example:

  • $2Na + Cl_2 ightarrow 2Na^{+} + 2Cl^{-}$

    • Oxidation and reduction happen simultaneously; cannot have one without the other.

Oxidizing and Reducing Agents

• Oxidizing Agent:

  • Definition: Substance that accepts electrons from another substance.

  • Example: $Cl_2 + 2e^{-} <br>ightarrow 2Cl^{-}$

  • Role: The substance that is reduced.

• Reducing Agent:

  • Definition: Substance that donates electrons to another substance.

  • Example: $Na <br>ightarrow Na^{+} + e^{-}$

  • Role: The substance that is oxidized.

Oxidation Numbers (O.N.) Guidelines

1. Pure Element:

   • O.N. is zero (e.g., Na)

2. Monatomic Ion:

   • O.N. equals the charge of the ion (e.g., Cl^{-}).

3. Neutral Compound:

   • Sum of O.N. equals zero (e.g., HCl)

4. Polyatomic Ion:

   • Sum of ON equals the total charge on the ion (e.g., $NO_3^{-}$)

   • Generally, the negative oxidation number is assigned to the more electronegative element.

5. Hydrogen:

   • Assigned +1, except in metal hydrides (e.g., $NaH, BH_3$) where it is -1.

6. Oxygen:

   • Usually assigned -2; in peroxides (e.g., $O_2^{2-}$) is -1; and +2 in compounds like OF2.

7. Halogens (F, Cl, Br, I):

   • Usually -1; fluorine is always -1. When bonding to more electronegative atoms, they adopt a positive charge.

Understanding Oxidation and Reduction

• Definitions:

  • Oxidation:

  • Loss of electrons

  • Oxidation number increases (becomes more positive)

  • Mnemonic: OIL (Oxidation Is Loss)

  • Reduction:

  • Gain of electrons

  • Oxidation number decreases (becomes more negative)

  • Mnemonic: RIG (Reduction Is Gain)

Redox Reaction Examples and Identification

• Example: Identify Oxidizing Agent in Reaction

  • Reaction: $Zn(s) + Pt^{2+}(aq) <br>ightarrow Pt(s) + Zn^{2+}(aq)$

  • Options:

  • A. Pt(s)

  • B. Zn^{2+}(aq)

  • C. Pt^{2+}(aq)

  • D. Zn(s)

  • E. None of these, as this is not a redox reaction.

• Example: Species getting oxidized

  • Reaction: $2Ag^{+}(aq) + Zn(s) <br>ightarrow Zn^{2+}(s) + 2Ag(s)$

  • Options:

  • A. Ag(s)

  • B. Ag^{+}(aq)

  • C. Zn^{2+}(aq)

  • D. Zn(s)

  • E. None, as this is not a redox.

• Example: Identify Oxidizing Agent and Reduced Species

  • Reaction: $2H^{+}(aq) + Mn(s) <br>ightarrow Mn^{2+}(aq) + H_2(g)$

  • Options:

  • A. H^{+} is oxidizing and Mn reduced

  • B. H^{+} is oxidizing and H^{+} reduced

  • C. Mn oxidizing and H^{+} reduced

  • D. Mn oxidizing and Mn reduced

  • E. Mn oxidizing and Mn^{2+} reduced

Guidelines for Redox Reactions

• Key Principles:

  • Oxidation and reduction always occur simultaneously.

  • Total number of electrons lost by one substance equals the total number of electrons gained by another substance.

  • For a redox reaction to occur, an electron acceptor must be present.

Rules for Assigning Oxidation Numbers

1. Oxidation numbers must add up to the charge on the molecule, formula unit, or ion

2. Atoms in free elements have oxidation numbers of zero

3. Metals in Group 1A, 2A, and Al have +1, +2, and +3 oxidation numbers, respectively

4. H and F in compounds have +1 and -1 oxidation numbers, respectively

5. Oxygen typically has -2 oxidation number

6. Group 7A elements usually have -1 oxidation number

7. Group 6A elements typically have -2 oxidation number

8. Group 5A elements are usually assigned -3 oxidation number

9. In cases of rule conflicts or ambiguity, apply the rule with the lower oxidation number and disregard the conflicting rule

Oxidation State Definitions

• Oxidation State: Used interchangeably with oxidation number; indicates charge on monatomic ions. For example, Iron(III) means +3 oxidation state of Fe or Fe^{3+}.

Examples of Assigning Oxidation Numbers

• Example 1: Assigning Oxidation Numbers in $ClO_4^{-}$

  • Oxygen (4 atoms) x (-2) = -8

  • Adding $Cl (1 atom) x (x) = x$ gives $-8 + x = -1$, leading to $x = +7$; thus Cl has an oxidation state of +7.

• Example 2: Assigning oxidation numbers in $Li_2O$

  • $Li (2 atoms) x (+1) = +2$ and

  • $O (1 atom) x (-2) = -2$, balancing to zero: $+2 - 2 = 0$

Balancing Redox Reactions

• Ion Electron Method:

  • Goal: Balance mass and charge in redox equations

  • Steps:

  1. Write skeleton equation

  2. Break into two half-reactions (oxidation and reduction)

  3. Balance each half-reaction separately

  4. Recombine to get a balanced net ionic equation

• Specific Approach in Acidic Solutions:

  1. Divide into half-reactions

  2. Balance atoms excluding H and O

  3. Balance O by adding H2O

  4. Balance H by adding H+

  5. Balance charge by adding electrons

  6. Equalize electron gain and loss, then combine half-reactions

  7. Cancel any common species

Example: Balancing Redox in Acidic Medium

• Example: Reactants: $Fe^{2+} + MnO_4^{-} <br>ightarrow Mn^{2+} + Fe^{3+}$

• Breakdown:

  • Oxidation Half-Reaction: $Fe^{2+} <br>ightarrow Fe^{3+} + e^{-}$

  • Reduction Half-Reaction: $MnO<em>4^{-} + 8H^{+} + 5e^{-} ightarrow Mn^{2+} + 4H</em>2O$

• Complete reaction:

  • Combine both half-reactions and balance the electrons.

  • Final balanced equation: $5Fe^{2+} + MnO<em>4^{-} + 8H^{+} ightarrow 5Fe^{3+} + Mn^{2+} + 4H</em>2O$

Galvanic Cell Overview

• Definition: A device converting chemical energy into electrical energy through spontaneous redox reactions generating a current for work.

• Cell Components:

  • Anode: Site of oxidation; reducing agent

  • Cathode: Site of reduction; oxidizing agent

  • Salt Bridge: Contains strong electrolyte that allows ion flow without extensive mixing

  • Porous Disk: Facilitates ion flow through tiny passages

Electrochemical and Cell Potential

• Cell Potential (E_cell):

  • The driving force on electrons, measured in volts (V); calculated as the difference in potential between oxidizing agent and reducing agent.

  • Formula:
    $ext{E_cell} = ext{E}^o ext{(cathode)} - ext{E}^o ext{(anode)}$

• Understanding Half-Reactions:

  • All half-reactions listed in standard tables are given as reduction processes.

  • When a half-reaction is reversed, the sign of $E^o$ is also reversed, while multiplying by an integer does not affect $E^o$.

Practical Example of Galvanic Cell Reaction

• Example:

  • Overall Reaction: $2Fe^{3+} + Cu(s) <br>ightarrow Cu^{2+} + 2Fe^{2+}$

  • Half-Reactions:

  • $Fe^{3+} + e^{-} <br>ightarrow Fe^{2+} ext{ with E}^o = 0.77 V$

  • $Cu^{2+} + 2e^{-} <br>ightarrow Cu ext{ with E}^o = 0.34 V$

  • To balance: Reverse the second reaction; thus:
    $Cu <br>ightarrow Cu^{2+} + 2e^{-}$

  • Overall Balanced Reaction:
    $Cu + 2Fe^{3+} <br>ightarrow Cu^{2+} + 2Fe^{2+}$

  • Cell Potential Calculation:
    $ext{E}^o_{ ext{cell}} = 0.77 V - 0.34 V = 0.43 V$

Line Notation for Electrochemical Cells

• Formatting Rules:

  • Anode components listed on the left, cathode components on the right, separated by double vertical lines (||).

  • Concentrations of aqueous solutions should be specified.

  • Example Notation:
    $Mg(s) | Mg^{2+}(aq) || Al^{3+}(aq) | Al(s)$

Cell Characteristics in Galvanic Cells

• The cell potential is typically positive for a spontaneous redox reaction, which drives the current flow and work output.

• Electrode details and the ions present in compartments are essential to identifying the nature of each half-reaction.

Work and Cell Potential

• Work vs. Maximum Work:

  • Work is never at maximum during current flow; also, energy is always lost in spontaneous processes, meaning actual work is less than maximum calculations.

• Maximum Cell Potential Relation:

  • Directly related to the free energy difference between reactants and products in the cell, calculated with:
    $ext{ΔG} = -nF ext{E}^o$

  • Where $F = 96,485 C/mol e^{-}$