Unit 13: Reduction-Oxidation (Redox)

Determining Oxidation Numbers

• Definition of Oxidation State: The oxidation number represents the charge an atom would have if all bonds to atoms of different elements were completely ionic.

• Uncombined Elements:     * The oxidation number for all non-bonded (uncombined) atoms is always zero ($0$).     * Examples provided:         * $\text{Na}^0$         * $\text{Al}^0$         * $\text{N}_2^0$         * $\text{O}_2^0$

• Bonded Elements (Compounds):     * Oxidation numbers for elements within compounds are found on the Periodic Table.     * Rule of Zero: The total charge of all atoms in a compound must equal zero.     * Example: Sodium Chloride ($\text{NaCl}$):         * Oxidation numbers: $\text{Na} = +1$, $\text{Cl} = -1$.         * Total charge: $(+1) + (-1) = 0$.     * Example: Sodium Perchlorate ($\text{NaClO}_4$):         * Oxidation values: $\text{Na} = +1$, $\text{Cl} = +7$, $\text{O} = -2$.         * Calculation: $(+1) + (+7) + 4(-2) = 0$.

• Polyatomic Ions:     * The sum of all oxidation numbers in a polyatomic ion must equal the overall charge of the ion, not zero.     * Example: Sulfate Ion ($\text{SO}_4^{2-}$):         * The total charge of the sulfate ion is $-2$.         * Oxidation numbers: $\text{S} = +6$, $\text{O} = -2$.         * Calculation check: $(+6) + 4(-2) = -2$.

• Rules for Hydrogen:     * Hydrogen is generally assigned an oxidation number of $+1$ in compounds.     * Exception: Hydrogen is $+1$ unless it is bonded with a metal (forming a hydride), in which case it is $-1$.     * Example Reaction: $\text{Zn}^0 + 2\text{H}^{+1}\text{Cl}^{-1} \rightarrow \text{H}_2^0 + \text{Zn}^{+2}\text{Cl}_2^{1-}$.

Reduction-Oxidation (Redox) Reactions

• Definition: Redox reactions are chemical processes characterized by the transfer of electrons between species.

• Oxidation:     * Defined as the loss of electrons ($e^-$).     * The oxidation number (charge) of the species goes up ($\uparrow$).     * Oxidized species in example: $\text{Fe}^0$.

• Reduction:     * Defined as the gain of electrons ($e^-$).     * The oxidation number (charge) of the species goes down ($\downarrow$).     * Reduced species in example: $\text{Cu}^{2+}$.

• Electron Transfer Pathways:     * Reaction 1: $\text{Fe} + \text{CuSO}_4 \rightarrow \text{Cu} + \text{FeSO}_4$.         * Electrons are transferred from $\text{Fe}^0$ to $\text{Cu}^{2+}$.     * Reaction 2: $\text{Fe} + \text{O}_2 \rightarrow \text{Fe}_2\text{O}_3^{3-2}$.         * Electrons are transferred from $\text{Fe}^0$ to $\text{O}^0$.

• Mnemonic Devices:     * LEO says GER:         * LEO: Loss of Electrons is Oxidation.         * GER: Gain of Electrons is Reduction.

Half-Reactions

• Half-reactions are used to show the separate reduction and oxidation components of a redox reaction.

• Conservation of Charge: Electrons ($e^-$) must be placed on the side of the equation that balances the total charge.

• Oxidation Half-Reaction Rules:     * Electrons ($e^-$) are always written on the Right (product) side.     * Example: $\text{Fe}^0 \rightarrow \text{Fe}^{2+} + 2e^-$.

• Reduction Half-Reaction Rules:     * Electrons ($e^-$) are always written on the Left (reactant) side.     * Example: $\text{Cu}^{2+} + 2e^- \rightarrow \text{Cu}^0$.

Balancing Redox Reactions

• Law of Conservation: In any redox reaction, Mass, Energy, and Charge must be conserved (equal on both sides).

• Step-by-Step Balancing Procedure:     1. First, balance the total charge by using coefficients in front of the ions.     2. Second, balance the total mass by using coefficients in front of the atoms.

• Example: Zinc and Silver:     * Unbalanced: $\text{Zn}^0 + \text{Ag}^+ \rightarrow \text{Zn}^{2+} + \text{Ag}^0$     * Balanced: $1\text{Zn}^0 + 2\text{Ag}^+ \rightarrow 1\text{Zn}^{2+} + 2\text{Ag}^0$

• Example: Magnesium and Aluminum:     * Unbalanced: $\text{Mg}^0 + \text{Al}^{3+} \rightarrow \text{Mg}^{2+} + \text{Al}^0$     * Balanced: $3\text{Mg}^0 + 2\text{Al}^{3+} \rightarrow 3\text{Mg}^{2+} + 2\text{Al}^0$

Electrochemical Cells: Voltaic Cell

• Nature of Reaction: Spontaneous.

• Energy Conversion: Chemical energy is converted into electrical energy.

• Materials: Constructed using two different metals.

• Electron Flow: Electrons flow from the metal higher on Table-J to the metal lower on current Table-J.

• Operational Principles:     * Anode (-):         * Site of Oxidation ("An OX").         * Reaction example: $\text{Zn} \rightarrow \text{Zn}^{2+} + 2e^-$.         * Physical Change: Loss of mass as atoms become ions.     * Cathode (+):         * Site of Reduction ("Red Cat").         * Reaction example: $\text{Cu}^{2+} + 2e^- \rightarrow \text{Cu}^0$.         * Physical Change: Gain of mass as ions become atoms.     * Salt Bridge: This component allows for the passage of ions and maintains electrical neutrality within the cell.

Electrochemical Cells: Electrolytic Cell

• Nature of Reaction: Non-spontaneous.

• Energy Conversion: Electrical energy is converted into chemical energy.

• Energy Requirement: Requires a power source (e.g., a battery) to force a non-spontaneous reaction.

• Electron Flow in Electrolytic Systems:     * Electrons flow into the positive ($+$) battery terminal.     * Electrons flow out of the negative ($-$) battery terminal.

• Electrolysis (Example: Molten NaCl):     * Reaction: $\text{NaCl} \rightarrow \text{Na}^+ + \text{Cl}^- \rightarrow \text{Na} + \text{Cl}$.     * Cathode (-): Terminal where positive ions (cations) are reduced. $\text{Na}^+ + e^- \rightarrow \text{Na}$.     * Anode (+): Terminal where negative ions (anions) are oxidized. $\text{Cl}^- \rightarrow \text{Cl} + e^-$.

• Electroplating:     * A process used to coat a surface with a thin layer of metal.     * The object to be plated is always the Cathode (-).     * Metal ions in solution are reduced onto the object: $\text{Ag}^+ + e^- \rightarrow \text{Ag}^0$.