REDOX REACTION

Oxidation and reduction are two fundamental concepts in chemistry known collectively as redox reactions, which involve the transfer of electrons between chemical species.

Definition of Oxidation:

• Electronic Concept: Oxidation is defined as the loss of one or more electrons by an atom or ion.

• Charge Shift: During oxidation, the oxidation number of an element increases (moves in a positive direction on the number line).

• Electrochemical Site: Oxidation occurs at the Anode, which is the negative ($(-)$) electrode in certain electrochemical contexts.

Definition of Reduction:

• Electronic Concept: Reduction is defined as the gain of one or more electrons by an atom or ion.

• Charge Shift: During reduction, the oxidation number of an element decreases (moves in a negative direction on the number line; "charge goes DOWN").

• Electrochemical Site: Reduction occurs at the Cathode, which is the positive ($+$) electrode.

Summary Mnemonic:

• Oxidation = Loss of Electrons (Charge increases).

• Reduction = Gain of Electrons (Charge decreases).

Prototypical Redox Reaction: Zinc and Copper Sulfate

• Chemical Equation: $Zn + CuSO<em>4 \rightarrow ZnSO</em>4 + Cu$

• Mechanism of Oxidation:

  • The zinc ($Zn$) atom begins with an oxidation state of $0$.

  • It loses two electrons to become $Zn^{2+}$ in the compound zinc sulfate ($ZnSO_4$).

  • Change: $0 \rightarrow +2$.

• Mechanism of Reduction:

  • The copper ($Cu$) ion in copper sulfate ($CuSO_4$) begins with an oxidation state of $+2$.

  • It gains two electrons to become a neutral copper ($Cu$) atom.

  • Change: $+2 \rightarrow 0$.

Applied Drills in Oxidation State Determination

• Drill 1: Tracking the Charge

  • Reaction 1: $Al^0 \rightarrow Al^{+3}$. Charge changes from $0$ to $+3$. Classification: Oxidation.

  • Reaction 2: $O^0 \rightarrow O^{-2}$. Charge changes from $0$ to $-2$. Classification: Reduction.

  • Reaction 3: $Fe^{+2} \rightarrow Fe^{+3}$. Charge changes from $+2$ to $+3$. Classification: Oxidation.

  • Reaction 4: $Ag^{+1} \rightarrow Ag^0$. Charge changes from $+1$ to $0$. Classification: Reduction.

• Drill 2: Magnesium in Hydrochloric Acid

  • Full Reaction: A magnesium strip is dipped into hydrochloric acid ($Mg + 2HCl \rightarrow MgCl<em>2 + H</em>2$).

  • Reactant Charges:

  • Magnesium ($Mg$): $0$ (as it is a pure element alone).

  • Hydrogen ($H$) in $HCl$: $+1$.

  • Chlorine ($Cl$) in $HCl$: $-1$.

  • Product Charges:

  • Magnesium ($Mg$) in $MgCl_2$: $+2$.

  • Hydrogen ($H$) in $H_2$: $0$.

  • Electron Exchange Analysis:

  • Element Oxidized: Magnesium ($Mg$) lost electrons.

  • Element Reduced: Hydrogen ($H$) gained electrons (specifically, the proton stole electrons from the metal).

  • Half-Reactions:

  • Oxidation Half-Reaction: $Mg^0 \rightarrow Mg^{2+} + 2e^{-}$ (Loses $2$ electrons).

  • Reduction Half-Reaction: $2H^{+} + 2e^{-} \rightarrow H_2$ (Gains $2$ electrons total).