What Do Rusting Iron, Burning Wood, and Photosynthesis Have in Common?
Oxidation and Reduction: More Than Just Electron Transfer
- Redox reactions involve two simultaneous processes: oxidation and reduction.
- These terms can be understood in different ways depending on the context.
Oxidation
- Electron Loss: When an atom or ion loses electrons, it undergoes oxidation.
- Oxygen Gain: Oxidation can also refer to the addition of oxygen.
- Hydrogen Loss: Oxidation may involve the loss of hydrogen.
- Sodium loses an electron and is oxidized: $$
\text{Na(s)} \rightarrow \text{Na⁺(aq)} + e⁻
$$ - Magnesium gains oxygen and is oxidized: $$
2\text{Mg(s)} + \text{O₂(g)} \rightarrow 2\text{MgO(s)}
$$ - Methane loses hydrogen during combustion: $$
\text{CH₄(g)} + 2\text{O₂(g)} \rightarrow \text{CO₂(g)} + 2\text{H₂O(g)}
$$
- While oxygen gain and hydrogen loss are traditional definitions of oxidation, they are less commonly used today compared to the electron transfer definition.
- Always consider the context of the reaction.
Reduction
- Electron Gain: Reduction occurs when an atom or ion gains electrons.
- Oxygen Loss: Reduction can also mean the removal of oxygen.
- Hydrogen Gain: Reduction may involve the addition of hydrogen.
- Chlorine gains electrons and is reduced: $$
\text{Cl₂(g)} + 2e⁻ \rightarrow 2\text{Cl⁻(aq)}
$$ - Copper(II) oxide loses oxygen and is reduced: $$
\text{CuO(s)} + \text{H₂(g)} \rightarrow \text{Cu(s)} + \text{H₂O(g)}
$$ - Ethene gains hydrogen in this reaction: $$
\text{C₂H₄(g)} + \text{H₂(g)} \rightarrow \text{C₂H₆(g)}
$$
Oxidation States: A Tool for Tracking Electron Transfers
- To systematically analyze redox reactions, chemists assign oxidation states to atoms.
- These numbers help track the movement of electrons in a reaction.
Rules for Assigning Oxidation States
- Elements in Their Standard State: The oxidation state of an atom in its elemental form is always 0.
- Ions: The oxidation state equals the ion’s charge.
- Fluorine: Fluorine always has an oxidation state of -1 in compounds.
- Oxygen: Oxygen typically has an oxidation state of -2, except in peroxides ($-1$) or when bonded to fluorine ($+2$).
- Hydrogen: Hydrogen usually has an oxidation state of +1, except in hydrides ($-1$).
- Neutral Compounds: The sum of oxidation states in a neutral compound must equal 0.
- Polyatomic Ions: The sum of oxidation states in a polyatomic ion must equal the ion’s charge.
Consider the compound $\text{H₂SO₄}$ (sulfuric acid):
- Hydrogen: Each H is +1 (total = +2).
- Oxygen: Each O is -2 (total = -8).
- Sulfur: To balance the total charge to 0, sulfur must be +6.
Assign oxidation states to all atoms in $\text{K₂Cr₂O₇}$.
Recognizing Oxidizing and Reducing Agents
In redox reactions, the species that is oxidized acts as the reducing agent, and the species that is reduced acts as the oxidizing agent.
Reaction Between Zinc and Copper(II) Sulfate
$$
\text{Zn(s)} + \text{CuSO₄(aq)} \rightarrow \text{ZnSO₄(aq)} + \text{Cu(s)}
$$
- Zinc is oxidized (reducing agent).
- Copper(II) ion is reduced (oxidizing agent).
- Do not confuse the oxidizing and reducing agents.
- Remember: the oxidizing agent is reduced, and the reducing agent is oxidized.
Transition Metals and Variable Oxidation States
Transition metals, such as iron and manganese, often exhibit multiple oxidation states, enabling them to participate in diverse redox reactions.
Manganese in Permanganate Ion
$$\text{MnO}_4^- (aq) + 8\text{H}^+ (aq) + 5e^- \rightarrow \text{Mn}^{2+} (aq) + 4\text{H}_2\text{O} (l)$$
Manganese is reduced from +7 in $MnO_4^-$ to +2 in $Mn^{2+}$.
Iron in Iron(III) Reduction
$$\text{Fe}^{3+} (aq) + e^- \rightarrow \text{Fe}^{2+} (aq)$$
Iron is reduced from +3 in $Fe^{3+}$ to +2 in $Fe^{2+}$.
Chromium in Dichromate Ion
$$\text{Cr}_2\text{O}_7^{2-} (aq) + 14\text{H}^+ (aq) + 6e^- \rightarrow 2\text{Cr}^{3+} (aq) + 7\text{H}_2\text{O} (l)$$
Chromium is reduced from +6 in $Cr_2O_7^{2-}$ to +3 in $Cr^{3+}$.
Copper in Copper(II) Reduction
$$\text{Cu}^{2+} (aq) + 2e^- \rightarrow \text{Cu} (s)$$
Copper is reduced from +2 in $Cu^{2+}$ to 0 in solid copper metal.
The variable oxidation states of transition metals are crucial in catalysis, batteries, and biological processes.
- In the reaction $\text{Fe(s)} + \text{HCl(aq)} \rightarrow \text{FeCl₂(aq)} + \text{H₂(g)}$, identify the oxidizing and reducing agents.
- Why are transition metals particularly versatile in redox reactions?
- How do oxidation states simplify the analysis of complex reactions?
