Rucete ✏ SAT Chemistry In a Nutshell
12. Oxidation–Reduction
This chapter introduces redox reactions—processes that involve the transfer of electrons. It explains oxidation states, how to identify redox reactions, and the roles of oxidizing and reducing agents. It also connects redox to combustion and gives you the tools to analyze reactions from an electron-transfer perspective.
What Are Redox Reactions?
- Redox = reduction + oxidation: both must happen together.
- Oxidation = loss of electrons
Reduction = gain of electrons
("LEO the lion says GER") - In redox, one substance gives up electrons (oxidized), the other gains (reduced).
- Example (net ionic):
2Ag⁺ + Cu → Cu²⁺ + 2Ag- Ag⁺ is reduced to Ag
- Cu is oxidized to Cu²⁺
Activity Series and Spontaneity
- Elements higher in the activity series are more likely to lose electrons (more easily oxidized).
- A single replacement reaction occurs only if the element doing the replacing is higher than the one being replaced.
- Example:
Au + CuCl₂ → no reaction (gold is less active)
Zn + FeCl₃ → ZnCl₂ + Fe (zinc is more active)
Oxidation States
- Help track electron transfer in both ionic and molecular compounds.
- Oxidation state rules (simplified):
- Elemental form = 0 (e.g., O₂, Na)
- Monatomic ion = its charge (e.g., Na⁺ = +1)
- Fluorine = −1
- Oxygen = −2 (except in peroxides: −1)
- Hydrogen = +1 (−1 in metal hydrides)
- Sum of oxidation states = 0 for neutral compounds; equals charge for ions
- Example 1: Na₂SO₄ → Na = +1, O = −2, S = +6
- Example 2: CO₂ → O = −2 × 2 = −4, C = +4
- Example 3: Cr₂O₇²⁻ → O = −2 × 7 = −14; Cr must be +6 each
Recognizing Redox Reactions
- A reaction is a redox reaction if there is a change in oxidation state for at least one element.
- If nothing is oxidized or reduced, the reaction is not redox.
- Example of redox:
Zn + Cu²⁺ → Zn²⁺ + Cu- Zn goes from 0 to +2 → oxidized
- Cu²⁺ goes from +2 to 0 → reduced
- Example of non-redox:
HCl + NaOH → NaCl + H₂O- No change in oxidation numbers → not a redox reaction
Oxidizing and Reducing Agents
- The substance that is reduced (gains electrons) = oxidizing agent
- The substance that is oxidized (loses electrons) = reducing agent
- Example:
In Zn + Cu²⁺ → Zn²⁺ + Cu- Zn = reducing agent
- Cu²⁺ = oxidizing agent
- Tip: Agents are opposite of what they do.
Combustion Reactions
- A special type of redox where a substance reacts with oxygen and releases energy (heat/light).
- Products are usually CO₂ and H₂O if the fuel contains C and H.
- Common fuels: methane (CH₄), propane (C₃H₈), gasoline, glucose
- Example:
CH₄ + 2O₂ → CO₂ + 2H₂O ΔH = −890.3 kJ/mol - Combustion always involves:
- Oxidation of the fuel (C → CO₂, H → H₂O)
- Oxygen is reduced (O₂ → O in H₂O or CO₂)
Heat of Combustion (ΔHc)
- The amount of energy released when 1 mole of a substance is burned completely in oxygen.
- Can be measured using a calorimeter.
- Fuels with higher ΔHc values release more energy per gram or per mole.
- Used in energy analysis of reactions, food calories, and fuel comparisons.
In a Nutshell
Redox reactions involve the transfer of electrons between species and can be tracked using oxidation states. Oxidation is loss of electrons, and reduction is gain. The oxidizing agent is reduced, and the reducing agent is oxidized. Combustion is a highly exothermic redox process where substances burn in oxygen. Understanding redox allows you to analyze chemical change at the electron level—a key part of SAT Chemistry.