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Oxidation-Reduction Reactions

Chapter 11: Oxidation-Reduction Reactions

5 min read Updated Mar 26, 2026
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🎯 Diagnostic: Test Your Starting Level 24 questions (2 per section). No prior reading required, see what you already know.

Aim to answer every question before checking. Missed questions point you to the sections you need most.

1. (11.1) The oxidation state of a pure element in its natural form is:
A. O₂, N₂, Na(s), and Fe(s) all have oxidation state 0. This is the starting reference for tracking electron transfer.
2. (11.1) The oxidation state of oxygen in most compounds (excluding peroxides) is:
C. -2 in most oxides. Exceptions: peroxides (-1), superoxides (-12\frac{1}{2}), and OF₂ (+2 because F is more electronegative).
3. (11.2) Oxidation is:
D. "OIL RIG": Oxidation Is Loss, Reduction Is Gain (of electrons). "LEO the lion says GER" works too.
4. (11.2) Reduction is:
B. Gaining negatively charged electrons makes the species more negative, i.e., "reduced" in oxidation number.
5. (11.3) An oxidizing agent:
A. The oxidizer oxidizes something else and gets reduced in the process. KMnO₄ is a classic oxidizing agent.
6. (11.3) A reducing agent:
C. Electron donors reduce something else. NADH and FADH₂ are common biological reducing agents.
7. (11.4) In balancing redox reactions in acidic solution:
D. Then balance charges with electrons, multiply to match electron counts, and add the half-reactions.
8. (11.4) The final step in balancing a redox equation is:
B. Electrons must cancel across the combined equation. After addition, simplify any H₂O or H⁺ that appears on both sides.
9. (11.5) To convert an acid-balanced redox equation to basic solution:
A. Each H⁺ + OH⁻ combines to H₂O, which cancels with existing water molecules. The net result has no free H⁺.
10. (11.5) In a redox reaction balanced for basic solution, typical species include:
C. Basic solutions are defined by excess OH⁻; H⁺ concentration is very low, so it should not appear in the net equation.
11. (11.6) A half-reaction represents:
D. Splitting the full reaction makes electron bookkeeping straightforward. Standard reduction potentials are tabulated on this basis.
12. (11.6) Electrons appear on which side of an oxidation half-reaction?
B. Oxidation produces electrons. In a reduction half-reaction they appear as reactants (left side).
13. (11.7) The activity series of metals ranks metals by:
A. A more-active metal will displace a less-active metal's cation from solution (e.g., Zn(s) + Cu²⁺ → Zn²⁺ + Cu(s)).
14. (11.7) An alkali metal like sodium near the top of the activity series:
C. Active metals like Na, K reduce water on contact, producing H₂ gas.
15. (11.8) A disproportionation reaction is one in which:
D. Disproportionation requires an element that can exist in three oxidation states relatively easily.
16. (11.8) Which is an example of disproportionation?
B. Oxygen starts at -1 in H₂O₂, goes to 0 in O₂ (oxidation) and -2 in H₂O (reduction).
17. (11.9) A common strong oxidizing agent in the lab is:
A. Mn(VII) is reduced during oxidation of organic or inorganic substrates. K₂Cr₂O₇ and H₂O₂ are also common strong oxidants.
18. (11.9) H₂O₂ can serve as:
C. Oxygen in H₂O₂ sits between 0 and -2, so it can move either direction.
19. (11.10) A common reducing agent is:
D. Hydride sources reduce aldehydes to alcohols and carboxylic acids to alcohols (LiAlH₄) or aldehydes/ketones to alcohols (NaBH₄).
20. (11.10) Alkali metals react with water by:
B. 2 Na + 2 H₂O → 2 NaOH + H₂. The reaction is vigorous; heavier alkali metals (K, Rb, Cs) can ignite the hydrogen.
21. (11.11) In biological redox, NAD⁺:
A. NAD⁺ picks up 2 electrons and 1 H⁺, becoming NADH. It ferries reducing equivalents to the ETC.
22. (11.11) FADH₂ is:
C. FAD + 2 H⁺ + 2 e⁻ → FADH₂. Complex II (succinate dehydrogenase) uses the FAD/FADH₂ couple.
23. (11.12) Complete combustion of a hydrocarbon in excess O₂ produces:
D. Incomplete combustion (limited O₂) produces CO or soot instead.
24. (11.12) In CH₄ + 2 O₂ → CO₂ + 2 H₂O:
B. Total electrons transferred: 8 (one C losing 8; four O each gaining 2).

Every time you take a breath, you are running a redox reaction. Oxygen enters your lungs, travels to your mitochondria, and accepts electrons at the end of the electron transport chain. Those electrons were stripped from the food you ate - glucose was oxidized, carbon dioxide was exhaled, and the energy released is what keeps you alive right now. Rusting iron, charging your phone, bleaching a stain, and burning gasoline are all the same fundamental process: one substance loses electrons while another gains them.

Oxidation-reduction (redox) reactions are among the most frequently tested topics on the MCAT because they bridge general chemistry, biochemistry, and biology. The Chem/Phys section will ask you to assign oxidation states, balance half-reactions, and predict spontaneous displacement reactions. The Bio/Biochem section expects you to recognize NAD+/NADH and FAD/FADH2 as electron carriers and to understand how cellular respiration is a series of coupled redox events. If you can track where electrons go, you can answer all of these questions.

This chapter builds from the ground up. We start with the rules for assigning oxidation states, move through balancing redox equations in both acidic and basic solution, and then connect everything to the activity series and standard reduction potentials. By the end, you will see how the same electron-transfer logic that explains a zinc strip dissolving in copper sulfate also explains why your cells need oxygen to make ATP.

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