Questions: Assigning Oxidation Numbers and Identifying Redox

Question 1

In H₂O₂ (hydrogen peroxide), what is the oxidation state of oxygen, and why does the standard rule not apply?

• A. −2, because oxygen is always −2 in any compound
• B. −1, because each oxygen is bonded to another oxygen of equal electronegativity, so neither atom wins the shared electrons outright
• C. 0, because the molecule is electrically neutral
• D. +1, because hydrogen contributes +1 to balance each oxygen

Answer: (B)

The standard rule (O = −2) assumes oxygen bonds to a less electronegative element and claims all shared electrons. In peroxides, each oxygen is bonded to another oxygen — identical electronegativity means the electrons are split equally, giving each oxygen −1 instead of −2. This is the classic peroxide exception. Students who memorize 'oxygen is always −2' without understanding the electronegativity basis will assign the wrong value here.

Question 2

In the reaction Zn + H₂SO₄ → ZnSO₄ + H₂, which species is reduced, and how do you know?

• A. Zinc, because it forms a compound and gains a positive charge
• B. Sulfur, because it is present on both sides as sulfate
• C. Hydrogen, because it goes from +1 in H₂SO₄ to 0 in H₂ — a decrease in oxidation number
• D. Oxygen, because it ends up in the sulfate product on both sides

Answer: (C)

Reduction means a decrease in oxidation number (gain of electron density — the 'R' in 'OIL RIG'). Hydrogen starts at +1 in H₂SO₄ and becomes 0 in elemental H₂ — a decrease, so hydrogen is reduced. Zinc goes from 0 to +2 (oxidized). Sulfur stays at +6 throughout; oxygen stays at −2 throughout. Neither undergoes a change, so neither participates in the electron transfer.

Question 3

Oxidation states represent the actual electrical charges that atoms carry within a compound.

• T. True
• F. False

Answer: False

Oxidation states are a bookkeeping fiction, not a physical measurement. They are assigned by pretending all bonds are fully ionic — giving all shared electrons to the more electronegative atom — even in covalent molecules where electrons are only partially shifted. For example, carbon in CO₂ is assigned +4, but it does not carry a genuine +4 charge. The real charge distribution is determined by partial electronegativities, bond polarity, and electron delocalization. Oxidation states are useful for tracking electron transfers in redox reactions, not for describing actual charge.

Question 4

If no atom's oxidation number changes on either side of a chemical equation, the reaction is not a redox reaction.

• T. True
• F. False

Answer: True

Redox reactions are defined by electron transfer, and oxidation numbers track electron ownership. If every atom has the same oxidation number on both sides of the equation, no electrons have changed hands — there has been no oxidation or reduction. The reaction might be a precipitation, an acid-base neutralization, or a complex-forming reaction, but it is not a redox reaction. Changing oxidation numbers are both necessary and sufficient to classify a reaction as redox.

Question 5

Explain why hydrogen has an oxidation state of +1 in water (H₂O) but −1 in sodium hydride (NaH), and what principle underlies this difference.

Model answer: In H₂O, oxygen is more electronegative than hydrogen, so oxygen 'wins' the shared electrons in the O–H bond. Hydrogen is assigned +1 because it is the electron-donor in the bond. In NaH, sodium is more electropositive (less electronegative) than hydrogen, so hydrogen 'wins' the shared electrons in the Na–H bond. Hydrogen is now the electron-acceptor, giving it an oxidation state of −1. The underlying principle is electronegativity: in the oxidation number system, the more electronegative atom in a bond is assigned all the shared electrons. When hydrogen bonds to an element more electropositive than itself (metals), hydrogen becomes the winner and takes on a negative oxidation state.

The peroxide and metal hydride exceptions both stem from the same principle: the standard rules (O = −2, H = +1) hold only when the usual electronegativity ordering applies. When oxygen bonds to oxygen, or hydrogen bonds to a metal more electropositive than itself, the usual hierarchy breaks down and the exceptions apply. Understanding the electronegativity basis means you can always derive the correct assignment rather than memorizing a list of exceptions.