Redox reactions of acids with metals

Chemical changes • Reactions of acids

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Describe the role of spectator ions in acid-metal reactions.

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Spectator ions (such as Cl-) remain unchanged in oxidation state and do not participate in electron transfer.

Key concepts

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Definition of redox

Redox describes reactions that involve transfer of electrons between species. One species loses electrons (oxidation) and another gains electrons (reduction). The total number of electrons lost equals the total number of electrons gained, conserving charge. Oxidation and reduction can be identified by change in oxidation number or by writing half-equations that show electrons explicitly. A reaction that shows a metal becoming a cation while hydrogen ions become hydrogen atoms (then H2) is a redox reaction.

General behaviour of metals with acids

Acids supply hydrogen ions (H+) in solution. A metal atom reacts by losing electrons and forming a positive ion (Mn+). Hydrogen ions gain those electrons and form hydrogen atoms that pair to make H2 gas. The general ionic equation is: metal -> metal^n+ + n e-; n H+ + n e- -> (n/2) H2. Metals higher than hydrogen in the reactivity series lose electrons readily and react with acids. Metals less reactive than hydrogen do not reduce H+ and therefore do not produce hydrogen gas with dilute acids.

Half-equations and electron bookkeeping

Half-equations separate the oxidation and reduction processes to show electron transfer clearly. Example: Mg -> Mg2+ + 2e- (oxidation); 2H+ + 2e- -> H2 (reduction). Combining the half-equations cancels electrons and yields the full ionic equation. Half-equations show which species loses electrons (oxidised) and which gains electrons (reduced). Accurate balancing of electrons ensures conservation of charge and mass.

Identification of oxidised and reduced species

Oxidised species show an increase in oxidation number or loss of electrons. Reduced species show a decrease in oxidation number or gain of electrons. In acid-metal reactions, the metal atom is oxidised and the hydrogen ion is reduced. Example: Zn + 2HCl -> ZnCl2 + H2. Zinc: Zn -> Zn2+ + 2e- (oxidised). Hydrogen: 2H+ + 2e- -> H2 (reduced). The chloride ions are spectator ions and do not change oxidation state.

Oxidation numbers as a diagnostic tool

Oxidation numbers track electron loss or gain. An element that increases its oxidation number during a reaction undergoes oxidation. An element that decreases its oxidation number undergoes reduction. In acid-metal reactions, the metal oxidation number increases from 0 (elemental metal) to a positive value in the salt. Hydrogen changes from +1 (in H+) to 0 (in H2), showing reduction.

Limitations and special cases

Some metals do not react with acids because they sit below hydrogen in the reactivity series or form protective oxide layers. Transition metals often form less predictable redox behaviour due to multiple oxidation states. Concentrated oxidising acids (for example, concentrated nitric acid) can oxidise metals without producing H2, because the acid itself acts as an oxidising agent producing other products; such reactions fall under broader redox chemistry and require separate analysis.

Key notes

Important points to keep in mind

Redox equals electron transfer: oxidation = loss, reduction = gain.

Metal atoms start with oxidation number 0 and increase when oxidised.

Hydrogen in acids starts at +1 and decreases to 0 when reduced to H2.

Write half-equations to show electrons explicitly and to identify oxidised/reduced species.

Balance electrons between half-equations before combining.

Spectator ions do not change oxidation state and can be omitted from ionic equations.

Metals below hydrogen in the reactivity series usually do not produce hydrogen gas with dilute acids.

Concentrated oxidising acids can undergo different redox pathways and may not yield H2.