MCAT General Chemistry Review - Alexander Stone Macnow, MD 2019-2020

Oxidation–Reduction Reactions
Concept Summary

Oxidation—Reduction Reactions

· Oxidation is a loss of electrons, and reduction is a gain of electrons; the two are paired together in what is known as an oxidation—reduction (redox) reaction.

· An oxidizing agent facilitates the oxidation of another compound and is reduced itself in the process; a reducing agent facilitates the reduction of another compound and is itself oxidized in the process.

o Common oxidizing agents almost all contain oxygen or a similarly electronegative element.

o Common reducing agents often contain metal ions or hydrides (H).

· To assign oxidation numbers, one must know the common oxidation states of the representative elements.

o Any free element or diatomic species has an oxidation number of zero.

o The oxidation number of a monatomic ion is equal to the charge of the ion.

o When in compounds, Group IA metals have an oxidation number of +1; Group IIA metals have an oxidation number of +2.

o When in compounds, Group VIIA elements have an oxidation number of —1 (unless combined with an element with higher electronegativity).

o The oxidation state of hydrogen is +1 unless it is paired with a less electronegative element, in which case it is —1.

o The oxidation state of oxygen is usually —2, except in peroxides (when its charge is —1) or in compounds with more electronegative elements.

o The sum of the oxidation numbers of all the atoms present in a compound is equal to the overall charge of that compound.

· When balancing redox reactions, the half-reaction method, also called the ion—electron method, is the most common.

o Separate the two half-reactions.

o Balance the atoms of each half-reaction. Start with all the elements besides H and O. In acidic solution, balance H and O using water and H+. In basic solution, balance H and O using water and OH.

o Balance the charges of each half-reaction by adding electrons as necessary to one side of the reaction.

o Multiply the half-reactions as necessary to obtain the same number of electrons in both half-reactions.

o Add the half-reactions, canceling out terms on both sides of the reaction arrow.

o Confirm that the mass and charge are balanced.

Net Ionic Equations

· A complete ionic equation accounts for all of the ions present in a reaction. To write a complete ionic reaction, split all aqueous compounds into their relevant ions. Keep solid salts intact.

· Net ionic equations ignore spectator ions to focus only on the species that actually participate in the reaction. To obtain a net ionic reaction, subtract the ions appearing on both sides of the reaction, which are called spectator ions.

o For reactions that contain no aqueous salts, the net ionic equation is generally the same as the overall balanced reaction.

o For double displacement (metathesis) reactions that do not form a solid salt, there is no net ionic reaction because all ions remain in solution and do not change oxidation number.

· Disproportionation (dismutation) reactions are a type of redox reaction in which one element is both oxidized and reduced, forming at least two molecules containing the element with different oxidation states.

· Oxidation—reduction titrations are similar in methodology to acid—base titrations. These titrations follow transfer of charge.

o Indicators used in such titrations change color when certain voltages of solutions are achieved.

o Potentiometric titration is a form of redox titration in which a voltmeter or external cell measures the electromotive force (emf) of a solution. No indicator is used, and the equivalence point is determined by a sharp change in voltage.