· 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.