17.1 Review of Redox Reaction
An oxidation-reduction (redox) reaction is a type of chemical reaction that involves a transfer of electrons between two species. An oxidation-reduction reaction is any chemical reaction in which the oxidation number of a molecule, atom, or ion changes by gaining or losing an electron. Redox reactions are common and vital to some of the basic functions of life, including photosynthesis, respiration, combustion, and corrosion or rusting.
Redox reactions are comprised of two parts, a reduced half and an oxidized half, that always occur together. The reduced half gains electrons and the oxidation number decreases, while the oxidized half loses electrons and the oxidation number increases. Simple ways to remember this include the mnemonic devices OIL RIG, meaning “oxidation is loss” and “reduction is gain,” and LEO says GER, meaning “loss of e– = oxidation” and “gain of e– = reduced.” There is no net change in the number of electrons in a redox reaction. Those given off in the oxidation half reaction are taken up by another species in the reduction half reaction.
The two species that exchange electrons in a redox reaction are given special names. The ion or molecule that accepts electrons is called the oxidizing agent; by accepting electrons it causes the oxidation of another species. Conversely, the species that donates electrons is called the reducing agent; when the reaction occurs, it reduces the other species. In other words, what is oxidized is the reducing agent and what is reduced is the oxidizing agent. (Note: the oxidizing and reducing agents can be the same element or compound, as in disproportionation reactions).
A good example of a redox reaction is the thermite reaction, in which iron atoms in ferric oxide lose (or give up) O atoms to Al atoms, producing Al2O3.
Fe2O3(s) + 2Al(s)→Al2O3(s) + 2Fe(l)
Oxidation Numbers
By definition, a redox reaction is one that entails changes in oxidation number (or oxidation state) for one or more of the elements involved. The oxidation number of an element in a compound is essentially an assessment of how the electronic environment of its atoms is different in comparison to atoms of the pure element.
The oxidation number is a number that indicates the number of electrons that an atom uses when it forms a compound.
The oxidation number is positive if the atom lends electrons, or shares with a more electronegative atom. And it will be negative when the atom gains electrons or shares them with an atom less electronegative.
The oxidation number is written in Roman numbers (remember it in the nomenclature of Stock): +I, +II, +III, +IV,-I,-II,-III,-IV, etc. But in this web also we use Arabic characters for them: +1, +2, +3, +4, -1, -2, -3, -4. etc., as if they were whole numbers.
In monoatomic ions the electrical charge coincides with the oxidation number. In the oxidation number we write the sign + or – it on the left of the number, like in the whole numbers. The charge of the ions, or charge number, must be written with the sign on the right of the digit: Ca2+ ion calcium(2+), CO32- ion carbonate(2-).
How can we know the oxidation number that corresponds to each atom? It is enough to know the oxidation number of the elements that have fixed oxidation number, which are a few, and it is very easy to deduce it from the electronic configurations. These oxidation numbers appear in the following table. The oxidation numbers of the other elements are deduced from the formulae or are indicated in the name of the compound.
Simple ionic compounds present the simplest examples to illustrate this formalism, since by definition the elements’ oxidation numbers are numerically equivalent to ionic charges. Sodium chloride, NaCl, is comprised of Na+ cations and Cl− anions, and so oxidation numbers for sodium and chlorine are, +1 and −1, respectively. Calcium fluoride, CaF2, is comprised of Ca2+ cations and F− anions, and so oxidation numbers for calcium and fluorine are, +2 and −1, respectively.
Covalent compounds require a more challenging use of the formalism. Water is a covalent compound whose molecules consist of two H atoms bonded separately to a central O atom via polar covalent O−H bonds. The shared electrons comprising an O−H bond are more strongly attracted to the more electronegative O atom, and so it acquires a partial negative charge in the water molecule (relative to an O atom in elemental oxygen). Consequently, H atoms in a water molecule exhibit partial positive charges compared to H atoms in elemental hydrogen. The sum of the partial negative and partial positive charges for each water molecule is zero, and the water molecule is neutral.
Imagine that the polarization of shared electrons within the O−H bonds of water were 100% complete—the result would be transfer of electrons from H to O, and water would be an ionic compound comprised of O2− anions and H+ cations. And so, the oxidations numbers for oxygen and hydrogen in water are −2 and +1, respectively. Applying this same logic to carbon tetrachloride, CCl4, yields oxidation numbers of +4 for carbon and −1 for chlorine. In the nitrate ion, NO−3NO3−, the oxidation number for nitrogen is +5 and that for oxygen is −2, summing to equal the 1− charge on the molecule: