The oxidation state of an atom, also called the oxidation number, is a numerical value that represents the charge an ion has or an atom appears to have when its electrons are counted based on a set of commonly accepted rules. The value (+ or -) of the oxidation state symbolizes the apparent charge for each atom in a chemically bonded molecule based on the distribution of the electrons in the molecule. For ionically bonded compounds, the oxidation state of an atom is equal to the ionic charge of the atom. The charge of an atom in an ionic compound represents the electron distribution of the compound. The oxidation state of an atom bound in a covalent compound is the average charge for the atom based on the electronegativity values for each atom involved in the molecule. In this case, the oxidation state is equal to the charge the atom in a covalent compound would have if the electrons in the molecule were distributed around the atoms with the greatest electronegativities.
The oxidation state of an atom is not necessarily equal to the charge of the atom. This difference is expressed in the notation of oxidation state versus ionic charge. The oxidation state is written as a superscript with the charge followed by the number, for example, aluminum (Al+3) or fluorine (F-1). Ionic charges, on the other hand, are written as a superscript with the number followed by the charge. Examples of ionic charge notation include the hydrogen ion (H1+) and the oxygen ion (O2-). Oxidation states do not represent an actual physical property of an atom, as ionic charges do. An ionic charge of 1- means the atom has gained an electron, whereas an oxidation state of -1 means that atom has a greater attraction to a bonding electron. Oxidation states are simply a way of keeping track of the electrons involved in a molecule during a chemical reaction.
Oxidation and reduction are two processes that are quite common in chemical reactions. Oxidation occurs when an atom loses one or more electrons, becoming more positive. Reduction occurs when an atom gains one or more electrons, becoming more negative. Oxidation always takes place at the same time as reduction. When an atom loses an electron during oxidation, that electron needs to go somewhere. It is picked up by another atom during the process of reduction. Reactions which involve oxidation and reduction are called oxidation-reduction reactions, or simply redox reactions.
Redox reactions involve the transfer of electrons between atoms. The atom that undergoes reduction, i.e., the atom that receives the electron in a reaction, is called the oxidizing agent. The oxidizing agent causes the oxidation of an atom by receiving an electron from that atom and is itself reduced. The atom that undergoes oxidation, i.e., the atom that loses the electron in a reaction, is called the reducing agent. The reducing agent causes the reduction of an atom by donating an electron and is itself oxidized.
When an atom is oxidized, it loses an electron and becomes more positive. When this occurs, its oxidation state increases. Likewise, when an atom is reduced, it gains an electron and becomes more negative. When this occurs, its oxidation state decreases. In every redox reaction, an atom must increase its oxidation state and an atom must decrease its oxidation state. A change in oxidation state of an atom indicates that a redox reaction has taken place. Assigning oxidation states to the chemical compounds involved in a redox reaction is important because it allows every electron to be accounted for, which in turn allows for the formation of balanced equations for the reaction.
As mentioned earlier, the oxidation state of an atom is not necessarily the same as the ionic charge of an atom. The oxidation state of an atom can be determined by studying the bonding characteristics of that atom, but this requires time and facilities that many people do not have. As a general rule, shared electrons in a compound are assumed to belong to the atom that has a higher electronegativity, but this does not help much in assigning oxidation states, either. To make the process simpler, a set of rules has been developed to determine the oxidation state of an atom involved in a reaction. First, the oxidation state of any atom in an uncombined element is equal to zero. For example, the oxidation state of an elemental carbon (C) atom is zero, as is the oxidation state of an elemental sulfur (S) atom. The second rule specifies that the oxidation state of any monoatomic ion is equal to its ionic charge. A chlorine ion (Cl1-) has an oxidation state of -1, for example, and a magnesium ion (Mg2+) has an oxidation state of +2. The third rule gives the oxidation states of many elements in a compound according to the element's position in the periodic table. Elements in group 1A on the periodic table always have an oxidation state of +1. Elements in group 2A on the periodic table always have an oxidation state of +2. Aluminum (Al) is always +3 and fluorine is always -1. The oxidation state of hydrogen is always +1 if it is combined with a nonmetal, and the oxidation state of oxygen is usually -2 in most ions and compounds. The fourth rule says that the oxidation states of all the atoms in a compound must total zero. The last rule specifies that the oxidation states of all the atoms in a polyatomic ion must total the charge of the ion.
Not every atom follows the rules for assigning oxidation states. Many nonmetals, for example, can have more than one oxidation number. For example, sulfur can have an oxidation state of +4 or +6. The Stock system of nomenclature eliminates some of the resulting confusion by placing the oxidation state of a substance in Roman numerals within parentheses as part of the name of a compound. For example, sulfur with an oxidation state of +4 forms the compound SO2, named sulfur (IV) oxide. The form of sulfur with an oxidation state of +6 forms the compound SO3, named sulfur (VI) oxide. In general, however, the rules for assigning oxidation states can be followed and used to name compounds, write formulas, or balance chemical equations.
With the assistance of the above rules, oxidation states can be determined for the atoms in almost any compound. The oxidation state of an atom can give valuable information about that atom. For example, the oxidation number describes how well an atom is going to bond with other atoms. A positive oxidation state indicates that the atom can undergo oxidation, i.e., lose electrons. A negative oxidation state indicates that an atom can undergo reduction, i.e., gain electrons. A negative oxidation state indicates a stronger attraction for electrons than does a positive oxidation state. The larger the negative value, the larger the attraction for the electrons. For example, an atom with an oxidation state of -3 will have a stronger attraction for electrons than one with an oxidation state of -1. Similarly, the larger the positive value of an atom's oxidation state, the smaller the attraction for the electrons. An atom with an oxidation state of +3 will have a weaker attraction for electrons than one with an oxidation state of +1. The oxidation number of an atom can also be used to predict how it will combine with other atoms as well as what the resulting chemical formula and balanced chemical equation will be.
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