By employing oxidation numbers, we may broaden our notion of redox to include oxidation and reduction in covalent chemical reactions (oxidation numbers are also called oxidation states). Each atom or ion in a chemical is assigned an oxidation number, which indicates the degree of oxidation. The number of oxidations might be positive, negative, or zero. The plus or minus sign must always be present. Higher positive oxidation numbers indicate that an atom or ion has been oxidized more. Higher negative oxidation values indicate a more reduced atom or ion.
What is an Oxidation Number?
The oxidation number (O.N.) plays a crucial role in predicting the course of reactions and naming molecules.
The term “oxidation number” refers to a number assigned to a molecule or element to indicate how many electrons were acquired or lost.
A negative value indicates that electrons were obtained, while a positive number indicates that electrons were lost. Since electrons have a negative charge and more electrons will make the compound more negative, a negative number is utilized to obtain electrons. The overall outcome would be more favorable if electrons were lost, hence the oxidation number is positive.
The O.N. of an atom is sometimes defined as the hypothetical charge assumed on that atom if the compound were made up of ions.
Oxidation Numbers Rule
The following rules are used to assign these oxidation numbers. Using oxidation number criteria, we may calculate the oxidation number of any atom or ion. It’s critical to understand that an oxidation number refers to a single atom in a molecule.
- The oxidation number of any uncombined element is zero i.e. O.N. of free element is always zero. For example, the oxidation number of each atom in S8, Cl2, and Zn is zero.
- The oxidation number of an element in a monatomic ion is always the same as the charge. For example, O.N. of Cl– is –1, Al3+ is +3.
- The sum of the oxidation numbers of all atoms in a neutral compound is zero.
- The sum of the oxidation numbers in an ion is equal to the charge on the ion.
- In either a compound or an ion, the more electronegative element is given the negative oxidation number.
- The charge of a polyatomic ion is equal to the sum of its oxidation numbers. The sum of the oxidation numbers for SO42-, for example, is -2.
- In compounds many atoms or ions have fixed oxidation numbers
- Group 1 elements have O.N. always +1
- Group 2 elements are always with O.N. +2
- Fluorine have O.N. always –1
- Hydrogen is +1 (except in metal hydrides such as NaH, where it is –1)
- Oxygen is –2 (except in peroxides e.g. BaO2, where it is –1 due to the structure [O-O] 2-, and in F2O, where it is +2 because F is more electronegative than O).
- Except when paired with a higher electronegativity element, the oxidation number of a Group VIIA element in a compound is -1. In HCl, the oxidation number of Cl is -1, but in HOCl, the oxidation number of Cl is +1.
Calculating Oxidation Number
The oxidation number or state of an atom/ion is the number of electrons obtained or lost by the molecule in comparison to the neutral atom. Electropositive metal atoms in groups I, II, and III lose a particular number of electrons while maintaining constant positive oxidation values.
More electronegative atoms in molecules receive electrons from less electronegative atoms and exhibit negative oxidation states. The oxidation state’s numerical value is equal to the number of electrons lost or gained.
An atom’s or ion’s oxidation number or oxidation state in a molecule/ion is assigned by:
i) Compiling the continuous oxidation states of other atoms/molecules/ions linked to it; and
ii) Equivalently relating molecules or ions’ total oxidation state to their total charge.
Compounds containing a metal and a nonmetal
The metal always has a positive O.N. , while the nonmetal always has a negative O.N.
In the case of sodium oxide, Na2O: Na = +1 and O = -2.
- If we don’t know the O.N. of one of the atoms, we can usually figure it out using the rule of an invariable O.N.
- In sodium sulfide, for example, O.N. of each Na atom = +1 for two sodium atoms = +2 Na2S has no overall charge, hence the total O.N. is zero.
- O.N. of S = -2
Non-metal compounds with non-metals
The sign of the O.N. in compounds comprising two separate nonmetals is determined by the electronegativity of each atom . The negative symbol is assigned to the most electronegative element.
- SO2 [Sulfur dioxide]
Because SO2 has no charge, the total Oxidation number is zero.
O.N. of each O atom = -2
Hence O.N. for two oxygen atoms = 2 (-2) = -4
O.N. of S = +4
- ICl3 [Iodine trichloride]
Chlorine is more electronegative than iodine, so chlorine is – and iodine is +
O.N. of each Cl atom = –1
Hence O.N. for three chlorine atoms = 3 × (–1) = –3
ICl3 has no charge, so the total O.N. is zero
O.N. of I = +3
- N2H4 [Hydrazine]
Nitrogen is more electronegative than hydrogen, so nitrogen is – and hydrogen is +
O.N. of each H atom = +1 (rule 2)
Hence, O.N. for four hydrogen atoms = 4 × (+1) = +4
N2H4 has no charge, so the total O.N. is zero.
O.N. of two N atoms = –4
O.N. of each N atom = –2
Compound ions are ions that contain two or more distinct atoms.
NO3– [Nitrate ion]
O.N. of each O atom = –2
Hence. O.N. for three oxygen atoms = 3 × (–2) = –6
NO3– has a charge of 1–, so the total O.N. of N and O atoms is –1
O.N. of the nitrogen atom + O.N. of the three oxygen atoms (–6) = –1
O.N. of N = +5
Oxidation Number in Monatomic Compounds
Monatomic elements are easy to determine the O.N. because it is equal to the charge seen on the element.
- LI+: O.N. = +1
- Ba2: O.N. = +2
- Fe: O.N. = 0
- N2: O.N. = 0 (even though the nitrogen is bound to another nitrogen, this is the free species of the element nitrogen)
- Cl–: O.N. = -1
Oxidation Number of Atoms in a Diatomic Molecule
A diatomic molecule may be homonuclear or heteronuclear.
Homonuclear diatomic molecule
The concept of oxidation number applies exclusively to heteroatoms that compose a molecule. As a result, the oxidation number of the atoms in a homonuclear diatomic molecule is zero. The oxidation number of hydrogen, oxygen, nitrogen, and chlorine molecules in their respective molecules is zero.
Heteronuclear diatomic molecule
All bonds generated between atoms in hetero diatomic compounds are termed ionic.
More electronegative atoms are thought to steal bonding electrons from less electronegative ones. As a result, the electronegative atom will have a negative oxidation state with a magnitude equal to the amount of electrons it has taken.
The electron is thought to have been transferred from the less electronegative atom to the more electronegative atom. As a result, the atom with the lowest electronegative charge will have a positive oxidation state equal to the amount of electrons lost.
|HCl as an example||H2O as an example|
|Chlorine has a higher electronegative potential than hydrogen. As a result, chlorine is supposed to remove one electron from hydrogen. Chlorine has an O.N. of -1 when it receives one electron, whereas hydrogen has an O.N.of +1 when it loses one electron.||Oxygen has a higher electronegative potential than hydrogen. As a result, the oxygen atom obtains one electron from each of the two hydrogen atoms and has an O.N. of -2. Both hydrogens that lose one electron have an O.N. of +1.|
Oxidation Number of Atoms in a Diatomic Molecule
Oxidation: An increase in the O.N. of an element in a compound. (Loss of electron(s) by species)
Reduction: A reduction in the O.N. of an element in a specific compound. (Gain of electron(s))
Oxidizing Agent: It is a substance that increases the O.N. of an element in a compound. (Acceptor of electron(s))
Reducing Agent: A reagent that reduces an element’s O.N. in a compound. (Donor of electron(s))
Redox Reactions: In redox reactions, the O.N. of the interacting species changes.