Oxidation Numbers and Oxidation States

This is part of Year 11 HSC Chemistry course under the topic of Predicting Reactions of Metals

HSC Chemistry Syllabus

Apply the definitions of oxidation and reduction in terms of electron transfer and oxidation numbers to a range of reduction and oxidation (redox) reactions

Oxidation Numbers and States Explained

This video explains what oxidation numbers are states are. How can you calculate oxidation numbers and states? How do oxidation numbers and states explain to us whether oxidation or reduction has occurred?

What Are Oxidation States?

An oxidation state, also known as an oxidation number, is a theoretical charge that an atom would have if all its bonds were completely ionic. Understanding oxidation states is crucial in redox chemistry as it aids in identifying the substances that are oxidised or reduced. This is especially useful when it's unclear whether a chemical species is gaining or losing electrons (OIL-RIG mnemonic).

Oxidation State Changes:

An increase in the oxidation state of a substance signifies that it has been oxidised

A decrease in the oxidation state indicates the substance has been reduced

The magnitude of change in oxidation state infers the number of electron(s) gained or lost. For example, an increase in oxidation state of 2 means two electrons are lost. Vice versa, a decrease in oxidation state of 3 means 3 electrons are gained (reduction).

Rules for Determining Oxidation Numbers

Memorising a set of rules can greatly simplify the process of calculating oxidation numbers. Here is a quick overview:

Elemental Form Rule: Atoms in their elemental form have an oxidation number of zero. Examples include Na (sodium), Mg (magnesium), and Cu (copper) in their neutral states, as well as diatomic gases like `O_2` (oxygen), `N_2` (nitrogen), and `Cl_2` (chlorine).

Monoatomic Ions: The oxidation state of a monoatomic ion is equal to its charge. Sodium ions (`Na^+`) and magnesium ions (`Mg^{2+}`), for example, have oxidation numbers of +1 and +2 respectively.

Neutral Compounds: The sum of oxidation numbers/states for the atoms in a neutral molecule or a compound equals zero.

Polyatomic ions: Similar to monoatomic ions, polyatomic ions such as hydrogen (`OH^–`), carbonate (`CO_3^{2–}`), and phosphate (`PO_4^{3–}`) have oxidation numbers equal to their charges of -1, -2, and -3 respectively.

Electronegativity rule: When analysing the oxidation state of two covalently bonded atoms, we consider the bond as being purely ionic, with the more electronegative atom taken to have all the electrons in the bond(s). For example, in carbon dioxide `O=C=O`, the four electrons in each of the `C=O` double bonds are assigned to the oxygen atom, which is more electronegative than carbon. So, carbon would be considered to have 'lost' the four valence electrons it contributed, giving it an oxidation state of +4. When the covalently bonded atoms are identical and thus have the same electronegativity, the electrons are considered to be evenly split amongst these atoms. This is consistent with the elemental form rule for diatomic gases.

Rules For Specific Atoms and Groups of Elements

For Group 1 elements (alkaline metals): oxidation number is `+1` in all ionic compounds.

For Group 2 elements (alkaline earth metals): oxidation number is `+2` in all ionic compounds.

For hydrogen: oxidation number is `+1` in covalent compounds (with non-metals) and `-1` in compounds with metals or boron.

For oxygen: oxidation number is `-2` in most compounds except in peroxides and compounds containing fluorine.

For group 17 elements (halogens): oxidation number is `-1` in all compounds except when combined with more electronegative atoms e.g. oxygen, fluorine and a more electronegative halogen. Fluorine always has an oxidation number of `-1` because it is the most electronegative element.

Examples

Carbonic Acid (`H_2CO_3`)

The sum of oxidation number in this compound is zero as the compound is neutral. Each oxygen atom has an oxidation number of –2, and each hydrogen atom has an oxidation number of +1. With three oxygen atoms, we have a total of –6, and two hydrogen atoms contribute a total oxidation number of +2. This means the carbon atom must have an oxidation number of +4 to achieve the total 0 state for the compound.

Sulfate Ion (`SO_4^{2-}`)

Following the rules, the sulfate ion as a whole has an oxidation state of –2. Oxygen typically has an oxidation state of –2, except in peroxides. With four oxygen atoms, we have a total of –8. Consequently, sulfur must have an oxidation state of +6 to achieve the –2 state for the entire ion.

Dichromate Ion (`Cr_2O_7^{2–}`)

The dichromate ion, with a total charge of –2 and seven oxygens, leads to oxygen having a total of –14. To balance this, each chromium must have an oxidation state of +6.

Nitrate Ion (`NO_3^–`)

For the nitrate ion, with a charge of –1 and three oxygens contributing –6, nitrogen has an oxidation state of +5.

Permanganate Ion (`MnO_4^–`)

The permanganate ion has a charge of –1. With four oxygens totalling –8, manganese's oxidation state is +7.

Hypochlorous Acid (`HClO`)

The total oxidation state is zero as the compound is neutral. The oxygen atom has an oxidation state of –2, and the hydrogen atom +1. In this case, the chlorine atom has an oxidation state of +1 to achieve a net state of zero. Note that the chlorine does not have a typical oxidation state of –1 in this compound because it is combined with oxygen which is more electronegative.

Applying Oxidation Numbers to Reactions

In compounds, oxidation numbers must sum up to the compound's overall charge. For example, in calcium carbonate (`CaCO_3`), calcium has an oxidation number of +2, while the carbonate ion has –2, balancing the compound to a neutral charge.

Magnesium and Copper Sulfate Reaction

Consider the reaction between magnesium and copper sulfate forming magnesium sulfate and elemental copper:

$$Mg(s) + CuSO_4(aq) \rightarrow MgSO_4(aq) + Cu(s)$$

Here's how we apply oxidation states:

Magnesium (Mg) starts with an oxidation state of 0.
Copper (Cu) in copper sulfate has an oxidation state of +2.
Sulfate (`SO_4^{2–}`) maintains an oxidation state of –2 throughout the reaction.

After the reaction:

Magnesium appears in magnesium sulfate (`MgSO_4`) with an oxidation state of +2
Copper is reduced to its elemental form with an oxidation state of 0.

The increase from 0 to +2 for the oxidation state of magnesium means it has been oxidised, while the decrease from +2 to 0 for copper means it has been reduced. However, in this example the application of oxidation state is not particularly useful as it is rather easy to identify which metal or metal ion is losing or gaining electrons.

Magnesium and Water Reaction

Consider the reaction between magnesium and water:

$$Mg(s) + 2H_2O(l) \rightarrow Mg(OH)_2(aq) + H_2(g)$$

While it is easy to identify that magnesium undergoes oxidation to become magnesium ion, the species undergoes reduction is more unclear.

Let's assign oxidation states:

Magnesium (Mg) starts with an oxidation state of 0.
Hydrogen (H) starts with an oxidation state of +1. Two hydrogen atoms contribute combined oxidation state of +2.
Oxygen (O) starts with an oxidation state of –2.