Oxidation Number – Rules, Examples, and How to Find It

An oxidation number or oxidation state is a number assigned to an atom that describes its degree of oxidation and helps track electrons in compounds and chemical reactions. Oxidation numbers may be positive, negative, or zero. They are especially important for understanding oxidation-reduction reactions, naming compounds, predicting chemical behavior, and balancing redox equations.

An oxidation number is a bookkeeping value rather than necessarily the actual electrical charge on an atom. Chemists assign oxidation numbers by imagining that bonding electrons belong entirely to the more electronegative atom. Simple rules make it possible to determine oxidation numbers from most chemical formulas without drawing detailed structures.

Key Takeaways: Oxidation Number

• An oxidation number describes the formal degree of oxidation of an atom.
• Oxidation number and oxidation state are largely synonymous terms.
• An element in its uncombined elemental form has an oxidation number of 0.
• The oxidation number of a monatomic ion equals its ionic charge.
• The sum of oxidation numbers equals the overall charge of a molecule or ion.
• Oxygen usually has an oxidation number of −2, while hydrogen usually has an oxidation number of +1.
• Oxidation increases oxidation number, while reduction decreases oxidation number.

Why Oxidation Numbers Are Important

Oxidation numbers provide a simple way of keeping track of electrons in chemical compounds and reactions. They are useful even when a substance does not contain discrete ions.

Chemists use oxidation numbers to:

• Identify which atoms undergo oxidation and reduction
• Determine oxidizing and reducing agents
• Balance oxidation-reduction reactions
• Name compounds containing metals with multiple oxidation states
• Describe the chemistry of transition metals
• Compare related compounds of an element
• Predict likely formulas and reactions
• Follow electron transfer in electrochemistry
• Organize trends in inorganic chemistry

For example, iron occurs commonly as Fe²⁺ and Fe³⁺. Oxidation numbers distinguish iron(II) compounds from iron(III) compounds and help predict how one form converts into the other.

History of the Oxidation Number Concept

The idea behind oxidation numbers developed long before chemists understood electrons.

During the late 18th century, Antoine Lavoisier associated oxidation with reaction with oxygen. Chemists soon recognized that elements could combine with different amounts of oxygen, producing different degrees or stages of oxidation.

In 1835, German chemist Friedrich Wöhler used the term Oxydationsstufe, or oxidation grade, to describe these different degrees of oxidation.

As atomic theory and electrochemistry developed during the 19th and early 20th centuries, oxidation became connected with changes in electrical character rather than simply the addition of oxygen.

German chemist Alfred Stock introduced his influential system of inorganic nomenclature in 1919. In Stock nomenclature, Roman numerals distinguish different oxidation levels, such as iron(II) chloride and iron(III) chloride.

American chemist Wendell M. Latimer helped establish the modern use of the terms oxidation state and oxidation number, particularly through his 1938 book The Oxidation States of the Elements and Their Potentials in Aqueous Solutions. Later treatments connected oxidation numbers explicitly with the assignment of bonding electrons according to electronegativity.

Modern IUPAC recommendations define oxidation state using an ionic approximation of chemical bonds, providing a rigorous basis for the familiar rules taught in introductory chemistry.

Oxidation Number Definition

The oxidation number of an atom is the hypothetical charge assigned to that atom after bonding electrons are allocated according to agreed rules.

For a simple introductory treatment, imagine that every bond becomes completely ionic. Assign both electrons in a bond to the more electronegative atom. If the two atoms are identical, divide the bonding electrons equally between them.

The resulting hypothetical charge is the oxidation number.

For example, consider hydrogen chloride:

HCl

Chlorine is more electronegative than hydrogen, so both bonding electrons are assigned to chlorine.

• Hydrogen has oxidation number +1.
• Chlorine has oxidation number −1.

The molecule itself is neutral because:

(+1) + (−1) = 0

This does not mean HCl consists of isolated H⁺ and Cl⁻ ions when it exists as a molecular gas. The oxidation numbers are formal bookkeeping values.

Oxidation Number vs Oxidation State

The terms oxidation number and oxidation state are usually used interchangeably.

Current IUPAC terminology describes oxidation number as largely synonymous with oxidation state. A subtle distinction is sometimes useful:

• Oxidation state emphasizes the state or degree of oxidation of an atom.
• Oxidation number emphasizes the numerical value assigned to that state.

So, saying that iron has an oxidation state of +3 and saying that its oxidation number is +3 generally mean the same thing.

Oxidation Number vs Valence Electrons

An oxidation number is not the same thing as the number of valence electrons.

Valence electrons are electrons in the outer portion of an atom that participate in bonding and chemical reactions.

An oxidation number is a formal number assigned to an atom after considering the atoms to which it is bonded.

For example, a neutral sulfur atom has six valence electrons. But sulfur takes several oxidation numbers in compounds:

• H₂S: −2
• elemental S: 0
• SO₂: +4
• SO₃: +6

Sulfur does not suddenly have −2, +4, or +6 valence electrons. Its valence-electron configuration helps explain its chemistry, while its oxidation number describes its formal oxidation level in a particular substance.

Oxidation Number vs Formal Charge

Oxidation number and formal charge are different bookkeeping concepts.

Formal charge assumes electrons in a covalent bond are divided equally between the bonded atoms.

Oxidation number assigns bonding electrons to the more electronegative atom.

Consider the ammonium ion, NH₄⁺.

Nitrogen has five valence electrons. In NH₄⁺, nitrogen forms four N−H bonds and has no lone pairs. For formal charge, nitrogen receives one electron from each bond:

Formal charge on N = 5 − 4 = +1

Each hydrogen has a formal charge of 0, so the formal charges add to the overall +1 charge of NH₄⁺.

For oxidation number, nitrogen is more electronegative than hydrogen, so the bonding electrons are assigned to nitrogen. Hydrogen has an oxidation number of +1:

N + 4(+1) = +1

N = −3

Thus, nitrogen in NH₄⁺ has:

• Formal charge = +1
• Oxidation number = −3

The values differ because formal charge divides bonding electrons equally, while oxidation number assigns them to the more electronegative atom.

Rules for Assigning Oxidation Numbers

The following rules solve most oxidation-number problems in general chemistry.

1. A Free Element Has an Oxidation Number of 0

Any element in its uncombined form has oxidation number 0.

Examples:

• Na = 0
• Fe = 0
• H₂ = 0
• O₂ = 0
• Cl₂ = 0
• S₈ = 0

It does not matter whether the element is monatomic or molecular.

2. A Monatomic Ion Has an Oxidation Number Equal to Its Charge

Examples:

• Na⁺ = +1
• Mg²⁺ = +2
• Fe³⁺ = +3
• Cl⁻ = −1
• S²⁻ = −2

3. The Sum of Oxidation Numbers in a Neutral Compound Is 0

For H₂O:

2(+1) + (−2) = 0

For CO₂:

(+4) + 2(−2) = 0

4. The Sum of Oxidation Numbers in a Polyatomic Ion Equals the Ion Charge

For sulfate, SO₄²⁻:

S + 4(−2) = −2

S = +6

5. Group 1 Metals Usually Have an Oxidation Number of +1

Examples include lithium, sodium, potassium, rubidium, and cesium.

NaCl: Na = +1

K₂SO₄: K = +1

Rare exotic compounds exist that do not follow this simple introductory rule.

6. Group 2 Metals Usually Have an Oxidation Number of +2

Examples include magnesium, calcium, strontium, and barium.

MgO: Mg = +2

CaCl₂: Ca = +2

7. Fluorine Has an Oxidation Number of −1 in Compounds

Fluorine is the most electronegative element, so it receives the bonding electrons.

Examples:

HF: F = −1

CF₄: F = −1

OF₂: F = −1

The last example produces an important exception to the usual rule for oxygen.

8. Oxygen Usually Has an Oxidation Number of −2

Examples:

H₂O: O = −2

CO₂: O = −2

SO₃: O = −2

There are important exceptions.

Peroxides: Oxygen is −1.

H₂O₂: O = −1

Na₂O₂: O = −1

Superoxides: The average oxidation number of oxygen is −1/2.

KO₂: O = −1/2

Oxygen bonded to fluorine: Oxygen has a positive oxidation number.

OF₂: O = +2

O₂F₂: O = +1

9. Hydrogen Usually Has an Oxidation Number of +1

Examples:

H₂O: H = +1

HCl: H = +1

NH₃: H = +1

When hydrogen bonds to an element less electronegative than hydrogen, its oxidation number is usually −1.

Examples include many metal hydrides:

NaH: H = −1

CaH₂: H = −1

10. Chlorine, Bromine, and Iodine Are Usually −1 When Bonded to More Electropositive Elements

For example:

NaCl: Cl = −1

HBr: Br = −1

KI: I = −1

However, these halogens can have positive oxidation numbers when bonded to oxygen, fluorine, or another more electronegative element.

For example, chlorine is +7 in ClO₄⁻.

Quick Oxidation Number Rules

Always remember that the total oxidation numbers must equal the overall electrical charge.

How to Find an Oxidation Number

A reliable method is:

1. Write the chemical formula and overall charge.
2. Assign oxidation numbers that are fixed or strongly predictable.
3. Multiply each oxidation number by the number of atoms of that element.
4. Set their sum equal to the overall charge.
5. Solve for the unknown oxidation number.
6. Check the result.

General Equation

For a neutral compound:

sum of oxidation numbers = 0

For an ion:

sum of oxidation numbers = ion charge

Worked Example 1: Oxidation Number of Sulfur in H₂SO₄

Find the oxidation number of sulfur in sulfuric acid.

Hydrogen is +1.

Oxygen is −2.

Let sulfur equal x.

2(+1) + x + 4(−2) = 0

2 + x − 8 = 0

x = +6

Answer: sulfur has oxidation number +6.

Worked Example 2: Oxidation Number of Manganese in KMnO₄

Potassium is +1.

Oxygen is −2.

Let manganese equal x.

(+1) + x + 4(−2) = 0

1 + x − 8 = 0

x = +7

Answer: manganese has oxidation number +7.

Worked Example 3: Oxidation Number of Nitrogen in NH₄⁺

Hydrogen is +1.

The total charge is +1.

Let nitrogen equal x.

x + 4(+1) = +1

x + 4 = 1

x = −3

Answer: nitrogen has oxidation number −3.

Worked Example 4: Oxidation Number of Chromium in Cr₂O₇²⁻

Oxygen is −2.

The ion has charge −2.

Let chromium equal x.

2x + 7(−2) = −2

2x − 14 = −2

2x = 12

x = +6

Answer: chromium has oxidation number +6.

Worked Example 5: Oxygen in Hydrogen Peroxide

Find the oxidation number of oxygen in H₂O₂.

Hydrogen is +1.

Let oxygen equal x.

2(+1) + 2x = 0

2 + 2x = 0

x = −1

Answer: oxygen is −1 in a peroxide.

This example shows why memorizing the major exceptions is important.

Worked Example 6: Oxygen in OF₂

Fluorine is always −1 in its compounds.

Let oxygen equal x.

x + 2(−1) = 0

x = +2

Answer: oxygen has oxidation number +2.

Oxygen does not automatically have an oxidation number of −2. Fluorine takes priority because it is more electronegative.

Average Oxidation Numbers and Mixed Oxidation States

Sometimes a calculation gives a fractional oxidation number.

Consider magnetite:

Fe₃O₄

Oxygen contributes:

4(−2) = −8

Therefore, the three iron atoms must total +8.

Average oxidation number of Fe = +8/3

This does not mean each iron atom literally has an oxidation state of +2.667.

Magnetite is a mixed-valence compound containing both Fe(II) and Fe(III). The fraction +8/3 is an average produced by the empirical formula.

Fractional oxidation numbers often signal averaging over equivalent positions, mixed oxidation states, delocalized bonding, or a simplified representation of a more complicated structure.

Using the Periodic Table to Find Oxidation Numbers

The periodic table helps predict oxidation numbers because elements in the same group have similar valence-electron configurations.

However, the periodic table alone usually cannot tell the exact oxidation number of an element in a particular compound. The chemical formula and bonding partners are also necessary.

Common Main-Group Oxidation Numbers

These values describe trends, not absolute rules.

For example, thallium commonly forms both +1 and +3 compounds, while lead commonly forms +2 and +4 compounds.

Fluorine is an important exception within Group 17. It has an oxidation number of −1 in compounds and does not show the positive oxidation numbers common to chlorine, bromine, and iodine.

Some noble gases also form compounds. Xenon, for example, occurs in positive oxidation states including +2, +4, +6, and +8.

Transition Metal Oxidation Numbers

Transition metals frequently have multiple oxidation numbers because their outer s and d electrons have relatively similar energies.

Common examples include:

• Titanium: +2, +3, +4
• Vanadium: +2, +3, +4, +5
• Chromium: +2, +3, +6
• Manganese: +2, +3, +4, +6, +7
• Iron: +2, +3
• Cobalt: +2, +3
• Nickel: +2
• Copper: +1, +2

Less common oxidation states occur for many of these elements.

This variability is why Stock nomenclature is particularly useful for transition-metal compounds.

FeCl₂ is iron(II) chloride.

FeCl₃ is iron(III) chloride.

The Roman numeral tells the oxidation number of iron.

Which Elements Have the Fewest Oxidation States?

The answer depends on whether only common compounds or every known exotic compound is counted.

Fluorine provides the simplest case. Its oxidation number is −1 in compounds and 0 in elemental F₂.

Alkali metals are also extremely consistent in ordinary chemistry. Sodium and potassium, for example, occur overwhelmingly as +1 in compounds and 0 as elemental metals.

Among the familiar first-row transition metals, scandium chemistry is dominated by +3, while zinc chemistry is dominated by +2.

Rare and unusual compounds can complicate statements that an element has “only” one oxidation state, so it is safer to distinguish common chemistry from all experimentally known compounds.

Which Elements Have the Most Oxidation States?

Transition metals near the middle of the d block tend to display especially wide ranges of oxidation numbers.

Manganese is the classic introductory example. Compounds containing manganese are known across a broad range of oxidation states, with +2, +3, +4, +6, and +7 particularly important in general chemistry. More unusual manganese compounds extend the range still further.

Having many oxidation states is not the same thing as having the highest oxidation number.

The highest experimentally identified formal oxidation state is +9 for iridium in the gas-phase ion IrO₄⁺, reported in 2014. This is an unusual research example rather than a typical oxidation state encountered in introductory chemistry. (nature.com)

For everyday chemistry, values such as Mn(VII), Cr(VI), Os(VIII), and Ru(VIII) are more representative examples of very high positive oxidation states.

Oxidation Numbers and Redox Reactions

Changes in oxidation number identify oxidation and reduction.

Oxidation: oxidation number increases.

Reduction: oxidation number decreases.

Consider:

Zn + Cu²⁺ → Zn²⁺ + Cu

Zinc changes from:

0 → +2

Its oxidation number increases, so zinc is oxidized.

Copper changes from:

+2 → 0

Its oxidation number decreases, so copper is reduced.

This leads to the mnemonic:

OIL RIG

Oxidation Is Loss of electrons.

Reduction Is Gain of electrons.

The oxidation-number method expresses the same idea by tracking changes in assigned oxidation states.

Using Oxidation Numbers to Balance Redox Reactions

Oxidation numbers provide a convenient way to determine how many electrons are transferred.

Consider the reaction between Fe²⁺ and permanganate ion in acidic solution.

Fe²⁺ → Fe³⁺

Iron changes from +2 to +3, a loss of one electron.

MnO₄⁻ → Mn²⁺

Manganese changes from +7 to +2, a gain of five electrons.

Therefore, five Fe²⁺ ions are required for every MnO₄⁻ ion.

The balanced equation in acidic solution is:

MnO₄⁻ + 5 Fe²⁺ + 8 H⁺ → Mn²⁺ + 5 Fe³⁺ + 4 H₂O

The total increase in iron oxidation numbers equals the total decrease in manganese oxidation number.

Oxidation Numbers in Organic Chemistry

Oxidation numbers also help describe oxidation and reduction of carbon.

For a carbon atom:

• Each C−H bond contributes −1 to carbon.
• Each C−O, C−N, or C−halogen bond generally contributes +1 per bond order when the other atom is more electronegative.
• A C−C bond contributes 0.

For example, carbon in methane, CH₄, has an oxidation number of −4.

Carbon in carbon dioxide, CO₂, has an oxidation number of +4.

Combustion therefore oxidizes carbon from a relatively low oxidation state toward a higher one.

For molecules containing several carbon atoms, individual carbon atoms can have different oxidation numbers. An average oxidation number for carbon may hide these differences.

Oxidation Number Notation

When writing an oxidation number as a numerical value, the sign normally comes first:

+2

+3

−1

−2

In Stock nomenclature, oxidation numbers are written as Roman numerals in parentheses after the element name:

iron(II) chloride

iron(III) oxide

copper(I) oxide

copper(II) sulfate

The Roman numeral represents oxidation state, not the number of atoms in the formula.

Common Misconceptions About Oxidation Numbers

Misconception 1: Oxidation Number Is the Actual Charge on an Atom

Not usually.

For a monatomic ion, oxidation number and charge are the same. But in a covalent molecule, oxidation number is normally a formal bookkeeping value.

Misconception 2: Oxidation Number and Formal Charge Are the Same

They use different rules.

Formal charge divides bonding electrons equally. Oxidation number assigns bonding electrons according to electronegativity.

Misconception 3: Oxygen Is Always −2

Oxygen is usually −2, but important exceptions include peroxides, superoxides, and compounds with fluorine.

Misconception 4: Hydrogen Is Always +1

Hydrogen is usually +1 when bonded to nonmetals but is often −1 when bonded to sufficiently electropositive elements, as in NaH.

Misconception 5: Metals Always Have Positive Oxidation Numbers

Positive oxidation numbers dominate ordinary metal chemistry, but unusual compounds containing zero or negative oxidation states of metals are known.

Misconception 6: Oxidation Requires Oxygen

Oxidation means an increase in oxidation number or, equivalently in many reactions, loss of electrons.

Oxygen does not have to participate.

For example:

Fe²⁺ → Fe³⁺ + e⁻

is oxidation even though oxygen is absent.

Misconception 7: A Fractional Oxidation Number Means an Atom Has a Fractional Charge

Usually not.

A fractional oxidation number often represents an average over several atoms, as in Fe₃O₄.

Misconception 8: The Periodic Table Gives a Single Oxidation Number for Every Element

Many elements have several possible oxidation numbers. This is particularly true for transition metals and heavier main-group elements.

Study Tips for Oxidation Numbers

Start with the rules that rarely change:

1. Free element = 0
2. Monatomic ion = ion charge
3. Group 1 = +1
4. Group 2 = +2
5. Fluorine = −1
6. Oxygen = usually −2
7. Hydrogen = usually +1
8. Sum = overall charge

Then learn the main exceptions for oxygen and hydrogen.

For an unknown element, use algebra rather than trying to memorize every oxidation number.

The most useful final check is simple: Do all the oxidation numbers add up to the charge of the species?

Frequently Asked Questions

What is an oxidation number?

An oxidation number is a formal number assigned to an atom that represents its degree of oxidation. It is determined by assigning bonding electrons according to electronegativity and other agreed rules.

Is oxidation number the same as oxidation state?

For most purposes, yes. IUPAC treats the terms as largely synonymous. Oxidation state emphasizes the chemical condition, while oxidation number emphasizes its numerical value.

Can an oxidation number be negative?

Yes. Negative oxidation numbers occur when an atom is formally assigned additional bonding electrons. Oxygen is commonly −2, chlorine is commonly −1, and nitrogen is −3 in NH₃.

Can an oxidation number be zero?

Yes. Every element in its uncombined elemental form has oxidation number 0.

Examples include O₂, H₂, Fe, Cu, S₈, and Cl₂.

Can oxidation numbers be fractions?

Average oxidation numbers can be fractional. This often occurs in mixed-valence compounds or formulas that average several atomic environments.

What is the oxidation number of oxygen?

Oxygen is usually −2. It is −1 in peroxides, has fractional negative values in some oxygen species such as superoxides, and has a positive oxidation number when bonded to fluorine.

What is the oxidation number of hydrogen?

Hydrogen is usually +1. It is typically −1 in compounds where it bonds to elements that are less electronegative than hydrogen, including many metal hydrides.

What is the oxidation number of fluorine?

Fluorine has oxidation number −1 in its compounds and 0 in elemental F₂.

What is the oxidation number of a monatomic ion?

It equals the charge of the ion. For example, Fe³⁺ has oxidation number +3 and Cl⁻ has oxidation number −1.

What is the difference between ionic charge and oxidation number?

For a monatomic ion they are identical. In molecules and polyatomic ions, oxidation number is often a formal value assigned to individual atoms and does not necessarily represent an actual localized electrical charge.

How do you know whether an atom is oxidized or reduced?

Compare its oxidation number before and after the reaction.

An increase means oxidation.

A decrease means reduction.

Which element has the highest oxidation number?

The highest experimentally identified formal oxidation state is +9 for iridium in IrO₄⁺. This is a specialized gas-phase species rather than a common compound.

Why do transition metals have several oxidation numbers?

Transition metals have s and d electrons with relatively similar energies. Different numbers of these electrons can participate in bonding or be removed, allowing many transition metals to form compounds in several oxidation states.

Glossary

Oxidation: An increase in oxidation number or loss of electrons.

Oxidation number: A number assigned to an atom that describes its formal degree of oxidation.

Oxidation state: The formal state of oxidation of an atom, generally synonymous with oxidation number.

Reduction: A decrease in oxidation number or gain of electrons.

Redox reaction: A reaction involving both oxidation and reduction.

Formal charge: A bookkeeping charge calculated by dividing covalent bonding electrons equally between bonded atoms.

Valence electron: An electron in the outer part of an atom that participates in chemical bonding and reactions.

Electronegativity: The tendency of an atom to attract bonding electrons.

Peroxide: A compound containing an O−O unit in which oxygen commonly has oxidation number −1.

Stock nomenclature: A naming system that uses Roman numerals to specify oxidation states, as in iron(III) chloride.

References

• Ellis, John E. (2006). “Adventures with Substances Containing Metals in Negative Oxidation States”. Inorganic Chemistry. 45 (8): 3167–3186. doi:10.1021/ic052110i
• Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. doi:10.1016/C2009-0-30414-6. ISBN 978-0-08-037941-8.
• International Union of Pure and Applied Chemistry (2009). The IUPAC Compendium of Chemical Terminology: The Gold Book. Blackwell Scientific. doi:10.1351/goldbook.O04365. ISBN 978-0-9678550-9-7.
• Karen, P.; McArdle, P.; Takats, J. (2014). “Toward a comprehensive definition of oxidation state (IUPAC Technical Report)”. Pure Appl. Chem. 86 (6): 1017–1081. doi:10.1515/pac-2013-0505