Skip to main content
Chemistry LibreTexts

Oxidation States II

  • Page ID
  • \( \newcommand{\vecs}[1]{\overset { \scriptstyle \rightharpoonup} {\mathbf{#1}} } \)

    \( \newcommand{\vecd}[1]{\overset{-\!-\!\rightharpoonup}{\vphantom{a}\smash {#1}}} \)

    \( \newcommand{\id}{\mathrm{id}}\) \( \newcommand{\Span}{\mathrm{span}}\)

    ( \newcommand{\kernel}{\mathrm{null}\,}\) \( \newcommand{\range}{\mathrm{range}\,}\)

    \( \newcommand{\RealPart}{\mathrm{Re}}\) \( \newcommand{\ImaginaryPart}{\mathrm{Im}}\)

    \( \newcommand{\Argument}{\mathrm{Arg}}\) \( \newcommand{\norm}[1]{\| #1 \|}\)

    \( \newcommand{\inner}[2]{\langle #1, #2 \rangle}\)

    \( \newcommand{\Span}{\mathrm{span}}\)

    \( \newcommand{\id}{\mathrm{id}}\)

    \( \newcommand{\Span}{\mathrm{span}}\)

    \( \newcommand{\kernel}{\mathrm{null}\,}\)

    \( \newcommand{\range}{\mathrm{range}\,}\)

    \( \newcommand{\RealPart}{\mathrm{Re}}\)

    \( \newcommand{\ImaginaryPart}{\mathrm{Im}}\)

    \( \newcommand{\Argument}{\mathrm{Arg}}\)

    \( \newcommand{\norm}[1]{\| #1 \|}\)

    \( \newcommand{\inner}[2]{\langle #1, #2 \rangle}\)

    \( \newcommand{\Span}{\mathrm{span}}\) \( \newcommand{\AA}{\unicode[.8,0]{x212B}}\)

    \( \newcommand{\vectorA}[1]{\vec{#1}}      % arrow\)

    \( \newcommand{\vectorAt}[1]{\vec{\text{#1}}}      % arrow\)

    \( \newcommand{\vectorB}[1]{\overset { \scriptstyle \rightharpoonup} {\mathbf{#1}} } \)

    \( \newcommand{\vectorC}[1]{\textbf{#1}} \)

    \( \newcommand{\vectorD}[1]{\overrightarrow{#1}} \)

    \( \newcommand{\vectorDt}[1]{\overrightarrow{\text{#1}}} \)

    \( \newcommand{\vectE}[1]{\overset{-\!-\!\rightharpoonup}{\vphantom{a}\smash{\mathbf {#1}}}} \)

    \( \newcommand{\vecs}[1]{\overset { \scriptstyle \rightharpoonup} {\mathbf{#1}} } \)

    \( \newcommand{\vecd}[1]{\overset{-\!-\!\rightharpoonup}{\vphantom{a}\smash {#1}}} \)

    \(\newcommand{\avec}{\mathbf a}\) \(\newcommand{\bvec}{\mathbf b}\) \(\newcommand{\cvec}{\mathbf c}\) \(\newcommand{\dvec}{\mathbf d}\) \(\newcommand{\dtil}{\widetilde{\mathbf d}}\) \(\newcommand{\evec}{\mathbf e}\) \(\newcommand{\fvec}{\mathbf f}\) \(\newcommand{\nvec}{\mathbf n}\) \(\newcommand{\pvec}{\mathbf p}\) \(\newcommand{\qvec}{\mathbf q}\) \(\newcommand{\svec}{\mathbf s}\) \(\newcommand{\tvec}{\mathbf t}\) \(\newcommand{\uvec}{\mathbf u}\) \(\newcommand{\vvec}{\mathbf v}\) \(\newcommand{\wvec}{\mathbf w}\) \(\newcommand{\xvec}{\mathbf x}\) \(\newcommand{\yvec}{\mathbf y}\) \(\newcommand{\zvec}{\mathbf z}\) \(\newcommand{\rvec}{\mathbf r}\) \(\newcommand{\mvec}{\mathbf m}\) \(\newcommand{\zerovec}{\mathbf 0}\) \(\newcommand{\onevec}{\mathbf 1}\) \(\newcommand{\real}{\mathbb R}\) \(\newcommand{\twovec}[2]{\left[\begin{array}{r}#1 \\ #2 \end{array}\right]}\) \(\newcommand{\ctwovec}[2]{\left[\begin{array}{c}#1 \\ #2 \end{array}\right]}\) \(\newcommand{\threevec}[3]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \end{array}\right]}\) \(\newcommand{\cthreevec}[3]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \end{array}\right]}\) \(\newcommand{\fourvec}[4]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \\ #4 \end{array}\right]}\) \(\newcommand{\cfourvec}[4]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \\ #4 \end{array}\right]}\) \(\newcommand{\fivevec}[5]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \\ #4 \\ #5 \\ \end{array}\right]}\) \(\newcommand{\cfivevec}[5]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \\ #4 \\ #5 \\ \end{array}\right]}\) \(\newcommand{\mattwo}[4]{\left[\begin{array}{rr}#1 \amp #2 \\ #3 \amp #4 \\ \end{array}\right]}\) \(\newcommand{\laspan}[1]{\text{Span}\{#1\}}\) \(\newcommand{\bcal}{\cal B}\) \(\newcommand{\ccal}{\cal C}\) \(\newcommand{\scal}{\cal S}\) \(\newcommand{\wcal}{\cal W}\) \(\newcommand{\ecal}{\cal E}\) \(\newcommand{\coords}[2]{\left\{#1\right\}_{#2}}\) \(\newcommand{\gray}[1]{\color{gray}{#1}}\) \(\newcommand{\lgray}[1]{\color{lightgray}{#1}}\) \(\newcommand{\rank}{\operatorname{rank}}\) \(\newcommand{\row}{\text{Row}}\) \(\newcommand{\col}{\text{Col}}\) \(\renewcommand{\row}{\text{Row}}\) \(\newcommand{\nul}{\text{Nul}}\) \(\newcommand{\var}{\text{Var}}\) \(\newcommand{\corr}{\text{corr}}\) \(\newcommand{\len}[1]{\left|#1\right|}\) \(\newcommand{\bbar}{\overline{\bvec}}\) \(\newcommand{\bhat}{\widehat{\bvec}}\) \(\newcommand{\bperp}{\bvec^\perp}\) \(\newcommand{\xhat}{\widehat{\xvec}}\) \(\newcommand{\vhat}{\widehat{\vvec}}\) \(\newcommand{\uhat}{\widehat{\uvec}}\) \(\newcommand{\what}{\widehat{\wvec}}\) \(\newcommand{\Sighat}{\widehat{\Sigma}}\) \(\newcommand{\lt}{<}\) \(\newcommand{\gt}{>}\) \(\newcommand{\amp}{&}\) \(\definecolor{fillinmathshade}{gray}{0.9}\)

    Oxidation state is a number assigned to an element in a compound according to some rules. This number enable us to describe oxidation-reduction reactions, and balancing redox chemical reactions. When a covalent bond forms between two atoms with different electronegativities the shared electrons in the bond lie closer to the more electronegative atom:


    • The oxidation number of an atom is the charge that results when the electrons in a covalent bond are assigned to the more electronegative atom
    • It is the charge an atom would possess if the bonding were ionic

    In HCl (above) the oxidation number for the hydrogen would be +1 and that of the Cl would be -1


    For oxidation numbers we write the sign first to distinguish them from ionic (electronic) charges

    Oxidation States

    Oxidation state (or oxidation numbers) do not refer to real charges on the atoms, except in the case of actual ionic substances. Oxidation numbers can be determined using the following rules:

    1. The oxidation number for an element in its elemental form is 0 (holds true for isolated atoms and elemental substances which bond identical atoms: e.g. Cl2, etc)
    2. The oxidation number of a monoatomic ion is the same as its charge (e.g. oxidation number of Na+ = +1, and that of S2- is -2)
    3. In binary compounds (two different elements) the element with greater electronegativity is assigned a negative oxidation number equal to its charge in simple ionic compounds of the element (e.g., the chlorine in PCl3 is more electronegative than the phosphorous. In simple ionic compounds Cl has an ionic charge of 1-, thus, its oxidation state is -1)
    4. The sum of the oxidation numbers is zero for an electrically neutral compound and equals the overall charge for an ionic species.
    5. Alkali metals exhibit only an oxidation state of +1 in compounds
    6. Alkaline earth metals exhibit only an oxidation state of +2 in compounds

    Rules for Assigning Oxidation States

    The oxidation state (OS) of an element corresponds to the number of electrons, e-, that an atom loses, gains, or appears to use when joining with other atoms in compounds. In determining the OS of an atom, there are seven guidelines to follow:

    1. The OS of an individual atom is 0.
    2. The total OS of all atoms in: a neutral species is 0 and in an ion is equal to the ion charge.
    3. Group 1 metals have an OS of +1 and Group 2 an OS of +2
    4. The OS of fluorine is -1 in compounds
    5. Hydrogen generally has an OS of +1 in compounds
    6. Oxygen generally has an OS of -2 in compounds
    7. In binary metal compounds, Group 17 elements have an OS of -1, Group 16 of -2, and Group 15 of -3.

    (Note: The sum of the OSs is equal to zero for neutral compounds and equal to the charge for polyatomic ion species.)

    Example 20.1.1

    Determine the OSs of the elements in the following reactions:

    1. \(Fe_{(s)} + O_{2(g)} \rightarrow Fe_2O_{3(g)}\)
    2. \(Fe^{2+}\)
    3. \(Ag_{(s)} + H_2S \rightarrow Ag_2S_{(g)} + H_{2(g)}\)


    1. Fe and O2 are free elements; therefore, they each have an OS of 0 according to Rule #1. The product has a total OS equal to 0, and following Rule #6, O has an OS of -2, which means Fe has an OS of +3.
    2. The OS of Fe corresponds to its charge; therefore, the OS is +2.
    3. Ag has an OS of 0, H has an OS of +1 according to Rule #5, S has an OS of -2 according to Rule #7, and hence Ag in Ag2S has an OS of +1.
    Example 20.1.2

    Determine the OS of the bold element in each of the following:

    1. Na3PO3
    2. H2PO4-


    1. The oxidation numbers of Na and O are +1 and -2. Because sodium phosphite is neutral, the sum of the oxidation numbers must be zero. Letting x be the oxidation number of phosphorus, 0= 3(+1) + x + 3(-2). x=oxidation number of P= +3.
    2. Hydrogen and oxygen have oxidation numbers of +1 and -2. The ion has a charge of -1, so the sum of the oxidation numbers must be -1. Letting y be the oxidation number of phosphorus, -1= y + 2(+1) +4(-2), y= oxidation number of P= +5.
    Example 20.2.3

    Determine which element is oxidized and which element is reduced in the following reactions (be sure to include the OS of each):

    1. Zn + 2H+ → Zn2+ + H2
    2. 2Al + 3Cu2+→2Al3+ +3Cu
    3. CO32- + 2H+→ CO2 + H2O


    1. Zn is oxidized (Oxidation number: 0 → +2); H+ is reduced (Oxidation number: +1 → 0)
    2. Al is oxidized (Oxidation number: 0 → +3); Cu2+ is reduced (+2 → 0)
    3. This is not a redox reaction because each element has the same oxidation number in both reactants and products: O= -2, H= +1, C= +4.

    (For further discussion, see the article on oxidation numbers).

    An atom is oxidized if its oxidation number increases, the reducing agent, and an atom is reduced if its oxidation number decreases, the oxidizing agent. The atom that is oxidized is the reducing agent, and the atom that is reduced is the oxidizing agent. (Note: the oxidizing and reducing agents can be the same element or compound).

    Oxidation Numbers and Nomenclature

    Compounds of the alkali (oxidation number +1) and alkaline earth metals (oxidation number +2) are typically ionic in nature. Compounds of metals with higher oxidation numbers (e.g., tin +4) tend to form molecular compounds

    • In ionic and covalent molecular compounds usually the less electronegative element is given first.
    • In ionic compounds the names are given which refer to the oxidation (ionic) state
    • In molecular compounds the names are given which refer to the number of molecules present in the compound

    Figure 20.1.1: Example of nomenclature based on oxidation states.




    magnesium hydride


    dihydrogen sulfide


    iron(II) fluoride


    oxygen difluoride


    manganese(III) oxide


    dichlorine trioxide

    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.

    Oxidation-Reduction Reactions

    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).


    Figure 20.1.1: A thermite reaction taking place on a cast iron skillet. A thermite reaction, using about 110 g of the mixture, taking place. The cast-iron skillet was destroyed in the process. Image by Schuyler S and used with permission by Wikipedia.

    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 \(Al_2O_3\) (Figure 20.1.1).

    \[Fe_2O_{3(s)} + 2Al_{(s)} \rightarrow Al_2O_{3(s)} + 2Fe_{(l)}\]

    Another example of the redox reaction is the reaction between zinc and copper sulfate.

    Example 20.1.4: Identifying Oxidized Elements

    Using the equations from the previous examples, determine what is oxidized in the following reaction.

    \[Zn + 2H^+ \rightarrow Zn^{2+} + H_2\]


    The OS of H changes from +1 to 0, and the OS of Zn changes from 0 to +2. Hence, Zn is oxidized and acts as the reducing agent.

    Example 20.1.5: Identifying Reduced Elements

    What is reduced species in this reaction?

    \[Zn + 2H^+ \rightarrow Zn^{2+} + H_2\]


    The OS of H changes from +1 to 0, and the OS of Zn changes from 0 to +2. Hence, H+ ion is reduced and acts as the oxidizing agent.

    Combination Reactions

    Combination reactions are among the simplest redox reactions and, as the name suggests, involves "combining" elements to form a chemical compound. As usual, oxidation and reduction occur together. The general equation for a combination reaction is given below:

    \[A + B \rightarrow AB \]

    Example 20.1.6: Combination Reaction

    Equation: H2 + O2 → H2O
    Calculation: 0 + 0 → (2)(+1) + (-2) = 0
    Explanation: In this equation both H2 and O2 are free elements; following Rule #1, their OSs are 0. The product is H2O, which has a total OS of 0. According to Rule #6, the OS of oxygen is usually -2. Therefore, the OS of H in H2O must be +1.

    Decomposition Reactions

    A decomposition reaction is the reverse of a combination reaction, the breakdown of a chemical compound into individual elements:

    \[AB \rightarrow A + B\]

    Example 20.1.7: Decomposition Reaction

    Consider the decomposition of water:

    \[H_2O \rightarrow H_2 + O_2\]

    Calculation: (2)(+1) + (-2) = 0 → 0 + 0
    Explanation: In this reaction, water is "decomposed" into hydrogen and oxygen. As in the previous example the H2O has a total OS of 0; thus, according to Rule #6 the OS of oxygen is usually -2, so the OS of hydrogen in H2O must be +1.

    Single Replacement Reactions

    A single replacement reaction involves the "replacing" of an element in the reactants with another element in the products:

    \[A + BC \rightarrow AB + C\]

    Example 20.1.8: Single Replacement Reaction


    \[Cl_2 + Na\underline{Br} \rightarrow Na\underline{Cl} + Br_2\]

    Calculation: (0) + ((+1) + (-1) = 0) -> ((+1) + (-1) = 0) + 0
    Explanation: In this equation, Br is replaced with Cl, and the Cl atoms in Cl2 are reduced, while the Br ion in NaBr is oxidized.

    Double Replacement Reactions

    A double replacement reaction is similar to a double replacement reaction, but involves "replacing" two elements in the reactants, with two in the products:

    \[AB + CD \rightarrow AD + CB \]

    Example 20.1.9: Double Replacement Reaction

    Equation: Fe2O3 + HCl → FeCl3 + H2O
    Explanation: In this equation, Fe and H trade places, and oxygen and chlorine trade places.

    Combustion Reactions

    Combustion reactions almost always involve oxygen in the form of O2, and are almost always exothermic, meaning they produce heat. Chemical reactions that give off light and heat and light are colloquially referred to as "burning."

    \[C_xH_y + O_2 \rightarrow CO_2 + H_2O\]

    Although combustion reactions typically involve redox reactions with a chemical being oxidized by oxygen, many chemicals "burn" in other environments. For example, both titanium and magnesium burn in nitrogen as well:

    \[ 2Ti(s) + N_{2}(g) \rightarrow 2TiN(s)\]

    \[3 Mg(s) + N_{2}(g) \rightarrow Mg_3N_{2}(s) \]

    Moreover, chemicals can be oxidized by other chemicals than oxygen, such as Cl2 or F2; these processes are also considered combustion reactions

    Disproportionation Reactions

    Disproportionation Reactions: In some redox reactions a single substance can be both oxidized and reduced. These are known as disproportionation reactions, with the following general equation:

    \[2A \rightarrow A^{+n} + A^{-n}\]

    Where n is the number of electrons transferred. Disproportionation reactions do not need begin with neutral molecules, and can involve more than two species with differing oxidation states (but rarely).

    Example 20.1.10: Disproportionation Reaction

    Disproportionation reactions have some practical significance in everyday life, including the reaction of hydrogen peroxide, H2O2 poured over a cut. This a decomposition reaction of hydrogen peroxide, which produces oxygen and water. Oxygen is present in all parts of the chemical equation and as a result it is both oxidized and reduced. The reaction is as follows:

    \[2H_2O_{2}(aq) \rightarrow 2H_2O(l) + O_{2}(g)\]

    Explanation: On the reactant side, H has an OS of +1 and O has an OS of -1, which changes to -2 for the product H2O (oxygen is reduced), and 0 in the product O2 (oxygen is oxidized).


    • Remember the 7 rules of OSs (these are vital to understanding redox reactions)
    • Oxidation signifies a loss of electrons and reduction signifies a gain of electrons.
    • Balancing redox reactions is an important step that changes in neutral, basic, and acidic solutions.
    • Remember the various types of redox reactions
      • Combination and decomposition
      • Displacement reactions (single and double)
      • Combustion
      • Disproportionation
    • The oxidizing agent undergoes reduction and the reducing agent undergoes oxidation.


    1. Petrucci, et al. General Chemistry: Principles & Modern Applications. 9th ed. Upper Saddle River, New Jersey: Pearson/Prentice Hall, 2007.
    2. Sadava, et al. Life: The Science of Biology. 8th ed. New York, NY. W.H. Freeman and Company, 2007

    Contributors and Attributions

    • Christopher Spohrer (UCD), Christina Breitenbuecher (UCD), Luvleen Brar (UCD)

    Oxidation States II is shared under a CC BY-NC-SA 4.0 license and was authored, remixed, and/or curated by LibreTexts.

    • Was this article helpful?