Ion Charge / Oxidation State Lookup
Look up common oxidation states and simple oxidation number rules.
What you can search
Use the dropdown for common examples or type an element name, symbol, or atomic number. The lookup shows common oxidation states, periodic position, and a rule hint so you can connect a charge value to the way it is used in formulas.
| Search | Example | What to notice |
|---|---|---|
| Symbol | Fe | Iron often uses +2 and +3 |
| Name | oxygen | Oxygen is usually -2 |
| Atomic number | 17 | Chlorine often forms -1 but has positive states in oxyanions |
| Group rule | Na, Mg, Al | Many main-group metal charges follow group patterns |
Rules used
Worked examples
FeCl3: chlorine is usually -1. Three chlorine atoms total -3, so iron must be +3.
SO4^2-: oxygen is usually -2. Four oxygen atoms total -8. The ion charge is -2, so sulfur is +6.
Na2O: sodium is +1. Two sodium atoms total +2, so oxygen is -2.
How to use the charge chips
The charge chips show common states, not every possible state in advanced chemistry. For simple monatomic ions, the ion charge often matches the oxidation state. For covalent compounds and polyatomic ions, oxidation state is a bookkeeping tool: assign the common fixed values first, then solve the missing element by charge balance.
Where this lookup is useful
- Naming compounds with Roman numerals, such as iron(III) chloride.
- Balancing redox half-reactions and tracking electron transfer.
- Predicting simple ionic formulas from charges.
- Checking polyatomic ion oxidation numbers.
- Reviewing periodic trends before quizzes and lab reports.
Common mistakes
- Assuming every transition metal has only one charge.
- Forgetting free elements such as O2, H2, Fe(s), and Cl2 have oxidation state 0.
- Using -2 for oxygen in peroxides without checking the exception.
- Making oxidation numbers add to 0 for a polyatomic ion instead of to the ion charge.
- Confusing oxidation state with formal charge in Lewis structures.
Result checking
After assigning oxidation states, add every atom contribution. Neutral compounds must total 0. Ions must total the written charge. If the total does not match, check subscripts first, then check whether an exception such as peroxide, hydride, or a transition-metal Roman numeral is involved.
Related Chemistry Tools
FAQs
How to calculate oxidation number?
Apply the standard rules in order. (1) A free element has oxidation number 0 (Na, O2, H2, S8). (2) A monatomic ion has oxidation number equal to its charge (Na+ = +1, Cl- = -1). (3) Hydrogen is +1 with non-metals and -1 in metal hydrides such as NaH and CaH2. (4) Oxygen is -2 in most compounds, -1 in peroxides (H2O2, Na2O2), -1/2 in superoxides (KO2), and +2 in OF2. (5) Group 1 metals are +1, Group 2 are +2, Al is +3. (6) The sum of oxidation numbers equals 0 for a neutral compound and equals the charge for a polyatomic ion. Worked examples: H2SO4 → 2(+1) + S + 4(-2) = 0 → S = +6. KMnO4 → +1 + Mn + 4(-2) = 0 → Mn = +7. Cr2O7^2- → 2 Cr + 7(-2) = -2 → Cr = +6.
Which compound has the atom with the highest oxidation number?
Among common compounds, oxidation states peak around +7 (KMnO4, HClO4) and reach +8 in rare oxides such as OsO4, RuO4, and XeO4. Group-by-group ceiling: Group 7 (Mn, Cl) → +7; Group 8 (Os, Ru, Xe) → +8 (the highest stable oxidation state observed in molecular compounds). In introductory courses, KMnO4 is the usual answer with Mn at +7. It is a strong oxidizer; in acidic solution it is reduced to Mn^2+: MnO4- + 8 H+ + 5 e- → Mn^2+ + 4 H2O. Heuristic for spotting high oxidation states: look for the central atom surrounded by many highly electronegative atoms — usually oxygens or fluorines.
How to find the oxidation state of an element?
Set the unknown to x, apply the standard oxidation-number rules to every other atom, and solve so the sum equals 0 (for a neutral compound) or the ion's charge (for an ion). Examples: HNO3 → +1 + x + 3(-2) = 0 → x = +5; PO4^3- → x + 4(-2) = -3 → x = +5; SO3^2- → x + 3(-2) = -2 → x = +4; CO3^2- → x + 3(-2) = -2 → x = +4. For organic molecules, treat C-H bonds as if H takes its usual +1, C-O bonds as if O takes -2, and C-C bonds as electron-neutral; the central carbon's oxidation state then follows from the local environment.
What is the oxidation number of oxygen?
Usually -2, with a few well-known exceptions. -2 in most oxides: H2O, CO2, MgO, Na2O, SiO2, Fe2O3, SO3. -1 in peroxides: H2O2, Na2O2, BaO2 (O-O bond means each O has only one electron from the other O). -1/2 in superoxides: KO2, RbO2, CsO2 (an unpaired electron is shared over O2-). Positive when bonded to fluorine: O = +2 in OF2 and O = +1 in O2F2 (because F is more electronegative). 0 in elemental forms: O2 and O3 (ozone). Always check the formula before assuming -2 — peroxides and oxygen fluorides are the common exam traps.
What does it mean when a substance is oxidizing?
An oxidizing agent (oxidant) is a species that gains electrons in a reaction — that is, it is itself reduced and causes the other species to be oxidized. Common oxidizing agents: O2, F2, Cl2, KMnO4, K2Cr2O7, HNO3, H2O2, O3. They typically contain an element in a high oxidation state (Mn^+7 in MnO4-, Cr^+6 in Cr2O7^2-) or a highly electronegative element (F2, Cl2, O2). Mnemonic: OIL RIG — Oxidation Is Loss of electrons, Reduction Is Gain. Examples: 2 Mg + O2 → 2 MgO (Mg loses electrons, O gains; O2 is the oxidant). MnO4- + 8 H+ + 5 e- → Mn^2+ + 4 H2O (Mn goes from +7 to +2, gaining 5 electrons; KMnO4 is the oxidant). Practical use: chlorine and ozone for water disinfection; KMnO4 in titrimetric analysis; concentrated HNO3 to dissolve metals.
What is the oxidation number for O2?
Zero. Free elements always have oxidation number 0, regardless of how many atoms are bonded — including O2, N2, H2, the halogens (F2, Cl2, Br2, I2), and elemental allotropes such as P4, S8, and graphite. Reasoning: in O2, both atoms are identical and equally electronegative, so the shared electrons are split evenly and neither atom has 'gained' or 'lost' anything relative to the free atom. Don't confuse O2 (0) with O in oxides like H2O or CO2 (-2). The change from 0 to -2 is exactly what makes O2 an oxidizing agent: 2 Mg + O2 → 2 MgO has Mg going from 0 → +2 and O from 0 → -2.
Does the mass change in an ion?
Technically yes — but by an amount too small to matter for chemistry. The electron mass is 9.11 × 10^-31 kg ≈ 5.49 × 10^-4 amu, about 1/1836 of a proton. So Na (22.990 amu) becomes Na+ at 22.989 amu when it loses one electron — a change near 0.0024%. Periodic-table atomic masses include the electron count of the neutral atom, and stoichiometry treats Na and Na+ as 22.99 g/mol indistinguishably. The picture is different in nuclear chemistry, where rearrangements involve protons and neutrons and the mass change becomes measurable as binding-energy difference (E = mc^2).
What event changes an atom into an ion?
An atom becomes an ion by gaining or losing electrons — a process called ionization. Losing electrons gives a cation. Metals are the typical electron donors: Na → Na+ + e-; Mg → Mg^2+ + 2 e-; Al → Al^3+ + 3 e-. Gaining electrons gives an anion. Non-metals near full octets are the typical acceptors: Cl + e- → Cl-; O + 2 e- → O^2-; N + 3 e- → N^3-. Triggers in practice: a chemical reaction (2 Na + Cl2 → 2 NaCl); dissolution of an ionic solid in water (NaCl → Na+(aq) + Cl-(aq)); an electric current during electrolysis; high-energy photons (photoionization, used in mass spectrometry). Through all of this the proton count is unchanged, so the element's identity stays the same — only the electron count and the net charge change.