Common Ions and Their Charges: A Reference List

Writing formulas, naming compounds, balancing ionic equations and predicting precipitates all depend on knowing the charges of common ions. This reference list gathers the ions you’ll meet most often, grouped so the patterns are easy to see. Many charges can be worked out from the periodic table rather than memorised; the polyatomic ions are the ones that genuinely need learning. Each section explains the reasoning, so the list makes sense rather than being a random table.

How ions get their charges

An ion is an atom or group of atoms with an overall charge, because it has gained or lost electrons (see what is an ion?).

• Cations are positive (fewer electrons than protons). Most metals form cations.
• Anions are negative (more electrons than protons). Most non-metals form anions.

For main-group elements, the charge usually comes from the group number: atoms lose or gain electrons to reach the electron arrangement of the nearest noble gas.

1. Main-group cations (predictable from the group)

Group	Charge	Ions
1	+1	H⁺, Li⁺, Na⁺, K⁺, Rb⁺, Cs⁺
2	+2	Be²⁺, Mg²⁺, Ca²⁺, Sr²⁺, Ba²⁺
13	+3	Al³⁺ (and Ga³⁺)

Why: group 1 atoms have one outer electron, group 2 two, and aluminium three. Losing them leaves a full shell. See alkali metals and alkaline earth metals.

2. Main-group anions (predictable from the group)

Simple anions are named with the ending -ide:

Group	Charge	Ions
15	−3	N³⁻ (nitride), P³⁻ (phosphide)
16	−2	O²⁻ (oxide), S²⁻ (sulfide), Se²⁻ (selenide)
17	−1	F⁻ (fluoride), Cl⁻ (chloride), Br⁻ (bromide), I⁻ (iodide)
1	−1	H⁻ (hydride), in compounds such as NaH

Why: these atoms gain enough electrons to fill their outer shell: group 17 needs one, group 16 two, group 15 three.

3. Transition and post-transition metal cations

These often have more than one common charge, which is why their compound names include a Roman numeral (see naming compounds with transition metals).

Metal	Common ions	Notes
Iron	Fe²⁺, Fe³⁺	iron(II) pale green; iron(III) yellow-brown
Copper	Cu⁺, Cu²⁺	Cu²⁺ usually blue in solution
Chromium	Cr²⁺, Cr³⁺	Cr³⁺ is the most stable
Manganese	Mn²⁺	Pale pink; higher states in MnO₂ and MnO₄⁻
Cobalt	Co²⁺, Co³⁺	Co²⁺ pink in water
Nickel	Ni²⁺	Green in water
Zinc	Zn²⁺	Only one common charge
Silver	Ag⁺	Only one common charge
Mercury	Hg₂²⁺, Hg²⁺	Mercury(I) exists as a pair of atoms
Tin	Sn²⁺, Sn⁴⁺
Lead	Pb²⁺, Pb⁴⁺	Pb²⁺ is more common

Transition metal charges can’t be predicted simply from group number (see transition metals); learn the common ones, and read others from Roman numerals in names.

4. Polyatomic cations

Ion	Formula	Charge
Ammonium	NH₄⁺	+1
Oxonium (hydronium)	H₃O⁺	+1

Ammonium behaves much like a group 1 metal ion in its compounds: ammonium salts are almost all soluble (see dative covalent bonds for how it forms).

5. Polyatomic anions (learn these)

A polyatomic ion is a group of covalently bonded atoms carrying an overall charge. It stays together as a unit in reactions and in formulas.

Charge −1

Ion	Formula
Hydroxide	OH⁻
Nitrate	NO₃⁻
Nitrite	NO₂⁻
Hydrogencarbonate (bicarbonate)	HCO₃⁻
Hydrogensulfate	HSO₄⁻
Dihydrogenphosphate	H₂PO₄⁻
Chlorate	ClO₃⁻
Perchlorate	ClO₄⁻
Hypochlorite	ClO⁻
Permanganate (manganate(VII))	MnO₄⁻
Cyanide	CN⁻
Thiocyanate	SCN⁻
Ethanoate (acetate)	CH₃COO⁻

Charge −2

Ion	Formula
Carbonate	CO₃²⁻
Sulfate	SO₄²⁻
Sulfite	SO₃²⁻
Thiosulfate	S₂O₃²⁻
Hydrogenphosphate	HPO₄²⁻
Chromate	CrO₄²⁻
Dichromate	Cr₂O₇²⁻
Peroxide	O₂²⁻
Ethanedioate (oxalate)	C₂O₄²⁻

Charge −3

Ion	Formula
Phosphate	PO₄³⁻
Phosphite	PO₃³⁻

Patterns that make polyatomic ions easier

-ate and -ite

For oxyanions (ions containing oxygen), -ate has one more oxygen than -ite, with the same charge:

• nitrate NO₃⁻ / nitrite NO₂⁻
• sulfate SO₄²⁻ / sulfite SO₃²⁻
• chlorate ClO₃⁻ / chlorite ClO₂⁻

The prefixes per- (one more O than -ate) and hypo- (one fewer O than -ite) extend the series for chlorine, bromine and iodine (see naming oxyanions).

Adding hydrogen

Adding H⁺ to a polyatomic anion reduces its negative charge by one, and the name gains “hydrogen”:

• CO₃²⁻ carbonate → HCO₃⁻ hydrogencarbonate
• SO₄²⁻ sulfate → HSO₄⁻ hydrogensulfate
• PO₄³⁻ phosphate → HPO₄²⁻ hydrogenphosphate → H₂PO₄⁻ dihydrogenphosphate

Period 3 oxyanions

Look at the “-ate” ions of period 3 elements: SiO₄⁴⁻, PO₄³⁻, SO₄²⁻, ClO₄⁻. Each has four oxygens, and the charge goes −4, −3, −2, −1 as you move across. (ClO₄⁻ is perchlorate, because chlorate is ClO₃⁻.) Patterns like this reduce the amount of pure memorisation.

How to learn the list efficiently

The main-group ions don’t need memorising at all — the periodic table gives them to you. Focus your effort on the polyatomic ions, and learn them in a smart order:

1. Start with a core of seven that appear constantly: hydroxide (OH⁻), nitrate (NO₃⁻), sulfate (SO₄²⁻), carbonate (CO₃²⁻), phosphate (PO₄³⁻), ammonium (NH₄⁺) and hydrogencarbonate (HCO₃⁻).
2. Derive the -ite ions from the -ate ions: same charge, one fewer oxygen. Nitrite and sulfite then come for free.
3. Derive the hydrogen ions by adding H⁺ and reducing the charge by one: hydrogensulfate, hydrogenphosphate, dihydrogenphosphate.
4. Add the rest gradually as you meet them in practicals and questions: permanganate, dichromate, thiosulfate, ethanoate.

Flashcards with the name on one side and the formula and charge on the other work well, especially if you test yourself in both directions. Writing the formulas of real compounds that contain each ion — Na₂CO₃, Ca(OH)₂, (NH₄)₂SO₄ — fixes the charges in memory far better than learning ions on their own. Revisit the list briefly each week, testing yourself in both directions, until recall is instant and automatic.

A quick check: charges from compounds you know

You can often recover an ion’s charge from a familiar compound:

• NaOH → Na⁺, so hydroxide is OH⁻.
• CaCO₃ → Ca²⁺, so carbonate is CO₃²⁻.
• H₂SO₄ → two H⁺, so sulfate is SO₄²⁻.
• H₃PO₄ → three H⁺, so phosphate is PO₄³⁻.
• HNO₃ → one H⁺, so nitrate is NO₃⁻.

Using ion charges

Once you know the charges, you can:

• Write formulas by balancing charges (see the criss-cross method and how to write chemical formulas).
• Name compounds (see how to name ionic compounds).
• Predict precipitates in reactions using solubility patterns (see solubility rules).
• Calculate molar masses of ionic compounds (see the molar mass calculator).

Common mistakes

• Confusing -ide, -ite and -ate: sulfide S²⁻, sulfite SO₃²⁻, sulfate SO₄²⁻.
• Giving transition metals a single fixed charge — check the Roman numeral.
• Forgetting that polyatomic ions stay together — you can’t change SO₄²⁻ to balance an equation.
• Mixing up ammonium (NH₄⁺) and ammonia (NH₃) — ammonia is a neutral molecule.

Key takeaways

• Main-group charges follow the group: +1, +2, +3 for groups 1, 2, 13; −3, −2, −1 for groups 15, 16, 17.
• Transition metals often have several charges; names show them with Roman numerals.
• Polyatomic ions must be learned, but patterns help: -ate has one more O than -ite, and adding H⁺ reduces the charge by one.
• Recover charges from familiar compounds like NaOH, CaCO₃, H₂SO₄.