How to Predict the Charge of an Ion from the Periodic Table

To write the formula of an ionic compound, name a salt or balance an ionic equation, you need to know the charges on the ions involved. You don’t have to memorise dozens of charges: for most main-group elements, the periodic table tells you the answer directly. This guide gives a step-by-step method, explains why it works, and shows how to handle the trickier cases.

The quick rule

Group	Typical ion charge	Examples
1	+1	Li⁺, Na⁺, K⁺
2	+2	Mg²⁺, Ca²⁺, Ba²⁺
13	+3	Al³⁺ (Ga³⁺)
14	usually form covalent bonds, not simple ions	(C, Si)
15	−3	N³⁻, P³⁻
16	−2	O²⁻, S²⁻
17	−1	F⁻, Cl⁻, Br⁻, I⁻
18	no ions (noble gases)	—

Why the rule works

Atoms tend to form ions with the same electron arrangement as the nearest noble gas, because that arrangement is especially stable and the energy balance in forming ionic compounds favours it. See how atoms become ions.

• Metals (left side) have few outer electrons. They lose them to reach the previous noble gas configuration, forming positive ions. The charge equals the number of electrons lost.
• Non-metals (right side) have nearly full outer shells. They gain electrons to reach the next noble gas configuration, forming negative ions. The charge equals minus the number gained.

For example:

• Magnesium (2, 8, 2) loses 2 → Mg²⁺ (2, 8), like neon.
• Oxygen (2, 6) gains 2 → O²⁻ (2, 8), like neon.

See electron shells explained and how to find valence electrons.

Step-by-step method

1. Find the element on the periodic table and note its group.
2. Metal or non-metal? Metals are on the left and centre; non-metals on the right. See metals, non-metals and metalloids.
3. For metals in groups 1, 2, 13: charge = +(group number) for groups 1–2, and +3 for group 13.
4. For non-metals in groups 15–17: charge = group number − 18 (so 15 → −3, 16 → −2, 17 → −1).
5. For transition metals: look up the charge or read it from the Roman numeral in the compound’s name.
6. Check that the resulting ion has a noble gas configuration (for main-group elements).

Worked examples: single ions

Element	Group	Metal/non-metal	Ion
Potassium	1	metal	K⁺
Calcium	2	metal	Ca²⁺
Aluminium	13	metal	Al³⁺
Nitrogen	15	non-metal	N³⁻ (nitride)
Sulfur	16	non-metal	S²⁻ (sulfide)
Bromine	17	non-metal	Br⁻ (bromide)

Note the naming: negative ions of single elements end in -ide.

Using ion charges to write formulas

In an ionic compound, the total positive and negative charges must balance to zero.

Method: the “swap and drop” (criss-cross) shortcut

1. Write each ion with its charge.
2. Use the size of each charge as the subscript of the other ion.
3. Simplify to the lowest whole-number ratio.

Example 1: aluminium oxide Al³⁺ and O²⁻ → Al₂O₃ (2 × +3 = +6; 3 × −2 = −6 ✓)

Example 2: magnesium chloride Mg²⁺ and Cl⁻ → MgCl₂

Example 3: calcium oxide Ca²⁺ and O²⁻ → Ca₂O₂ → simplify → CaO

Example 4: sodium nitride Na⁺ and N³⁻ → Na₃N

See how to name ionic compounds.

Transition metals: variable charges

Transition metals can often form more than one ion, because their d electrons can be lost in different numbers. Their charges can’t be read simply from the group number.

Metal	Common ions
Iron	Fe²⁺, Fe³⁺
Copper	Cu⁺, Cu²⁺
Chromium	Cr²⁺, Cr³⁺ (and higher oxidation states in compounds)
Manganese	Mn²⁺ (and higher oxidation states in compounds)
Zinc	Zn²⁺ only
Silver	Ag⁺ only

The Roman numeral in a name gives the charge:

• iron(II) chloride: Fe²⁺ → FeCl₂
• iron(III) chloride: Fe³⁺ → FeCl₃
• copper(I) oxide: Cu⁺ → Cu₂O
• copper(II) sulfate: Cu²⁺ → CuSO₄

Zinc and silver form only one common ion each, so no Roman numeral is used. See transition metals and electron configurations of ions.

Special cases

Hydrogen

Hydrogen can form H⁺ (as in acids, although in water it’s always attached to water molecules as H₃O⁺) or H⁻ (the hydride ion, in compounds with very reactive metals, such as sodium hydride, NaH). See what is an acid?.

Group 14

Carbon and silicon usually form covalent bonds rather than C⁴⁺ or C⁴⁻ ions, because gaining or losing four electrons would require too much energy. Tin and lead, lower in the group, form Sn²⁺/Sn⁴⁺ and Pb²⁺ ions.

Heavier p-block metals

Some heavier p-block metals form ions with a charge two less than expected, because their outer s electrons are reluctant to be lost (the “inert pair effect”): thallium forms Tl⁺ as well as Tl³⁺; lead most often forms Pb²⁺ rather than Pb⁴⁺.

Noble gases

Group 18 elements already have full shells and don’t form simple ions in ordinary chemistry. See noble gases.

Common polyatomic ions

Many ions contain several atoms. Their charges must be learned:

Ion	Formula	Charge
hydroxide	OH⁻	−1
nitrate	NO₃⁻	−1
carbonate	CO₃²⁻	−2
sulfate	SO₄²⁻	−2
phosphate	PO₄³⁻	−3
ammonium	NH₄⁺	+1

When writing formulas with more than one polyatomic ion, use brackets: calcium hydroxide is Ca(OH)₂; ammonium sulfate is (NH₄)₂SO₄. See polyatomic ions list.

Common mistakes

• Using the group number as the charge for non-metals. Chlorine is in group 17, but its ion is Cl⁻, not Cl¹⁷⁻. Subtract 18 from the group number.
• Giving noble gases a charge. Group 18 elements don’t form simple ions.
• Forgetting that transition metals vary. “Iron chloride” is ambiguous; always state iron(II) or iron(III).
• Leaving a formula unsimplified. Ca₂O₂ should be written CaO; ionic formulas show the simplest ratio.
• Dropping brackets around polyatomic ions. Mg(OH)₂ means two hydroxide ions; MgOH₂ would mean something completely different.
• Writing charges on neutral compounds. The formula of a compound, such as NaCl, has no overall charge, because the charges balance.

Why group 14 and the middle of the table are different

The further an element is from a noble gas, the more electrons it must lose or gain to reach a noble gas configuration, and the more energy that costs. Losing four electrons from carbon, or gaining four, would require so much energy that no lattice energy could repay it. Carbon therefore shares electrons in covalent bonds instead. The same logic explains why simple ions become less common towards the middle of each period.

Practice questions

Predict the charge of the ion formed by each element:

1. Lithium
2. Barium
3. Phosphorus
4. Iodine
5. Selenium

Write formulas for:

6. potassium sulfide
7. aluminium chloride
8. calcium nitride
9. iron(III) oxide
10. ammonium carbonate

Answers: (1) Li⁺; (2) Ba²⁺; (3) P³⁻; (4) I⁻; (5) Se²⁻; (6) K₂S; (7) AlCl₃; (8) Ca₃N₂; (9) Fe₂O₃; (10) (NH₄)₂CO₃.

Key takeaways

• Group 1, 2 and 13 metals form +1, +2 and +3 ions; group 15, 16 and 17 non-metals form −3, −2 and −1 ions.
• Ions usually have the electron arrangement of the nearest noble gas.
• Charges in an ionic compound must balance; the swap-and-drop method gives the formula.
• Transition metals often form several ions; Roman numerals in names show the charge.
• Learn common polyatomic ions separately, and watch for special cases such as hydrogen, group 14 and heavier p-block metals.