Learn a reliable method for using the periodic table to predict common ion charges and write balanced ionic formulas. This guide includes a main-group ion-charge chart, worked examples, checks for transition metals and polyatomic ions, and a practical review routine for homework, laboratory calculations and exam revision.
Overview
Ionic compounds form when positively charged ions, called cations, combine with negatively charged ions, called anions. The final chemical formula must be electrically neutral: the total positive charge has to equal the total negative charge.
The periodic table helps you predict the charges of many main-group ions because elements in the same group often have the same number of valence electrons. Metals usually lose electrons and form positive ions, while non-metals usually gain electrons and form negative ions. This is a useful starting pattern, but it is not a complete rule for every element. Transition metals can have more than one charge, and some ions consist of several atoms joined together.
To write an ionic formula, use this sequence:
- Identify the cation and anion.
- Write the charge of each ion.
- Choose subscripts that make the total charge zero.
- Reduce the subscripts to the simplest whole-number ratio.
- Check brackets, charges and formula conventions.
Do not carry the charge symbols into the finished formula. Charges are used to work out the ratio, whereas subscripts show how many ions are present.
What to track
Main-group ion charges
The following reference table covers common charges used in introductory chemistry. Group labels can vary between periodic tables, so the element examples are included as a cross-check.
| Periodic-table group or family | Typical ion charge | Examples |
|---|---|---|
| Group 1 metals | +1 | Li+, Na+, K+ |
| Group 2 metals | +2 | Mg2+, Ca2+, Ba2+ |
| Group 13 metals | Usually +3 | Al3+ |
| Group 15 non-metals | Usually −3 | N3−, P3− |
| Group 16 non-metals | Usually −2 | O2−, S2− |
| Group 17 halogens | −1 | F−, Cl−, Br−, I− |
Group 18 elements are generally unreactive in basic ionic-formula exercises because their outer electron shells are already relatively stable. The charges in the table are common predictions, not permission to ignore the compound’s stated name or supplied data.
Common polyatomic ions
A polyatomic ion is a charged group of atoms that behaves as one unit. Memorise or look up the ions required by your course rather than trying to predict every charge from the periodic table alone. Common examples include hydroxide, OH−; nitrate, NO3−; sulfate, SO42−; carbonate, CO32−; and ammonium, NH4+.
When more than one polyatomic ion is needed, place the ion in brackets before adding the subscript. For example, calcium nitrate contains Ca2+ and NO3−. Two nitrate ions are needed to balance one calcium ion, so the formula is Ca(NO3)2, not CaNO32.
Worked formula examples
Sodium chloride: Sodium forms Na+ and chlorine forms Cl−. One of each balances the charges, giving NaCl.
Magnesium oxide: Magnesium is Mg2+ and oxide is O2−. The charges cancel in a one-to-one ratio, so the formula is MgO. Do not write Mg2O2; the ratio must be simplified.
Aluminium oxide: Aluminium is Al3+ and oxide is O2−. The lowest combination that balances both charges is two aluminium ions and three oxide ions: Al2O3.
Calcium hydroxide: Calcium is Ca2+ and hydroxide is OH−. Two hydroxide ions are needed, so the formula is Ca(OH)2.
Iron(III) chloride: The Roman numeral III states that iron has a +3 charge: Fe3+. Chloride is Cl−, so three chloride ions balance one iron ion. The formula is FeCl3. Iron(II) chloride would instead be FeCl2, which shows why transition-metal names must be read carefully.
Cadence and checkpoints
Ion charges do not change from week to week in the way a weather measurement or mission status might, so this reference is best revisited at useful learning checkpoints rather than on an arbitrary daily schedule.
- Before a new topic: Review the main-group pattern and the polyatomic ions listed by your course.
- During practice: Check every formula by adding its total positive and negative charge.
- After a set of questions: Record recurring errors, such as forgetting brackets or using the wrong transition-metal charge.
- Before an assessment: Practise unfamiliar combinations rather than memorising only familiar examples.
- After a curriculum or reference update: Confirm notation, group numbering and the required list of ions against your current specification or classroom reference.
A simple study tracker can include the compound name, cation, anion, charges, proposed formula, charge check and any correction. This separates a mistake in identifying an ion from a mistake in balancing the formula.
How to interpret changes
Most errors are not caused by a change in the periodic table. They usually come from changing the charge into a subscript incorrectly or overlooking information in the compound’s name.
Use the following checks:
- Check the ion identities. “Oxide” is O2−, while “oxygen” may refer to the element rather than a named ion. Likewise, chloride is Cl−, not neutral chlorine gas.
- Check charge balance. In Al2O3, the positive total is 2 × +3 = +6 and the negative total is 3 × −2 = −6.
- Check the simplest ratio. If all subscripts share a common factor, reduce them. Ionic formulas represent the simplest ratio of ions.
- Check brackets. Use brackets when a polyatomic ion appears more than once. One unchanged polyatomic ion does not need brackets: NaOH is correct.
- Check variable-charge metals. Roman numerals in names such as copper(II) or iron(III) specify the cation charge. Do not infer it solely from the metal’s position.
- Check the order. Write the cation first and the anion second, even though the anion may be written first in some descriptive explanations.
A formula that looks familiar can still be wrong. For example, CaCl is not neutral because one Ca2+ and one Cl− leave a net charge of +1. The balanced formula is CaCl2.
When to revisit
Return to this guide whenever you begin a new set of ionic-compound questions, encounter a transition metal or polyatomic ion, or notice that your charge checks are failing. A quarterly review is reasonable for a general reference page, while students may benefit from a short review at the start and end of each chemistry unit.
Editors and teachers should check the page when a curriculum specification, textbook convention or classroom ion list changes. The core charge-balance method is stable, but terminology, group-numbering conventions and the examples expected in assessments can differ between courses.
For your next practice question, write the ions separately, mark their charges, balance the total charge, simplify the subscripts and then perform one final inspection for brackets and Roman numerals. That five-step habit is more reliable than trying to memorise every formula as an isolated fact.
