Difficulty: Introductory | Prerequisites: Basic atomic structure (protons, neutrons, electrons), familiarity with the periodic table layout.
This is the foundational material for understanding how and why atoms form compounds. Ionic bonding sits at one end of the bonding spectrum and governs the chemistry of salts, minerals, and most inorganic compounds. If you can identify valence electrons and predict ionic charges from an element's position on the periodic table, you have the toolkit for writing ionic formulas and naming ionic compounds. This material is tested heavily in general chemistry exams and is prerequisite to understanding covalent bonding, electronegativity, and molecular geometry.
Atoms form ionic bonds by transferring electrons: metals lose electrons to become positively charged cations, nonmetals gain electrons to become negatively charged anions, and the opposite charges hold the compound together. You can predict an element's ionic charge from its group number on the periodic table, then use those charges to write balanced ionic formulas where the total positive charge equals the total negative charge.
Atomic number (Z)
The number of protons in the nucleus of an atom. It defines which element you are looking at. Think of it as the element's identity number.
Atomic mass (amu)
The weighted average mass of an element's naturally occurring isotopes, measured in atomic mass units. In simple terms, this tells you roughly how heavy one atom of that element is.
Valence electrons
The electrons in the outermost energy level (shell) of an atom. These are the electrons involved in bonding. For main-group elements, the group number tells you how many valence electrons the atom has (Group 1 = 1, Group 2 = 2, Group 13 = 3, and so on up to Group 18 = 8).
Octet rule
Atoms tend to gain, lose, or share electrons until they have eight electrons in their outermost shell, matching the electron configuration of a noble gas. Think of it as atoms wanting a "full outer shell" for stability.
Cation
A positively charged ion, formed when an atom loses one or more electrons. Metals form cations. In simple terms, fewer electrons than protons means a net positive charge.
Anion
A negatively charged ion, formed when an atom gains one or more electrons. Nonmetals form anions. More electrons than protons means a net negative charge.
Ionic bond
The electrostatic attraction between a cation and an anion. It forms when one atom transfers electrons to another, typically between a metal and a nonmetal.
Ionic compound
A compound held together by ionic bonds, consisting of a lattice of alternating cations and anions. Table salt (NaCl) is the classic example.
Electron transfer
The process by which one atom gives up electrons and another atom accepts them, producing ions. This is the mechanism behind ionic bond formation.
Electrostatic attraction
The force of attraction between opposite charges. This is what holds ionic compounds together in their crystal lattice structure.
Atomic number = number of protons = number of electrons in a neutral atom. Elements are arranged in order of increasing atomic number.
Atomic mass is not the same as atomic number. Atomic mass reflects the total mass of protons and neutrons (averaged across isotopes). Two elements can have similar atomic masses but very different atomic numbers (e.g. Ar = 39.94 amu at Z=18, K = 39.09 amu at Z=19).
Groups (columns) determine chemical behaviour. Elements in the same group have the same number of valence electrons.
Group 1 (alkali metals): 1 valence electron
Group 2 (alkaline earth metals): 2 valence electrons
Group 13: 3 valence electrons
Group 14: 4 valence electrons
Group 15: 5 valence electrons
Group 16 (chalcogens): 6 valence electrons
Group 17 (halogens): 7 valence electrons
Group 18 (noble gases): 8 valence electrons (full shell, generally unreactive)
Periods (rows) tell you how many electron shells the atom has.
Ionic bonds form between metals (which lose electrons) and nonmetals (which gain electrons).
The metal atom loses its valence electrons to achieve a noble gas configuration, becoming a cation.
The nonmetal atom gains enough electrons to fill its outer shell, becoming an anion.
The electrostatic attraction between the cation and anion holds the ionic compound together.
Group 1 metals lose 1 electron, forming 1+ ions (Na+, K+).
Group 2 metals lose 2 electrons, forming 2+ ions (Mg2+, Ca2+).
Group 13 metals lose 3 electrons, forming 3+ ions (Al3+).
Group 15 nonmetals gain 3 electrons, forming 3- ions (N3-).
Group 16 nonmetals gain 2 electrons, forming 2- ions (O2-, S2-).
Group 17 nonmetals gain 1 electron, forming 1- ions (F-, Cl-, Br-).
Determine the charge on each ion from the periodic table.
Find the ratio that makes total positive charge equal total negative charge.
Write the cation first, anion second, with subscripts showing the ratio.
The formula must be electrically neutral.
Worked examples from the lab:
Na+ and O2-: need 2 sodium ions to balance 1 oxide ion, so Na2O (sodium oxide).
K+ and F-: charges already balance 1:1, so KF (potassium fluoride).
Mg2+ and Cl-: need 2 chloride ions to balance 1 magnesium ion, so MgCl2 (magnesium chloride).
Ca2+ and N3-: need 3 calcium ions and 2 nitride ions (total charge: 6+ and 6-), so Ca3N2 (calcium nitride).
Al3+ and S2-: need 2 aluminium ions and 3 sulfide ions (total charge: 6+ and 6-), so Al2S3 (aluminium sulfide).
Name the metal (cation) first, unchanged.
Name the nonmetal (anion) second, changing the ending to "-ide."
Examples: NaCl = sodium chloride, CaO = calcium oxide, Al2S3 = aluminium sulfide.
For metals with variable charges (transition metals), include a Roman numeral: FeCl2 = iron(II) chloride, FeCl3 = iron(III) chloride.
Group | Valence Electrons | Typical Ionic Charge | Examples |
|---|---|---|---|
1 (Alkali metals) | 1 | 1+ | Na+, K+, Li+ |
2 (Alkaline earth) | 2 | 2+ | Mg2+, Ca2+, Sr2+ |
13 | 3 | 3+ | Al3+ |
15 | 5 | 3- | N3-, P3- |
16 (Chalcogens) | 6 | 2- | O2-, S2- |
17 (Halogens) | 7 | 1- | F-, Cl-, Br-, I- |
Cross the charges: if the cation is X+ and the anion is Y-, the subscript on the cation equals the anion's charge number, and the subscript on the anion equals the cation's charge number. Then simplify to the lowest whole-number ratio.
Example: Al3+ and O2- gives Al2O3 (cross the 3 and the 2).
Table salt (NaCl) is the most familiar ionic compound: sodium transfers one electron to chlorine, producing the crystal lattice you shake onto food. Calcium carbonate (CaCO3) is the ionic compound in limestone, marble, and chalk. The ionic compound calcium phosphate (Ca3(PO4)2) is the main mineral component of your bones and teeth. Understanding ionic bonding is also why engineers choose specific salts as electrolytes in batteries.
Students often confuse atomic number with atomic mass. Atomic number is the proton count and is always a whole number. Atomic mass is the weighted average of isotopes and is usually not a whole number. Two neighbouring elements can have very similar masses (Ar = 39.94, K = 39.09) despite being completely different elements.
Students sometimes assign the wrong charge to an ion by counting valence electrons as the charge. The ionic charge is how many electrons the atom gains or loses to reach a full octet, not the number of valence electrons itself. Chlorine has 7 valence electrons but forms a 1- ion (it gains one), not a 7- ion.
Writing the formula in the wrong order is a common slip. The cation (metal) always comes first in the formula. Strontium phosphide is Sr3P2, not P2Sr3.
Students sometimes forget to reduce formulas to the lowest ratio. If you cross charges and get Mg2O2, simplify to MgO.
Predicting ionic charges from periodic table position is almost guaranteed to appear on a general chemistry exam.
Writing and balancing ionic formulas (charge neutrality) is a core skill tested in multiple formats: fill-in-the-blank, matching, and free response.
Naming ionic compounds (metal name + nonmetal root + "-ide") is frequently tested, especially the distinction between main-group naming and transition-metal naming with Roman numerals.
Distinguishing ionic from covalent compounds based on metal/nonmetal classification is a common exam question (covered in detail in Part 2 of these notes).
True or False: An element's atomic number equals the number of neutrons in its nucleus.
Answer: False. Atomic number = number of protons.
Fill in the blank: A sodium atom (Group 1) forms an ion with a charge of ______.
Answer: 1+
True or False: In an ionic compound formula, the anion is written first.
Answer: False. The cation (metal) is always written first.
Fill in the blank: Calcium (Ca2+) and chlorine (Cl-) form the ionic compound with the formula ______.
Answer: CaCl2
True or False: Noble gases (Group 18) readily form ionic bonds.
Answer: False. Noble gases have full outer shells and are generally unreactive.
Q: How many valence electrons does sulphur (Group 16) have, and what ionic charge does it typically form?
A: Sulphur has 6 valence electrons. It gains 2 electrons to complete its octet, forming a 2- ion (S2-).
Q: Write the formula for aluminium oxide. Show your reasoning.
A: Aluminium forms Al3+ (Group 13, loses 3 electrons). Oxygen forms O2- (Group 16, gains 2 electrons). Cross the charges: Al2O3. Check: 2(3+) + 3(2-) = 6+ + 6- = 0. The formula is Al2O3.
Q: Potassium and nitrogen form an ionic compound. Predict the formula and name it.
A: K+ (Group 1) and N3- (Group 15). Three potassium ions balance one nitride ion: K3N, potassium nitride.
Q: Argon (Ar, Z=18) and potassium (K, Z=19) have very similar atomic masses (39.94 and 39.09 amu). Explain why their atomic masses do not follow the same order as their atomic numbers.
A: Atomic mass depends on the number of protons and neutrons, plus the relative abundance of each isotope. Potassium's most abundant isotopes happen to have fewer neutrons on average than argon's, so potassium is slightly lighter despite having one more proton.
Q: Strontium (Sr) is in Group 2 and phosphorus (P) is in Group 15. Predict the ionic compound formula.
A: Sr2+ and P3-. Cross charges to get Sr3P2, strontium phosphide.
This connects directly to covalent bonding (Part 2 of these notes), where atoms share electrons instead of transferring them. The dividing line between ionic and covalent depends on electronegativity difference, which is a periodic trend you will study next. Understanding ionic charges also feeds into stoichiometry, where you balance chemical equations, and into solution chemistry, where ionic compounds dissociate into their component ions when dissolved in water.
Ionic bonding, ionic compounds, valence electrons, cation, anion, electron transfer, octet rule, ionic formula, charge balancing, cross-multiplication method, naming ionic compounds, metal and nonmetal bonding, electrostatic attraction, crystal lattice, periodic table groups, main-group ions, alkali metals, alkaline earth metals, halogens, chalcogens, noble gas configuration, general chemistry lab 4, Purdue CHM, ionic charge prediction