Difficulty: Introductory | Prerequisites: Basic atomic structure, periodic table familiarity
This topic sits right at the start of chemical bonding. Before you can make sense of molecules, reactions, or anything in organic chemistry later, you need to understand how and why atoms form ions, and how those ions combine into compounds. The core idea: atoms gain or lose electrons to reach a stable, noble-gas electron configuration, and the charges that result determine the formulas of the compounds they form. If you are comfortable with electron configurations and the layout of the periodic table, you have everything you need here.
Atoms form ions by losing or gaining electrons until they have the same electron configuration as the nearest noble gas. Metals lose electrons (forming cations), nonmetals gain them (forming anions). Ionic compound formulas are set by balancing the total positive and negative charges so they cancel to zero.
Ion
An atom or group of atoms that carries a net electrical charge because it has gained or lost one or more electrons.
In simple terms, it is an atom that is no longer electrically neutral.
Cation
A positively charged ion, formed when an atom loses one or more electrons.
Think of it as: the atom gave away electrons, so it has more protons than electrons and ends up positive.
Anion
A negatively charged ion, formed when an atom gains one or more electrons.
In simple terms, the atom picked up extra electrons, so the negative charges now outnumber the positive.
Electron configuration
The notation that describes how electrons are distributed among the energy levels and subshells of an atom (e.g. 1s² 2s² 2p⁶ 3s¹ for sodium).
Think of it as: the address system for every electron in the atom.
Electron-dot symbol (Lewis dot symbol)
A shorthand showing the element's symbol surrounded by dots representing its valence electrons.
In simple terms, it is a quick sketch that tells you how many outer electrons an atom has available for bonding.
Ionic charge
The charge an ion carries, written as a superscript number followed by + or − (e.g. Na¹⁺, Cl¹⁻).
This is determined by how many electrons the atom lost or gained.
Octet rule
Atoms tend to gain, lose, or share electrons until they have eight electrons in their outermost energy level (matching a noble-gas configuration).
Think of it as: atoms want a "full outer shell," which for most means eight electrons.
Ionic compound
A compound formed by the electrostatic attraction between cations and anions. The formula is always written so total positive charge equals total negative charge.
In simple terms, a metal gives electrons to a nonmetal, and the opposite charges hold them together.
Metals sit on the left side of the periodic table and have few valence electrons (1, 2, or 3).
They lose those valence electrons to reach the electron configuration of the nearest noble gas.
Example: Sodium (Na), atomic number 11, has the configuration 1s² 2s² 2p⁶ 3s¹. It loses 1 electron to become Na¹⁺ with configuration 1s² 2s² 2p⁶ (same as neon).
Example: Aluminum (Al), atomic number 13, configuration 1s² 2s² 2p⁶ 3s² 3p¹. It loses 3 electrons to become Al³⁺ with configuration 1s² 2s² 2p⁶.
Example: Calcium (Ca), atomic number 20, configuration 1s² 2s² 2p⁶ 3s² 3p⁶ 4s². It loses 2 electrons to become Ca²⁺ with configuration 1s² 2s² 2p⁶ 3s² 3p⁶ (same as argon).
Nonmetals sit on the right side of the periodic table and need only a few electrons to complete their outer shell.
They gain electrons to reach a noble-gas configuration.
Example: Chlorine (Cl), atomic number 17, configuration 1s² 2s² 2p⁶ 3s² 3p⁵. It gains 1 electron to become Cl¹⁻ with configuration 1s² 2s² 2p⁶ 3s² 3p⁶ (same as argon).
Example: Nitrogen (N), atomic number 7, configuration 1s² 2s² 2p³. It gains 3 electrons to become N³⁻ with configuration 1s² 2s² 2p⁶ (same as neon).
Example: Oxygen (O), atomic number 8, configuration 1s² 2s² 2p⁴. It gains 2 electrons to become O²⁻ with configuration 1s² 2s² 2p⁶.
Group 1A metals form 1+ ions.
Group 2A metals form 2+ ions.
Group 3A metals (like Al) form 3+ ions.
Group 7A nonmetals (halogens) form 1− ions.
Group 6A nonmetals form 2− ions.
Group 5A nonmetals form 3− ions.
The pattern: metals lose enough electrons to empty their outer shell; nonmetals gain enough to fill theirs. Both end up with eight electrons in the outermost energy level (the octet rule).
The total positive charge must equal the total negative charge. Use the "criss-cross" method: take the magnitude of each ion's charge and use it as the subscript for the other ion.
Lithium iodide: Li¹⁺ and I¹⁻ combine as LiI (charges already balance one-to-one).
Aluminum oxide: Al³⁺ and O²⁻ criss-cross to give Al₂O₃ (2 × 3+ = 6+, 3 × 2− = 6−).
Calcium sulfide: Ca²⁺ and S²⁻ give CaS (charges balance one-to-one because both are magnitude 2).
Magnesium bromide: Mg²⁺ and Br¹⁻ give MgBr₂.
Potassium nitride: K¹⁺ and N³⁻ give K₃N.
Sodium fluoride: Na¹⁺ and F¹⁻ give NaF.
Write the metal name first, unchanged.
Write the nonmetal name second, changing its ending to "-ide."
Examples: K₂S = potassium sulfide; BaF₂ = barium fluoride; MgO = magnesium oxide; Na₃N = sodium nitride; AlCl₃ = aluminum chloride; Mg₃P₂ = magnesium phosphide.
Sodium chloride (NaCl) is a representative ionic compound: white, solid, crystalline; density 2.16 g/cm³; melting point 801 °C.
Ionic compounds as a class tend to be hard, brittle solids with high melting points. They conduct electricity when dissolved in water or melted (because the ions become free to move), but not in their solid state.
Table salt (NaCl) is the most familiar ionic compound. The high melting point of ionic solids is why ceramics (ionic and covalent network solids) are used as heat shields and kiln linings, and why salt must be dissolved before it conducts electricity in applications like water purification.
Students often think ions are formed by gaining or losing protons. They are not. The number of protons never changes; only electrons are transferred.
Students sometimes write the ionic charge as a subscript (Na₁₊) rather than a superscript (Na¹⁺). The charge is always a superscript.
When writing formulas, students sometimes forget to reduce subscripts. If the criss-cross gives Ca₂O₂, simplify to CaO (divide both subscripts by 2).
Students often confuse the ion's charge with the subscript in the formula. The subscript tells you how many of that ion are present, not the charge itself.
⚠️ Be able to write the electron configuration of any element in the first 20 and then show the configuration of its ion after gaining or losing electrons.
⚠️ Know that the ion configuration matches the nearest noble gas.
⚠️ Given two ions, write the correct formula by balancing charges (criss-cross method).
⚠️ Given a formula, name the compound (metal name + nonmetal root + "-ide").
⚠️ Know the physical properties of a typical ionic compound (high melting point, crystalline solid, conducts when dissolved or molten).
True or False: A sodium atom loses one electron to form a Na¹⁻ ion. (False, it forms Na¹⁺.)
Fill in the blank: Oxygen gains ___ electrons to form an oxide ion with a ___ charge. (2; 2−)
True or False: The formula for aluminum oxide is AlO. (False, it is Al₂O₃.)
True or False: Ionic compounds typically have low melting points. (False, they have high melting points.)
Fill in the blank: The electron configuration of the Cl¹⁻ ion is ___. (1s² 2s² 2p⁶ 3s² 3p⁶)
Q: Write the electron configuration of a calcium ion (Ca²⁺). Which noble gas does it match?
A: 1s² 2s² 2p⁶ 3s² 3p⁶. This matches argon (Ar).
Q: What is the formula for magnesium bromide? Show your working.
A: Mg²⁺ and Br¹⁻. Criss-cross: Mg gets subscript 1, Br gets subscript 2. Formula: MgBr₂.
Q: Name the compound K₂S.
A: Potassium sulfide. (Potassium is the metal; sulfur becomes sulfide.)
Q: An atom has 13 protons and 10 electrons. What is the ion, and what is its charge?
A: It is Al³⁺ (aluminum ion). It lost 3 electrons (13 − 10 = 3), giving a 3+ charge.
Q: Why do ionic compounds conduct electricity when dissolved in water but not as solids?
A: In the solid state, ions are locked in a fixed crystal lattice and cannot move. When dissolved, the ions separate and become free to carry charge through the solution.
This connects directly to the next sections on variable-charge metals and polyatomic ions, which add complexity to the same naming and formula-writing rules. Understanding electron transfer here is also the foundation for electrochemistry (batteries, corrosion) later in the course. The distinction between ionic and covalent bonding, covered in the molecular compounds section, starts from the same question: how do atoms reach a stable electron configuration?
Ions, cations, anions, electron transfer, electron configuration, Lewis dot structure, electron-dot symbol, ionic bond, ionic compound, octet rule, noble-gas configuration, valence electrons, ion formation, NaCl, sodium chloride, charge balancing, criss-cross method, naming ionic compounds, "-ide" suffix, metal and nonmetal bonding, Chemistry Lab 9