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Structure 2.1 · The ionic model

1Forming ions: cations and anions 2.1.1 SL + HL

Metal atoms have few valence electrons and low ionization energies. When a metal reacts with a non-metal, its valence electrons are transferred to the non-metal atoms. The metal atom becomes a cation (positive ion) and the non-metal atom becomes an anion (negative ion). For main-group elements the ions usually have the electron configuration of a noble gas: sodium loses its single 3s electron to become Na⁺ (1s²2s²2p⁶, the configuration of neon), and chlorine gains one electron into its 3p sub-level to become Cl⁻ (1s²2s²2p⁶3s²3p⁶, the configuration of argon).

Predicting the charge of an ion from the electron configuration

Count the valence electrons. An atom with one, two or three valence electrons (groups 1, 2 and 13) loses them to form ions of charge 1+, 2+ and 3+. An atom with five, six or seven valence electrons (groups 15, 16 and 17) gains three, two or one electron to form ions of charge 3−, 2− and 1−. Removing more electrons than this would mean breaking into a complete inner shell, which requires a very large ionization energy; adding more electrons would place them in a new, higher shell.

Transition elements form ions with different charges

Transition elements such as iron and copper can form more than one stable ion because several of their successive ionization energies are close in value. When a first-row transition element forms an ion, the 4s electrons are removed before the 3d electrons:

The Roman numeral in the name gives the charge on the metal ion, so the name alone tells you which ion is present. Transition-element ions generally do not have noble-gas configurations, which shows that the noble-gas “rule” is a useful pattern for main-group elements, not a law.

Sizes of ions

A cation is smaller than its parent atom: Na → Na⁺ removes the whole third shell, and the remaining ten electrons are held by the same eleven protons. An anion is larger than its parent atom: the added electrons increase electron–electron repulsion while the nuclear charge is unchanged. Within a group, ionic radius increases because each ion has an extra occupied shell; the sulfide ion (2,8,8) is larger than the oxide ion (2,8). For isoelectronic ions (same electron configuration), the ion with more protons is smaller: N³⁻ > O²⁻ > F⁻ > Na⁺ > Mg²⁺ > Al³⁺, all with the configuration 2,8. These size differences matter because they control the strength of ionic bonding (2.1.3).

2The ionic bond, formulas and names 2.1.2 SL + HL

The attraction acts in all directions, so each ion attracts every oppositely charged ion around it; the bond is non-directional. This is why ions do not pair off into molecules but build an extended three-dimensional lattice (2.1.3).

Deducing formulas

An ionic compound is electrically neutral: the total positive charge equals the total negative charge. Choose the smallest whole-number ratio of ions that balances the charges. When a polyatomic ion appears more than once, it is placed in brackets.

Inside a polyatomic ion the atoms are held by covalent bonds; the ion as a whole carries the charge. A compound such as sodium ethanoate, CH₃COONa, or ammonium nitrate, NH₄NO₃, therefore contains both ionic and covalent bonds.

Naming binary ionic compounds

Name the cation first, then the anion; a monatomic anion takes the suffix -ide (chloride, oxide, nitride, sulfide). Where the metal can form more than one ion, give its charge as a Roman numeral: FeCl₃ is iron(III) chloride.

3Ionic lattices and the properties of ionic compounds 2.1.3 SL + HL

In a solid ionic compound, each ion is surrounded by ions of opposite charge in a regular, repeating three-dimensional arrangement called a lattice. In sodium chloride, each Na⁺ ion has six Cl⁻ nearest neighbours and each Cl⁻ has six Na⁺ nearest neighbours (Figure 2.2). Attractions between neighbouring opposite charges outweigh repulsions between the more distant like charges, so the lattice is strongly held together.

Because there are no discrete molecules, the formula of an ionic compound is an empirical formula: it gives the simplest whole-number ratio of ions in the lattice, not a count of atoms in a molecule.

Volatility (melting and boiling points)

To melt an ionic solid the ions must gain enough energy to move out of their fixed lattice positions; to boil it, they must separate further. Because the electrostatic attractions extend throughout the lattice and are strong, large amounts of energy are needed: ionic compounds have high melting and boiling points and low volatility. Sodium chloride melts at 801 °C; magnesium oxide, with doubly charged ions, melts at 2800 °C.

Electrical conductivity

A current requires mobile charged particles. In the solid, ions are held in fixed positions, so solid ionic compounds do not conduct. When molten or dissolved in water the ions are free to move, cations towards the negative electrode and anions towards the positive electrode, so the liquid conducts. The charge carriers are ions, not electrons.

Solubility

Many ionic compounds dissolve in water. Water molecules are polar: the partially negative oxygen atoms are attracted to cations and the partially positive hydrogen atoms to anions. If these ion–water attractions are strong enough to compensate for the energy needed to separate the ions from the lattice, the compound dissolves. Some ionic compounds with very strong lattices (for example calcium carbonate and zinc oxide) are insoluble. Ionic compounds do not dissolve in non-polar solvents such as hexane, because hexane molecules cannot interact strongly with ions.

Brittleness

Ionic crystals are hard but brittle. A sharp blow can displace one layer of ions by one ion position. Ions of the same charge then face each other, the repulsion splits the crystal along a plane (Figure 2.3).

4Lattice enthalpy: a measure of ionic bond strength 2.1.3 SL + HL

The attraction between two ions depends on the product of their charges and on the distance between their centres (the sum of the ionic radii). Qualitatively:

attraction ∝ (charge on cation × charge on anion) / (distance between ion centres)

• Higher ionic charges give stronger attraction and a larger lattice enthalpy. This is the larger effect: compare NaCl (788 kJ mol⁻¹) with MgO (3890 kJ mol⁻¹), where both ions carry double charges.
• Smaller ions allow their centres to approach more closely, giving stronger attraction. Down group 17 the anion becomes larger, so lattice enthalpy falls: LiF > LiCl > LiBr > LiI.

Review · Structure 2.1

5Misconceptions and the examiner’s view

6Past-paper practice

7Summary and knowledge organiser

Essential knowledge

• Every bonding model describes an electrostatic attraction: between oppositely charged ions (ionic), between shared electron pairs and two nuclei (covalent), or between a lattice of cations and delocalized electrons (metallic).
• Ionic compounds form giant lattices; their properties (high melting point, brittleness, conductivity only when molten or in solution, solubility in polar solvents) follow from strong, non-directional attraction between ions that can move only when the lattice breaks down.

Essential definitions and relationships

Essential observations and properties

Examination checklist

• Describe each bond type using “electrostatic attraction” and name exactly which particles are attracted.
• Write ionic formulae with brackets around polyatomic ions; give full electron configurations when “full” is asked.
• Ionic compounds conduct because ions move, and only when they are free to move (molten or aqueous) — never because electrons move.

Knowledge organiser · the ionic model