Formation of Ionic Bonds and Properties of Ionic Compounds

Block A ended with the periodic trends that make this chapter possible: electronegativity, ionization energy and electron affinity all measure, in one way or another, how tightly an atom holds on to its electrons and how much it wants more. Chemical bonding is nothing more than the consequence of those trends when two different atoms meet — the redistribution of valence electrons that leaves both atoms, taken together, at lower potential energy than they had apart. This chapter opens the bonding block with the case where that redistribution is most extreme: complete electron transfer between a metal and a non-metal, producing an ionic bond.

To follow the argument you should have to hand the trends in ionization energy (low for metals on the left of the table, especially Groups 1 and 2), electron affinity (favourable — energy-releasing — for non-metals on the right, especially Groups 16 and 17), and electronegativity (the single number, on the Pauling scale, that combines both tendencies into one predictive tool). Nothing here depends on wave mechanics or orbital overlap; those belong to the covalent-bonding chapters that follow. An ionic bond is best understood first as electrostatics: point charges of opposite sign, attracting according to Coulomb's law, arranged into the most stable repeating pattern available.

1. Electron Transfer and the Formation of Ions

An ionic bond forms when one or more electrons transfer completely from a metal atom to a non-metal atom, so that both atoms become ions carrying opposite charges that then attract each other electrostatically. The transfer is energetically favourable whenever the metal has a low ionization energy (little energy is needed to remove its outer electron) and the non-metal has a strongly negative electron affinity (a large amount of energy is released when it gains an electron) — which is exactly why ionic bonding is concentrated between Groups 1–2 and Groups 16–17 of the periodic table, the two regions where electronegativity differs most sharply.

Sodium, with electron configuration $[\text{Ne}]3s^{1}$, loses its single outer electron to reach the stable, fully-filled configuration of neon:

Chlorine, with configuration $[\text{Ne}]3s^{2}3p^{5}$, gains that electron to complete its own octet, reaching the configuration of argon:

Both ions are now isoelectronic with the nearest noble gas — the driving force behind electron transfer is precisely the exceptional stability of a complete octet (or, for the lightest elements, a complete first shell of two). The octet rule states that atoms tend to gain, lose or share electrons until each has eight electrons in its outer shell; the next chapter extends this rule to covalent bonding, and the chapter after that catalogues its genuine exceptions.

Key distinction. Ionic bonding is not "sharing with a strong preference" — it is complete transfer. The metal atom's outer electron(s) leave entirely; they are not merely pulled closer to the non-metal. This is the qualitative end-point of the electronegativity-difference scale that the covalent chapters will use to classify bonds that fall short of it.

2. Formulas of Ionic Compounds: the Charge-Balance Rule

An ionic compound is electrically neutral overall, even though it is built entirely from charged particles. This single requirement — total positive charge equals total negative charge — fixes the ratio of cations to anions in the formula, and is applied formally through the criss-cross method: the numerical value of each ion's charge (ignoring its sign) becomes the subscript of the other ion.

Worked example. Determine the formula of magnesium chloride from $\text{Mg}^{2+}$ and $\text{Cl}^{-}$.

Criss-crossing the charge magnitudes: the $2$ from $\text{Mg}^{2+}$ becomes the subscript on chloride, and the $1$ from $\text{Cl}^{-}$ becomes the (unwritten) subscript on magnesium:

Checking: one $\text{Mg}^{2+}$ contributes $+2$; two $\text{Cl}^{-}$ contribute $2\times(-1)=-2$. The charges cancel, confirming the formula. Electron-transfer bookkeeping agrees: one magnesium atom must lose two electrons ($[\text{Ne}]3s^2 \to [\text{Ne}]$) and each electron is accepted by a separate chlorine atom, so two chlorine atoms are needed per magnesium.

Worked example. Determine the formula of aluminium oxide from $\text{Al}^{3+}$ and $\text{O}^{2-}$.

Criss-crossing $3$ and $2$ gives $\text{Al}_2\text{O}_3$ directly; the subscripts $2$ and $3$ do not share a common factor, so this is already the simplest ratio.

Checking: $2\times(+3) + 3\times(-2) = 6-6=0$. When a polyatomic ion such as $\text{SO}_4^{2-}$ or $\text{NO}_3^{-}$ requires more than one unit in a formula, the whole ion is enclosed in parentheses before the subscript is applied, e.g. $\text{Al}_2(\text{SO}_4)_3$ for aluminium sulfate — never $\text{AlSO}_{4\,3}$, which would misassign the subscript to a single oxygen atom instead of the whole sulfate unit.

Warning. Criss-crossed subscripts must always be reduced to the smallest whole-number ratio sharing no common factor. $\text{Ca}^{2+}$ and $\text{O}^{2-}$ criss-cross formally to $\text{Ca}_2\text{O}_2$, but the correct formula is $\text{CaO}$: both subscripts share the factor $2$, which must be cancelled.

3. The Ionic Lattice: a Three-Dimensional Array, Not a Molecule

A single $\text{Na}^{+}$ ion attracts every $\text{Cl}^{-}$ ion in its vicinity, not just one partner — electrostatic attraction has no direction and no saturation point. The result is a giant ionic lattice: a regular, repeating three-dimensional arrangement in which each ion is surrounded by several oppositely charged neighbours, called its coordination number. In the rock-salt structure adopted by $\text{NaCl}$, each $\text{Na}^{+}$ ion is surrounded by six $\text{Cl}^{-}$ ions (and vice versa), each at the same distance, arranged octahedrally.

This structural fact has an immediate consequence for formula writing: $\text{NaCl}$ is not a "molecule" of sodium chloride in the sense that $\text{H}_2\text{O}$ is a molecule of water. The formula is an empirical formula, stating only the simplest whole-number ratio of ions present in the extended lattice — there is no discrete $\text{NaCl}$ unit that can be isolated. This is why chemists write "formula unit," not "molecule," when discussing an ionic compound's mass or its moles.

The stability of the lattice is quantified by the lattice enthalpy, $\Delta H_{\text{lattice}}$: the enthalpy change when one mole of the solid ionic lattice forms from its constituent gaseous ions, $\text{M}^{+}(g) + \text{X}^{-}(g) \longrightarrow \text{MX}(s)$ — always strongly exothermic, since gaseous ions separated by large distances are converted into a tightly bound solid. Sign convention. This chapter (like many textbooks) quotes lattice enthalpy for this formation direction, so the values are negative. The IB data booklet and IB mark schemes define lattice enthalpy for the reverse, dissociation direction, $\text{MX}(s) \longrightarrow \text{M}^{+}(g) + \text{X}^{-}(g)$, so IB values are positive (e.g. $+790~\text{kJ mol}^{-1}$ for $\text{NaCl}$, $+3791~\text{kJ mol}^{-1}$ for $\text{MgO}$). Always write down the equation you are using; every comparison of magnitude below is unaffected. Two structural factors govern its magnitude, both following directly from Coulomb's law, $F \propto \dfrac{q_1q_2}{r^2}$:

• Ionic charge. A larger magnitude of charge on either ion increases the electrostatic attraction, and therefore the magnitude of the lattice enthalpy, roughly in proportion to the product of the charges.
• Ionic radius. Smaller ions can approach more closely, decreasing $r$ and increasing the attraction (and the magnitude of the lattice enthalpy) according to the inverse-square law.

Worked example. Rank $\text{NaCl}$, $\text{MgO}$ and $\text{KBr}$ in order of increasing magnitude of lattice enthalpy, using only ionic charge and radius.

$\text{MgO}$ is built from $\text{Mg}^{2+}$ and $\text{O}^{2-}$, a $2+/2-$ pair; $\text{NaCl}$ and $\text{KBr}$ are both $1+/1-$ pairs. The charge product for $\text{MgO}$ is $2\times 2=4$, four times larger than the $1\times1=1$ product for the other two, so $\text{MgO}$ has by far the largest-magnitude lattice enthalpy (experimentally, around $-3800~\dfrac{\text{kJ}}{\text{mol}}$, against roughly $-790~\dfrac{\text{kJ}}{\text{mol}}$ for $\text{NaCl}$). Between $\text{NaCl}$ and $\text{KBr}$, both ion pairs carry $1+/1-$ charges, so radius decides: $\text{Na}^{+}$ and $\text{Cl}^{-}$ are both smaller than $\text{K}^{+}$ and $\text{Br}^{-}$ respectively (same trend down a group established in Block A), so $\text{NaCl}$ has the larger-magnitude lattice enthalpy of the two. In increasing magnitude: $\text{KBr} < \text{NaCl} < \text{MgO}$.

4. Physical Properties of Ionic Compounds

The giant-lattice structure and the strength of the electrostatic forces holding it together explain, directly and without exception, the characteristic physical properties of ionic compounds.

• High melting and boiling points. Melting a lattice requires breaking the attraction between every ion and its many neighbours throughout the whole structure, which demands a large amount of thermal energy. The stronger the lattice (larger ionic charge, smaller ionic radius), the higher the melting point — $\text{MgO}$ (melting point $2852~^{\circ}\text{C}$) melts far above $\text{NaCl}$ ($801~^{\circ}\text{C}$), exactly tracking the lattice-enthalpy comparison above.
• Brittleness. A crystal of an ionic solid shatters rather than bends under stress. An applied force that shifts one layer of ions relative to its neighbour brings ions of like charge into direct alignment; the resulting strong repulsion fractures the lattice along that plane rather than allowing a smooth, ductile deformation.
• Electrical conductivity — solid versus molten/dissolved. In the solid state, every ion is locked in a fixed lattice position and cannot migrate, so solid ionic compounds do not conduct electricity. Once molten or dissolved in water, the lattice breaks apart and the ions become mobile charge carriers, free to migrate towards an oppositely charged electrode — so the molten or aqueous state conducts well. This solid/molten contrast is one of the most frequently tested diagnostic properties distinguishing ionic from covalent (molecular) substances.
• Solubility in polar solvents. Water molecules, being polar (a property developed fully in the chapter on molecular polarity), can surround and stabilize individual ions as they separate from the lattice — a process called hydration — releasing enough energy to offset much of the lattice enthalpy for many ionic compounds, which is why so many ionic solids dissolve readily in water but are insoluble in non-polar solvents such as hexane.

Key distinction. "Ionic compounds conduct electricity" is an incomplete and exam-penalized statement. The correct claim is conditional: they conduct only when molten or in aqueous solution, where ions are free to move; the solid lattice does not conduct, because no charge carrier is mobile.

5. Common Exam Pitfalls

• Writing "molecule of NaCl." An ionic compound has no discrete molecular unit; the correct term is formula unit, and the formula itself states only the simplest ion ratio in an extended lattice.
• Forgetting to reduce criss-crossed subscripts. $\text{Ca}^{2+}$ with $\text{O}^{2-}$ criss-crosses to $\text{Ca}_2\text{O}_2$ before the common factor of $2$ is cancelled to give the correct $\text{CaO}$.
• Omitting parentheses around a polyatomic ion that needs more than one unit. $\text{Mg(NO}_3)_2$, not $\text{MgNO}_{3\,2}$ or $\text{MgNO}_6$ — the parentheses preserve the identity of the whole nitrate ion.
• Claiming solid ionic compounds conduct electricity. They do not; mobility of ions, available only once molten or dissolved, is the requirement for conduction, and this must be stated explicitly to earn credit.
• Attributing lattice strength to "ionic bond strength" in the abstract, without naming charge and radius specifically. Exam mark schemes require both factors to be named and applied in the correct direction (charge dominates when it differs; radius decides only when charges are equal).
• Describing ionic bonding as "sharing electrons unequally." That description belongs to a polar covalent bond. An ionic bond is complete transfer, with each atom becoming a discrete ion — the two are different points on the electronegativity-difference scale, not the same phenomenon by degree.

Summary

• An ionic bond forms by complete electron transfer from a metal (low ionization energy) to a non-metal (strongly exothermic electron affinity), each ion reaching a full outer shell/octet, most often isoelectronic with the nearest noble gas.
• Ionic formulas are fixed by overall charge neutrality, applied via the criss-cross method and always reduced to the smallest whole-number ratio; polyatomic ions needing more than one unit are enclosed in parentheses.
• Ionic solids form a giant lattice of alternating ions with no discrete molecular unit, so their formula is an empirical, whole-ratio formula, and their mass is reckoned per formula unit.
• Lattice enthalpy, $\Delta H_{\text{lattice}}$ for $\text{M}^{+}(g)+\text{X}^{-}(g)\to \text{MX}(s)$, is always strongly exothermic in this (formation) direction — the IB data booklet quotes the reverse, endothermic, dissociation direction with positive values — and increases in magnitude with larger ionic charge and smaller ionic radius, by Coulomb's law.
• The physical properties of ionic compounds — high melting/boiling points, brittleness, no conduction as a solid but good conduction molten or dissolved, and solubility in polar solvents — all follow directly from the strength and three-dimensional, non-directional nature of the lattice.
• The next chapter turns to the opposite extreme of the electronegativity-difference scale: covalent bonding, where electrons are shared rather than transferred, developed through Lewis structures.