General Chemistry I · Structure and Bonding

Ionic and Covalent Bonding

On this page 8 sections

In 30 seconds
Why this matters
The college version
Eli explains
Worked example
Key takeaway
Quick check
Study tools

In 30 seconds

Chemical bonds form when atoms transfer or share electrons to reach stable, lower-energy arrangements. Ionic bonds result from a full transfer of electrons between a metal and a nonmetal and are held together by ; covalent bonds share electrons, and the sharing is unequal (polar) when the atoms differ in electronegativity. Bond polarity creates dipole moments, with a partial negative charge (\(\delta^-\)) on the more electronegative atom and a partial positive charge (\(\delta^+\)) on the other.

Why this matters

Bond polarity is why water is a powerful solvent for ions and polar molecules, which underpins nearly all of biochemistry — from dissolving salts and sugars in blood to the folding of proteins via hydrogen bonding. Understanding lattice energy also explains why some ionic drugs dissolve readily (lower lattice energy, weaker crystal) while others resist dissolving, a key consideration in pharmaceutical formulation.

The college version

1. Ionic Bonding and Lattice Energy

An is the electrostatic attraction between oppositely charged ions. Lattice energy is the energy released when gaseous ions come together to form one mole of a solid ionic compound, or equivalently the energy required to separate one mole of the solid into gaseous ions. Coulomb's law shows why it is large for small, highly charged ions:

\[ E \propto \frac{Q_1 Q_2}{r} \]

where \(Q_1\) and \(Q_2\) are the ion charges and \(r\) is the distance between their centers. Compounds of small, multiply charged ions (e.g., \(MgO\)) have much larger lattice energies than those of large, singly charged ions (e.g., \(NaCl\)).

2. The Born–Haber Cycle

The is a thermochemical bookkeeping method that builds an ionic compound from its elements in hypothetical steps, each with a known enthalpy, so that lattice energy can be computed from the other steps. Because enthalpy is a state function, the sum of the step enthalpies around the cycle must equal the enthalpy of formation.

3. Electronegativity and Bond Polarity

Electronegativity (EN) is an atom's ability to attract shared electrons. The difference in electronegativity (\(\Delta EN\)) predicts bond type (approximate cutoffs):

\(\Delta EN = 0\) → nonpolar covalent (equal sharing).
\(0 < \Delta EN < \sim 0.4\) → essentially nonpolar covalent.
\(\sim 0.4 \le \Delta EN < \sim 2.0\) → polar covalent (unequal sharing).
\(\Delta EN \ge \sim 2.0\) → ionic (effectively transferred).

A dipole moment (\(\mu\)) measures the separation of charge in a polar bond, with magnitude \(\mu = Q \times r\), where \(Q\) is the magnitude of the partial charges and \(r\) is the distance between them. Diatomic molecules with polar bonds are always polar; in larger molecules, symmetry can cancel individual dipoles (see topic 22).

How it works

Identify the atoms and their electronegativities.
Compute \(\Delta EN\); classify the bond as nonpolar covalent, polar covalent, or ionic.
For polar bonds, mark the partial charges (\(\delta^-\) on the more electronegative atom).
For ionic compounds, use a Born–Haber cycle to find lattice energy from known enthalpies.
Predict relative lattice energies with Coulomb's law: smaller ions and higher charges → larger lattice energy.
Add up bond dipoles (topic 22) to decide if a whole molecule is polar.

Common confusions

Do not confuse	With	Difference
Lattice energy	Bond energy	Lattice energy applies to the whole ionic crystal; bond energy applies to a single covalent bond
Electronegativity	Electron affinity	EN is a relative scale for atoms in a bond; electron affinity is an energy measured for an isolated atom gaining an electron
Polar covalent	Ionic	Polar covalent shares electrons unequally; ionic transfers them fully
Dipole moment	Partial charge	The dipole moment is the product of charge and distance; \(\delta\) is the charge itself

Memory aids

"LARGER charge, SMALLER ions → LARGER lattice energy." For polarity, remember "ΔEN decides: 0 = even split, 2+ = full hand-off."

Quick review

Topic Recap

Bonding spans a spectrum from pure electron transfer (ionic) to equal sharing (nonpolar covalent), with polar covalent bonds in between. Electronegativity differences predict where a bond falls and produce partial charges and dipole moments. For ionic solids, the Born–Haber cycle ties measurable enthalpies together to reveal lattice energy, the quantity that makes ionic compounds stable.

Knowledge Check

What does lattice energy measure, and how does ionic charge affect it?
Which has the larger lattice energy, NaCl or MgO? Why?
Using electronegativities (\(H = 2.2\), \(O = 3.5\), \(F = 4.0\)), classify the O–H and F–F bonds.
In HF, which atom carries the partial negative charge, and why?
List the five steps of the Born–Haber cycle for an alkali-metal chloride.

Answers and Rationales

Lattice energy is the energy to separate one mole of an ionic solid into gaseous ions (or released when the lattice forms). It increases with the product of the ion charges (\(Q_1Q_2\)) — higher charges give much stronger attraction.
MgO. Both ions are doubly charged (\(Mg^{2+}, O^{2-}\)) versus singly charged (\(Na^+, Cl^-\)), and the ions are smaller, so \(E \propto Q_1Q_2/r\) is far larger.
O–H: \(\Delta EN = 3.5 - 2.2 = 1.3\) → polar covalent. F–F: \(\Delta EN = 0\) → nonpolar covalent.
Fluorine, because it is more electronegative (4.0 vs 2.2), so it pulls the shared electron pair toward itself.
(1) Sublimate the metal; (2) ionize the metal atom; (3) dissociate the halogen molecule; (4) add an electron to the halogen atom; (5) form the ionic lattice. Summing these gives the enthalpy of formation.

Eli explains

The same idea, in plain words

Explain it like I’m 10

Think of bonding like sharing a meal. In an ionic bond, one person simply hands over their food (electron) to the other — a clean transfer. In a nonpolar covalent bond, two equally hungry friends share the food evenly. In a polar covalent bond, the two friends still share, but one is much hungrier and pulls the plate closer to their side.

Where it stops being exact: real bonds are a continuum, not three boxes. Even "ionic" bonds have some shared-electron character, and "nonpolar" is a limit rarely reached exactly. The labels are useful approximations of where a bond falls on the transfer-to-share spectrum.

Simple Example

Sodium chloride is ionic: sodium gives its one valence electron to chlorine, forming \(Na^+\) and \(Cl^-\). Hydrogen gas (\(H_2\)) is nonpolar covalent: the two identical atoms share equally. Hydrogen chloride (\(HCl\)) is polar covalent: chlorine, being more electronegative, pulls the shared pair toward itself, giving H a \(\delta^+\) and Cl a \(\delta^-\).

Worked example

\[ \mu = Q \times r \]

where \(\mu\) is the dipole moment, \(Q\) the partial charge, and \(r\) the bond length.

Worked Example 1 — Born–Haber cycle for NaCl. Formation of \(NaCl(s)\) from \(Na(s)\) and \(\tfrac{1}{2}Cl_2(g)\) can be broken into steps (all values approximate, kJ/mol):

Sublime \(Na(s) \to Na(g)\): \(+108\)
Ionize \(Na(g) \to Na^+(g) + e^-\): \(+496\)
Dissociate \(\tfrac{1}{2}Cl_2(g) \to Cl(g)\): \(+122\)
Add electron \(Cl(g) + e^- \to Cl^-(g)\): \(-349\)
Form lattice \(Na^+(g) + Cl^-(g) \to NaCl(s)\): lattice energy \(U\) (unknown)

The sum of steps 1–5 equals the measured enthalpy of formation \(\Delta H_f^\circ = -411\text{ kJ/mol}\):

\[ 108 + 496 + 122 + (-349) + U = -411 \]

\[ U = -411 - (108 + 496 + 122 - 349) = -411 - 377 = -788\text{ kJ/mol} \]

So the lattice energy of NaCl is about −788 kJ/mol (released when the lattice forms). The large negative value is what makes ionic compounds so stable.

Worked Example 2 — Predicting bond polarity. For \(HCl\), \(EN(H) = 2.2\) and \(EN(Cl) = 3.2\), so \(\Delta EN = 1.0\) → polar covalent. For \(Cl_2\), \(\Delta EN = 0\) → nonpolar covalent. For \(NaCl\), \(\Delta EN = 3.2 - 0.9 = 2.3\) → ionic.

Common setup error: in the Born–Haber cycle, forgetting to divide the dissociation enthalpy of \(Cl_2\) by two (only half a mole of \(Cl_2\) is needed per mole of NaCl), or dropping the negative signs on electron affinity and lattice formation.

Key takeaways

High yield: Lattice energy grows with charge and shrinks with ionic size (\(E \propto Q_1Q_2/r\)).
High yield: The Born–Haber cycle sums step enthalpies to give lattice energy because enthalpy is a state function.
High yield: \(\Delta EN \approx 0.4\)–\(2.0\) → polar covalent; \(\ge 2.0\) → ionic; near 0 → nonpolar covalent.
Polar covalent bonds carry \(\delta^+\) and \(\delta^-\) partial charges; the more electronegative atom is \(\delta^-\).
A bond dipole is a vector; in symmetric molecules individual dipoles can cancel.
Ionic and covalent are endpoints of a continuum, not rigid categories.