Ionic Bonding

Nature of Ionic Bonds

Ionic bonding is a type of chemical bonding that arises from the electrostatic attraction between oppositely charged ions. These ions are formed when atoms transfer one or more electrons from one to another, resulting in:

• Positively charged ions (cations), typically formed from metals.
• Negatively charged ions (anions), typically formed from nonmetals or polyatomic groups.

The ionic bond itself is not a shared pair of electrons between two specific atoms (as in covalent bonding), but a widespread attraction between all cations and anions in the solid.

Formation of Ions by Electron Transfer

In ionic bonding, electron transfer is driven largely by the tendency of atoms to reach a particularly stable electron configuration, often similar to that of the nearest noble gas.

• Metals (e.g. Na, Mg, Ca) generally lose electrons:
• Sodium: $ \mathrm{Na \rightarrow Na^+ + e^-} $
• Magnesium: $ \mathrm{Mg \rightarrow Mg^{2+} + 2\,e^-} $
• Nonmetals (e.g. Cl, O, S) generally gain electrons:
• Chlorine: $ \mathrm{Cl + e^- \rightarrow Cl^-} $
• Oxygen: $ \mathrm{O + 2\,e^- \rightarrow O^{2-}} $

Electron transfer between atoms leads to ionic compounds. For example:

• Sodium and chlorine:
• $ \mathrm{Na \rightarrow Na^+ + e^-} $
• $ \mathrm{Cl + e^- \rightarrow Cl^-} $
• Overall: $ \mathrm{Na + Cl \rightarrow Na^+ + Cl^-} $

The resulting $ \mathrm{Na^+} $ and $ \mathrm{Cl^-} $ attract each other electrostatically and become neighbors in the ionic lattice of sodium chloride.

Typical Combinations that Form Ionic Compounds

• Metal + nonmetal (e.g. $ \mathrm{NaCl} $, $ \mathrm{CaO} $, $ \mathrm{MgCl_2} $)
• Metal + polyatomic anion (e.g. $ \mathrm{Na_2SO_4} $, $ \mathrm{CaCO_3} $)
• Polyatomic cation + nonmetal or polyatomic anion (e.g. $ \mathrm{NH_4Cl} $, $ \mathrm{(NH_4)_2SO_4} $)

The distinction between metal (cation former) and nonmetal (anion former) is central to predicting where ionic bonding is likely.

Electrostatic Interaction and Coulomb’s Law

The force holding ions together is the Coulomb attraction between charges:

$$
F = k \frac{|q_1 q_2|}{r^2}
$$

where:

• $ F $ is the magnitude of the electrostatic force,
• $ q_1 $ and $ q_2 $ are the charges on the ions,
• $ r $ is the distance between their centers,
• $ k $ is Coulomb’s constant.

Key consequences:

• Greater charges (e.g. $ \mathrm{2+} $ and $ \mathrm{2-} $) give stronger attraction than $ \mathrm{1+} $ and $ \mathrm{1-} $ at the same distance.
• Smaller ionic radii (smaller $ r $) give stronger attraction.

This explains why compounds like $ \mathrm{MgO} $ (with $ \mathrm{Mg^{2+}} $ and $ \mathrm{O^{2-}} $) usually have higher melting points than compounds with singly charged ions (e.g. $ \mathrm{NaCl} $).

Ionic Lattices and Crystal Structures

In the solid state, ionic compounds do not exist as discrete “molecules.” Instead, they form extended three-dimensional crystal lattices in which each cation is surrounded by anions and each anion is surrounded by cations.

Lattice Characteristics

• Regular, repeating arrangement of ions.
• Overall electrical neutrality: total positive charge equals total negative charge.
• Coordination number: number of nearest, oppositely charged neighbors an ion has in the lattice.

Examples:

• Sodium chloride ($ \mathrm{NaCl} $):
• Structure type: rock salt.
• Each $ \mathrm{Na^+} $ is surrounded by 6 $ \mathrm{Cl^-} $ ions.
• Each $ \mathrm{Cl^-} $ is surrounded by 6 $ \mathrm{Na^+} $ ions.
• Coordination numbers: 6:6.
• Cesium chloride ($ \mathrm{CsCl} $):
• Structure type: cesium chloride.
• Each $ \mathrm{Cs^+} $ is surrounded by 8 $ \mathrm{Cl^-} $ ions at the corners of a cube.
• Each $ \mathrm{Cl^-} $ is similarly surrounded by 8 $ \mathrm{Cs^+} $ ions.
• Coordination numbers: 8:8.
• Calcium fluoride ($ \mathrm{CaF_2} $):
• Structure type: fluorite.
• Each $ \mathrm{Ca^{2+}} $ is surrounded by 8 $ \mathrm{F^-} $ ions.
• Each $ \mathrm{F^-} $ is surrounded by 4 $ \mathrm{Ca^{2+}} $ ions.
• Coordination: 8:4.

The actual geometry depends mainly on the sizes (radii) of the cations and anions and their charges.

Formula Units

Because ionic solids are extended lattices, formulas represent the simplest whole-number ratio of ions, not discrete molecules. For example:

• $ \mathrm{NaCl} $ represents a 1:1 ratio of sodium ions to chloride ions.
• $ \mathrm{CaCl_2} $ represents 1 $ \mathrm{Ca^{2+}} $ to 2 $ \mathrm{Cl^-} $.

These are called formula units, not molecules.

Lattice Energy

Lattice energy is a measure of the strength of the ionic bonding in a crystal lattice.

Definition: Lattice energy ($U_\mathrm{lat}$) is the energy released when one mole of an ionic solid is formed from its gaseous ions:

$$
\mathrm{M^+(g) + X^-(g) \rightarrow MX(s)} \quad \Delta H = -U_\mathrm{lat}
$$

A large (more negative) lattice energy corresponds to a stronger ionic bond in the crystal.

Factors Affecting Lattice Energy

1. Ionic charges
• Higher charges generally increase lattice energy:
• $U_\mathrm{lat}(\mathrm{MgO}) > U_\mathrm{lat}(\mathrm{NaCl})$
• Compounds with doubly or triply charged ions usually have higher melting points and are often less soluble.
2. Ionic radii
• Smaller ions can get closer together (smaller $ r $) and therefore attract more strongly.
• Lattice energy increases as the sum of ionic radii decreases.

Qualitatively, this is often summarized with a dependence similar to Coulomb’s law:

$$
U_\mathrm{lat} \propto \frac{z^+ z^-}{r^+ + r^-}
$$

where $ z^+ $ and $ z^- $ are the ionic charges and $ r^+ $, $ r^- $ the ionic radii.

Properties of Ionic Compounds

The nature of ionic bonding and the lattice structure leads to characteristic macroscopic properties.

Melting and Boiling Points

• Ionic substances generally have high melting and boiling points.
• A large amount of energy is required to disrupt the extensive electrostatic network in the lattice.
• Higher lattice energies correlate with higher melting and boiling points.

Examples:

• $ \mathrm{NaCl} $ melts at about $801^\circ\mathrm{C}$.
• $ \mathrm{MgO} $ melts at a much higher temperature (over $2800^\circ\mathrm{C}$) due to its larger lattice energy.

Hardness and Brittleness

• Ionic crystals are usually hard because strong forces bind the ions in place.
• They are also brittle: when a force shifts layers of ions so that like charges (e.g. $ \mathrm{Na^+} $ next to $ \mathrm{Na^+} $) align, strong repulsion occurs and the crystal cleaves.

Electrical Conductivity

The electrical conductivity of ionic substances depends on the mobility of ions:

• Solid ionic compounds:
• Ions are fixed in place in the lattice.
• They do not conduct electricity in the solid state (they are electrical insulators).
• Molten (liquid) ionic compounds:
• Heating to the melting point frees ions to move.
• The melt (e.g. molten $ \mathrm{NaCl} $) conducts electrical current.
• Aqueous solutions of ionic compounds:
• When an ionic compound dissolves in water, its ions separate and become mobile.
• The solution conducts electricity; the dissolved ions are called electrolytes.

Solubility in Water and Other Solvents

Ionic compounds are often soluble in polar solvents like water, but solubility varies widely.

• Dissolution process in water:
• Water molecules, having partial charges, surround ions (hydration).
• Hydration stabilizes the separated ions in solution.
• The balance between lattice energy and hydration energy determines whether an ionic compound is soluble.
• Compounds such as $ \mathrm{NaCl} $ or $ \mathrm{KNO_3} $ are quite soluble.
• Others, such as $ \mathrm{AgCl} $ or $ \mathrm{BaSO_4} $, are only sparingly soluble or practically insoluble, often due to high lattice energies or unfavorable hydration.

Nonpolar solvents (like hexane) generally cannot stabilize ions and therefore usually do not dissolve ionic compounds.

Writing and Interpreting Ionic Formulas

Ionic formulas are constructed to ensure overall electrical neutrality.

Charge Balance

The sum of positive charges must equal the sum of negative charges.

Process:

1. Identify the likely ion formed by each element (including charge).
2. Find the smallest whole-number ratio that makes the total charge zero.
3. Write the formula without charges, using subscripts to indicate how many of each ion.

Examples:

• Sodium oxide from $ \mathrm{Na^+} $ and $ \mathrm{O^{2-}} $:
• Need two $ \mathrm{Na^+} $ to balance one $ \mathrm{O^{2-}} $.
• Formula: $ \mathrm{Na_2O} $.
• Calcium chloride from $ \mathrm{Ca^{2+}} $ and $ \mathrm{Cl^-} $:
• Need two $ \mathrm{Cl^-} $ per $ \mathrm{Ca^{2+}} $.
• Formula: $ \mathrm{CaCl_2} $.
• Aluminum oxide from $ \mathrm{Al^{3+}} $ and $ \mathrm{O^{2-}} $:
• Find smallest integers $x$ and $y$ such that $3x = 2y$.
• $x = 2$, $y = 3$.
• Formula: $ \mathrm{Al_2O_3} $.

Polyatomic Ions

Many ionic compounds contain polyatomic ions (groups of atoms with a net charge). The same charge-balance principle applies.

Examples:

• $ \mathrm{Na_2SO_4} $: 2 $ \mathrm{Na^+} $ and 1 $ \mathrm{SO_4^{2-}} $.
• $ \mathrm{Ca(NO_3)_2} $: 1 $ \mathrm{Ca^{2+}} $ and 2 $ \mathrm{NO_3^-} $.
• $ \mathrm{(NH_4)_2CO_3} $: 2 $ \mathrm{NH_4^+} $ and 1 $ \mathrm{CO_3^{2-}} $.

Parentheses are used around a polyatomic ion only when more than one of that ion is required.

Ionic Bonding vs. Covalent Character

In many real compounds, bonding is not purely ionic or purely covalent. Some ionic bonds have partial covalent character.

Factors That Increase Covalent Character in an “Ionic” Bond

• High charge and/or small size of the cation: strongly attracts (polarizes) the electron cloud of the anion.
• Large and easily polarizable anion.

As polarization increases, electron density becomes more shared between ions, and the bond gains covalent character.

Examples:

• $ \mathrm{NaCl} $: largely ionic.
• $ \mathrm{MgCl_2} $: still ionic, but with more covalent character.
• $ \mathrm{AlCl_3} $: shows considerable covalent character, especially in the vapor phase.

This continuum explains why some substances that are formally “ionic” may have lower melting points or solubilities than expected for a purely ionic lattice.

Typical Examples of Ionic Compounds in Everyday Life

• Sodium chloride ($ \mathrm{NaCl} $): table salt.
• Potassium nitrate ($ \mathrm{KNO_3} $): fertilizer component, also in some propellants.
• Calcium carbonate ($ \mathrm{CaCO_3} $): main component of limestone, marble, chalk, shells.
• Magnesium sulfate ($ \mathrm{MgSO_4} $): used as a bath additive (“Epsom salt”).
• Sodium bicarbonate ($ \mathrm{NaHCO_3} $): baking soda.

Their characteristic properties—high melting point (except hydrated or basic salts), hardness, ability to conduct electricity in solution, and behavior in water—reflect the underlying nature of ionic bonding and the crystal lattice.