The Ionic Bond

Ionic Bonding and Electron Transfer

An ionic bond results from the transfer of an electron from a metal atom to a non-metal atom.

Ionic Bonds

Ionic bonding is a type of chemical bond in which valence electrons are lost from one atom and gained by another. This exchange results in a more stable, noble gas electronic configuration for both atoms involved. An ionic bond is based on attractive electrostatic forces between two ions of opposite charge.

Cations and Anions

Ionic bonds involve a cation and an anion. The bond is formed when an atom, typically a metal, loses an electron or electrons, and becomes a positive ion, or cation. Another atom, typically a non-metal, is able to acquire the electron(s) to become a negative ion, or anion.

One example of an ionic bond is the formation of sodium fluoride, NaF, from a sodium atom and a fluorine atom. In this reaction, the sodium atom loses its single valence electron to the fluorine atom, which has just enough space to accept it. The ions produced are oppositely charged and are attracted to one another due to electrostatic forces.

At the macroscopic scale, ionic compounds form lattices, are crystalline solids under normal conditions, and have high melting points. Most of these solids are soluble in H₂O and conduct electricity when dissolved. The ability to conduct electricity in solution is why these substances are called electrolytes. Table salt, NaCl, is a good example of this type of compound.

Ionic bonds differ from covalent bonds. Both types result in the stable electronic states associated with the noble gases. However, in covalent bonds, the electrons are shared between the two atoms. All ionic bonds have some covalent character, but the larger the difference in electronegativity between the two atoms, the greater the ionic character of the interaction.

Lattice Energy

Lattice energy is a measure of the bond strength in an ionic compound.

Definition of Lattice Energy

Lattice energy is an estimate of the bond strength in ionic compounds. It is defined as the heat of formation for ions of opposite charge in the gas phase to combine into an ionic solid. As an example, the lattice energy of sodium chloride, NaCl, is the energy released when gaseous Na⁺ and Cl^– ions come together to form a lattice of alternating ions in the NaCl crystal.

[latex]\text{Na}^+ (g) + \text{Cl}^- (g) \rightarrow \text{NaCl} (s) \;\;\;\;\;\;\;\;\;\;\;\;\;\;\; \Delta H=-787.3\text{ kJ/mol}[/latex]

The negative sign of the energy is indicative of an exothermic reaction.

Alternatively, lattice energy can be thought of as the energy required to separate a mole of an ionic solid into the gaseous form of its ions (that is, the reverse of the reaction shown above).

Alternatively, lattice energy can be thought of as the energy required to separate a mole of an ionic solid into the gaseous form of its ions (that is, the reverse of the reaction shown above).

Lattice energy cannot be determined experimentally due to the difficulty in isolating gaseous ions. The energy value can be estimated using the Born-Haber cycle, or it can be calculated theoretically with an electrostatic examination of the crystal structure.

Factors Affecting Lattice Energy

In 1918, Born and Lande presented the following model for lattice energy:

[latex]E = - \frac {N_AMz^+z^-e^2}{4 \pi \epsilon_o r_o} (1-\frac {1}{n})[/latex]

In this equation, N_A is Avogadro’s constant; M is the Madelung constant, which depends on the crystal geometry; z⁺ is the charge number of the cation; z^– is the charge number of the anion; e is the elementary charge of the electron; n is the Born exponent, a characteristic of the compressibility of the solid; [latex]\epsilon _o[/latex] is the permittivity of free space; and r₀ is the distance to the closest ion.

This model emphasizes two main factors that contribute to the lattice energy of an ionic solid: the charge on the ions, and the radius, or size, of the ions. The effect of those factors is:

• as the charge of the ions increases, the lattice energy increases
• as the size of the ions increases, the lattice energy decreases

Lattice energies are also important in predicting the solubility of ionic solids in H₂O. Ionic compounds with smaller lattice energies tend to be more soluble in H₂O.

Formulas of Ionic Compounds

Ionic formulas must satisfy the noble gas configurations for the constituent ions and the product compound must be electrically neutral.

Ionic Compounds

An ionic bond is formed through the transfer of one or more valence electrons, typically from a metal to a non-metal, which produces a cation and an anion that are bound together by an attractive electrostatic force. On a macroscopic scale, ionic compounds, such as sodium chloride (NaCl), form a crystalline lattice and are solids at normal temperatures and pressures.

The charge on the cations and anions is determined by the number of electrons required to achieve stable noble gas electronic configurations. The ionic composition is then defined by the requirement that the resulting compound be electrically neutral overall.

For example, to combine magnesium (Mg) and bromine (Br) to get an ionic compound, we first note the electronic configurations of these atoms (valence level in indicated in italics):

Mg: 1s²2s²2p⁶3s2

Br: 1s²2s²2p⁶3s²3p⁶3d¹⁰4s24p5

In order to achieve noble gas configurations, the magnesium atom needs to lose its two valence electrons, while the bromine atom, which has 7 valence electrons, requires one additional electron to fill its outer shell. Therefore, for the resulting compound to be neutral, two bromine anions must combine with one magnesium cation to form magnesium bromide (MgBr₂). In addition, though any ratio of 2 bromine atoms to 1 magnesium atom will satisfy the two requirements above, the formula for ionic compounds is typically presented as the empirical formula, or the simplest whole-number ratio of atoms with positive integers.

Note that the cation always precedes the anion both in written form and in formulas. In the written form, while the cation name is generally the same as the element, the suffix of single-atom anions is changed to –ide, as in the case of sodium chloride. If the anion is a polyatomic ion, its suffix can vary, but is typically either –ate or –ite,as in the cases of sodium phosphate and calcium nitrite, depending on the identity of the ion.

More examples:

• lithium fluoride: Li⁺ and F^– combine to form LiF
• calcium chloride: Ca²⁺ and Cl^– combine to form CaCl₂
• iron (II) oxide: Fe²⁺ and O^2- combine to form FeO
• aluminum sulfide: Al³⁺ and S^2- combine to form Al₂S₃
• sodium sulfate: Na⁺ and SO₄^2- combine to form Na₂SO₄
• ammonium phosphate: NH⁴⁺ and PO₄^3- combine to form (NH₄)₃PO₄
• potassium chlorite: K⁺ and ClO₂^– combine to form KClO₂

Video Summary

Ionic vs Covalent Bond Character

Ionic bonds can have some covalent character.

Ionic vs Covalent Bonding

Chemical compounds are frequently classified by the bonds between constituent atoms. There are multiple kinds of attractive forces, including covalent, ionic, and metallic bonds. Ionic bonding models are generally presented as the complete loss or gain of one or more valence electrons from a metal to a nonmetal, resulting in cations and anions that are held together by attractive electrostatic forces.

In reality, the bond between these atoms is more complex than this model illustrates. The bond formed between any two atoms is not a purely ionic bond. All bonding interactions have some covalent character because the electron density remains shared between the atoms. The degree of ionic versus covalent character of a bond is determined by the difference in electronegativity between the constituent atoms. The larger the difference, the more ionic the nature of the bond. In the conventional presentation, bonds are designated as ionic when the ionic aspect is greater than the covalent aspect of the bond. Bonds that fall in between the two extremes, having both ionic and covalent character, are classified as polar covalent bonds. Such bonds are thought of as consisting of partially charged positive and negative poles.

Though ionic and covalent character represent points along a continuum, these designations are frequently useful in understanding and comparing the macroscopic properties of ionic and covalent compounds. For example, ionic compounds typically have higher boiling and melting points, and they are also usually more soluble in water than covalent compounds.