Chemical Bonding: Ionic and Covalent Bonds, Electron Configurations, and Ionic Radii - Chapter 10 - Part 1

Chemical Bonding

Introduction to Chemical Bonds

Chemical bonds are strong attractive forces that exist between certain atoms in a substance, holding them together to form compounds. There are three primary types of chemical bonds:

• Ionic bonds

• Covalent bonds

• Metallic bonds

Each type of bond involves different mechanisms of electron interaction and results in distinct physical and chemical properties.

Ionic Bonds

Describing Ionic Bonds

An ionic bond is a chemical bond formed by the electrostatic attraction between positive and negative ions. Ionic bonds typically form when one or more electrons are transferred from the valence shell of one atom (usually a metal) to the valence shell of another atom (usually a nonmetal).

• The atom that loses electrons becomes a cation (positively charged ion).

• The atom that gains electrons becomes an anion (negatively charged ion).

Example: Formation of NaCl:

• Na: $[Ne]3s^1$ loses one electron to become $Na^+$: $[Ne]$

• Cl: $[Ne]3s^2 3p^5$ gains one electron to become $Cl^-$: $[Ne]3s^2 3p^6$

Lewis Electron-Dot Symbols

A Lewis electron-dot symbol is a notation in which the valence electrons of an atom or ion are represented by dots placed around the chemical symbol. Dots are placed one to a side until all four sides are occupied.

Table: Lewis Electron-Dot Symbols for Atoms of the Second and Third Periods

Period	Li	Be	B	C	N	O	F	Ne
Second	•Li	•Be•	•B••	•C•••	•N••••	•O•••••	•F••••••	•Ne•••••••
Na	Mg	Al	Si	P	S	Cl	Ar
Third	•Na	•Mg•	•Al••	•Si•••	•P••••	•S•••••	•Cl••••••	•Ar•••••••

Formation of Ionic Bonds: Stepwise Process

• An electron is transferred between two separate atoms to give ions.

• The ions attract one another, forming an ionic bond. Both steps occur simultaneously.

Example: Using Lewis symbols to show electron transfer from Mg to F:

• Mg: $[Ne]3s^2$ loses two electrons to form $Mg^{2+}$

• Each F: $[He]2s^2 2p^5$ gains one electron to form $F^-$

Energy Involved in Forming Ionic Compounds

• Ionization energy: The energy required to remove an electron from an atom.

• Electron affinity: The energy change when an electron is added to an atom.

• Lattice energy: The change in energy when an ionic solid is separated into gas-phase ions.

Born-Haber Cycle

The Born-Haber cycle is a thermochemical cycle used to calculate lattice energy indirectly using Hess's law. It involves several steps, such as sublimation, ionization, dissociation, and formation of ions and the ionic solid.

• Sublimation of sodium: $\Delta H = 108$ kJ/mol

• Dissociation of chlorine: $\Delta H = 120$ kJ/mol

• Ionization of sodium: $\Delta H = 496$ kJ/mol

• Formation of chloride ion: $\Delta H = -349$ kJ/mol

• Formation of NaCl(s): $\Delta H = -U$ (lattice energy)

Using Hess's law, the sum of these enthalpy changes gives the overall enthalpy of formation for NaCl(s).

Properties of Ionic Substances

• Ionic substances are typically high-melting solids.

• The strength of the ionic bond depends on two factors, as given by Coulomb's law:

$E = \frac{kQ_1Q_2}{r}$

• The higher the ionic charge, the stronger the force.

• The smaller the ion, the stronger the force.

Example: The melting point of NaCl is 801°C, while that of MgO is 2800°C, due to higher charges and smaller ionic radii in MgO.

Coulomb's Law

Coulomb's law quantifies the potential energy between two charged particles:

$E = \frac{kQ_1Q_2}{r}$

• $k$ is a physical constant ($8.99 \times 10^9$ J·m/C$^2$).

• For ions with higher charges and smaller radii, the electrostatic force is much stronger.

Example Calculation:

$E = \frac{(8.99 \times 10^9 \text{ J·m/C}^2) \times (1.602 \times 10^{-19} \text{ C})^2}{2.82 \times 10^{-10} \text{ m}} = -8.18 \times 10^{-19} \text{ J}$

Electron Configurations of Ions

Main-group ions gain or lose electrons to achieve a noble-gas configuration.

• Example: Cl, $Z=17$; $Cl^-$ has 18 electrons: $1s^2 2s^2 2p^6 3s^2 3p^6$

• Transition metals often lose $ns$ electrons before $(n-1)d$ electrons, commonly forming $+2$ ions.

• Example: Mn ($Z=25$): $1s^2 2s^2 2p^6 3s^2 3p^6 3d^5 4s^2$; $Mn^{2+}$: $1s^2 2s^2 2p^6 3s^2 3p^6 3d^5$

Concept Check: Electron Configurations in Compounds

• Only ions with noble-gas configurations are commonly found in compounds.

• For example, $N^{3-}$: $[He]2s^2 2p^6$ is stable; $Na^+$: $[Ne]$ is stable.

Ionic Radii

Definition and Trends

Ionic radius is a measure of the size of the spherical region around the nucleus of an ion within which the electrons are most likely to be found. It can be measured in ionic compounds.

• A cation is always smaller than its neutral atom.

• An anion is always larger than its neutral atom.

Trends:

• Ionic radii increase down a group (more electron shells are added).

• Across a period, cations decrease in radius; when anions are reached, there is an abrupt increase, then the radius decreases again.

Isoelectronic Series

The term isoelectronic refers to different species having the same number and configuration of electrons. For an isoelectronic series, ionic radius decreases with increasing atomic number.

• Example: Ne, $Na^+$, and $F^-$ are isoelectronic.

• For $Br^-$, $Se^{2-}$, $Sr^{2+}$: $Sr^{2+} < Br^- < Se^{2-}$ (increasing radius)

• For $O^{2-}$, $F^-$, $Mg^{2+}$: $O^{2-} > F^- > Mg^{2+}$ (decreasing radius)

Covalent Bonds

Describing Covalent Bonds

A covalent bond is a chemical bond formed by sharing a pair of electrons between two atoms. This type of bond typically forms between nonmetal atoms.

• As atoms approach each other, the energy decreases to a minimum (bond formation), then increases if they get too close (repulsion).

• The distance at minimum energy is called the bond length.

Lewis Electron-Dot Formulas

A Lewis electron-dot formula uses dots to represent valence electrons. Electron pairs between atoms are bonding pairs (can be shown as lines), while pairs not involved in bonding are lone pairs or nonbonding pairs.

Coordinate Covalent Bonds

A coordinate covalent bond is formed when both electrons in the bond are donated by one atom. Once formed, the bond is indistinguishable from other covalent bonds.

The Octet Rule

Atoms tend to form covalent bonds in such a way that each atom (except hydrogen) achieves a full octet (eight electrons) in its valence shell. Hydrogen is an exception, achieving a duet (two electrons).

Types of Covalent Bonds

• Single bond: One pair of electrons shared between two atoms.

• Double bond: Two pairs of electrons shared between two atoms.

• Triple bond: Three pairs of electrons shared between two atoms.

Double bonds form primarily with C, N, and O; triple bonds form primarily with C and N.