Chapter 9: Chemical Bonding I – The Lewis Model

Chemical Bonding I: The Lewis Model

Introduction to Bonding Theories

Chemical bonding theories are essential for understanding how and why atoms combine to form molecules. These theories help explain molecular stability, predict molecular shapes, and determine the chemical and physical properties of compounds. The Lewis model is one of the simplest and most widely used bonding theories in general chemistry.

• Bonding theories explain the attachment and stability of atoms in molecules.

• They help predict why certain combinations (e.g., H2O) are stable while others (e.g., HO or H3O) are not.

• Bonding theories are used in pharmaceutical design, such as creating protease inhibitors to disable HIV-protease by simulating molecular shapes and interactions.

The Lewis Model

The Lewis theory focuses on the role of valence electrons in chemical bonding. It provides a visual representation of molecules using Lewis structures (electron dot structures), which allow prediction of molecular properties such as stability, shape, size, and polarity.

• Lewis theory emphasizes valence electrons in bonding.

• Lewis structures use dots to represent valence electrons around atomic symbols.

• These models help predict molecular behavior and properties.

Why Do Atoms Bond?

Atoms bond to lower their potential energy, resulting in more stable arrangements. The formation of a chemical bond occurs when the potential energy of bonded atoms is less than that of separate atoms. Key interactions include:

• Nucleus-to-nucleus repulsions

• Electron-to-electron repulsions

• Nucleus-to-electron attractions

Types of Bonds

Bonds are classified based on the types of atoms involved and the nature of electron interaction.

Types of Atoms	Type of Bond	Characteristic of Bond
Metal and nonmetal	Ionic	Electrons transferred
Nonmetal and nonmetal	Covalent	Electrons shared
Metal and metal	Metallic	Electrons pooled

Ionic Bonds

Ionic bonds form between metals and nonmetals through electron transfer. Metals lose electrons to become cations, while nonmetals gain electrons to become anions. The resulting oppositely charged ions attract each other, forming an ionic bond.

• Metals have low ionization energy (easy to lose electrons).

• Nonmetals have high electron affinity (easy to gain electrons).

• Example: NaCl formation from Na+ and Cl-.

Covalent Bonds

Covalent bonds occur between nonmetals, which have high ionization energies. Instead of transferring electrons, atoms share valence electrons, resulting in lower potential energy and stable molecules.

• Shared electrons are attracted by the nuclei of both atoms.

• Example: H2O molecules, where electrons are shared between H and O.

Metallic Bonds

Metallic bonds involve the pooling of valence electrons among metal atoms. The electrons are delocalized, forming a 'sea of electrons' that surrounds metal cations, resulting in strong metallic bonding.

• Metals have low ionization energy, allowing easy electron loss.

• Delocalized electrons provide conductivity and malleability.

• Example: Sodium metal (Na).

Valence Electrons and Bonding

Valence electrons are the outermost electrons and are most loosely held. Chemical bonding involves the transfer or sharing of these electrons, making them crucial in bond formation. Lewis theory specifically focuses on valence electron behavior.

Determining the Number of Valence Electrons

The number of valence electrons in a main group atom corresponds to its column number on the periodic table. Transition elements typically have two valence electrons due to their electron configuration.

Lewis Structures of Atoms

Lewis structures represent valence electrons as dots around the atomic symbol. The symbol stands for the nucleus and inner electrons, while dots indicate valence electrons.

• First two dots: s orbital electrons (paired).

• Next three dots: p orbital electrons (single on each side).

• Remaining dots: paired for p electrons.

• Example: Oxygen atom with six dots for six valence electrons.

Lewis Bonding Theory and the Octet Rule

Atoms bond to achieve a more stable electron configuration, often resulting in an outer shell with eight electrons (the octet rule). This is achieved by transferring or sharing electrons. Some exceptions exist, but the goal is to mimic noble gas configurations.

Stable Electron Arrangements and Ion Charge

Metals form cations by losing valence electrons, while nonmetals form anions by gaining them. The resulting electron configuration is often that of a noble gas.

• Example: Potassium (K) loses one electron to become K+, achieving an octet in the previous energy level.

• Electron configuration for K: $1s^2 2s^2 2p^6 3s^2 3p^6 4s^1$

• Electron configuration for K+: $1s^2 2s^2 2p^6 3s^2 3p^6$

Lewis Theory and Ionic Bonding

Lewis symbols can visually represent the transfer of electrons in ionic bonding. For example, the transfer from K to Cl forms K+ and Cl- ions, which are then attracted to each other.

• Example: $\text{K} + :\text{Cl}: \rightarrow \text{K}^+ [:\text{Cl}:]^-$

Lewis Theory Predictions for Ionic Bonding

Lewis theory predicts the number of electrons lost or gained to achieve stability (octet rule), allowing prediction of ionic compound formulas and bond strengths (using Coulomb's law).

Ionic Bonding and the Crystal Lattice

The formation of a crystal lattice releases extra energy, stabilizing the structure. Each cation is surrounded by anions and vice versa, maximizing electrostatic attractions for stability.

• Crystal lattice is held together by electrostatic attraction.

• Leads to highly stable ionic solids.

Crystal Lattice

Electrostatic attraction in ionic solids is nondirectional, meaning there is no specific anion-cation pair. The chemical formula is empirical, representing the ratio of ions for charge balance.

Lattice Energy

Lattice energy is the energy released when a solid crystal forms from separate ions in the gas state. It is always exothermic and depends on the size and charge of the ions.

• Lattice energy increases with higher ionic charge and decreases with larger ionic radius.

Determining Lattice Energy: The Born–Haber Cycle

The Born–Haber cycle is a series of hypothetical reactions used to calculate lattice energy by summing known enthalpy changes (using Hess's law).

• Key steps include atomization, ionization, electron affinity, and lattice formation.

• Equation: $\Delta H^\circ(\text{salt}) = \Delta H^\circ(\text{metal atoms, g}) + \Delta H^\circ(\text{nonmetal atoms, g}) + \Delta H^\circ(\text{cations, g}) + \Delta H^\circ(\text{anions, g}) + \Delta H^\circ(\text{crystal lattice})$

Born–Haber Cycle for NaCl

Example calculation for NaCl:

• Na(s) → Na(g)  +108 kJ

• ½ Cl2(g) → Cl(g)  +½(244 kJ)

• Na(g) → Na+(g)  +496 kJ

• Cl(g) → Cl-(g)  -349 kJ

• Na+(g) + Cl-(g) → NaCl(s)  ΔH(NaCl lattice)

• Na(s) + ½ Cl2(g) → NaCl(s)  -411 kJ

Calculation:

• $\Delta H^\circ(\text{NaCl, s}) = \Delta H^\circ(\text{Na atoms, g}) + \Delta H^\circ(\text{Cl–Cl bond energy}) + \Delta H^\circ(\text{IE of Na}) + \Delta H^\circ(\text{EA of Cl}) + \Delta H^\circ(\text{NaCl lattice})$

• NaCl lattice energy = (-411) - [(+108) + (+122)] + (+496) + (-349) = -788 kJ

Trends in Lattice Energy: Ion Size and Charge

Lattice energy is affected by both ion size and charge:

• Force of attraction is inversely proportional to distance (larger ions = lower lattice energy).

• Force of attraction is directly proportional to charge (higher charge = higher lattice energy).

• Ion charge generally has a greater effect than ion size.

Metal Chloride	Lattice Energy (kJ/mol)	Ion Radius (pm)
LiCl	-834	Li+: 241
NaCl	-788	Na+: 276
KCl	-701	K+: 314
CsCl	-657	Cs+: 348

Ionic Bonding Model versus Reality

Lewis theory predicts strong attractions between ions, resulting in high melting and boiling points for ionic compounds. All ionic compounds are solids at room temperature, and they are hard and brittle crystalline solids.

• Melting points generally > 300 °C.

• Liquid state conducts electricity; solid state does not.

• Many ionic compounds are water-soluble and conduct electricity in solution.

Properties of Ionic Compounds

• Hard and brittle crystalline solids

• High melting points

• Conductivity in liquid and aqueous states

• Solubility in water

Example: NaCl(s) ions are fixed; NaCl(l) ions can move and conduct electricity.

Hardness and Brittleness

Ionic solids are hard because ion positions in the crystal lattice are critical for stability. Displacing ions leads to repulsive forces, making the crystal brittle and prone to shattering when struck.

• Hardness is tested by rubbing materials together; the harder material cuts or does not streak.

• Ionic solids are harder than most molecular solids.

• Brittleness results from instability when ions are displaced.