lecture 15

Chemical Bonding and Lewis Structures

Introduction to the Lewis Structure Model

The Lewis Structure Model is a foundational concept in general chemistry for representing the valence electrons of atoms and molecules. It uses dots to symbolize valence electrons, which are the electrons involved in chemical bonding.

• Valence electrons are shown as dots around the chemical symbol of an element.

• Lewis electron-dot structures (Lewis structures) depict both the structural formula and the valence electrons.

• Lewis structures focus on valence electrons because chemical bonding involves the transfer or sharing of these electrons between atoms.

• Example: Oxygen atom with 6 valence electrons: O with 6 dots around it.

The Octet Rule: A Guideline for Molecule Formation

The octet rule states that atoms tend to gain, lose, or share electrons to achieve a noble gas-like configuration, typically with eight valence electrons (an octet).

• For main-group elements, especially nonmetals in period 2, the octet rule is strictly followed.

• The electron configuration for a noble gas is $ns^2np^6$.

• Expanded octets: Nonmetals in period 3 and below can have more than eight electrons when they are the central atom, due to available d-orbitals.

• Nonmetals not serving as the central atom generally follow the octet rule.

• Additional info: Quantum theory predicts the use of empty valence d-orbitals for expanded octets.

Exceptions to the Octet Rule

Some elements do not follow the octet rule due to their small size or electron configuration.

• Hydrogen (H): Can have only two valence electrons (a duet).

• Lithium (Li): Loses its one valence electron to achieve a stable configuration.

• Beryllium (Be): Commonly shares its two electrons in covalent bonds, resulting in four valence electrons.

• Boron (B): Commonly shares its three electrons in covalent bonds, resulting in six valence electrons.

Octet and Duet Stability

Atoms with eight valence electrons (octet) or two valence electrons (duet, for hydrogen and helium) are particularly stable.

• Octet: Full outer principal energy level, $n$.

• Duet: Only two valence electrons fill the principal level for small atoms like H and He.

• Example: Lewis symbol for oxygen: O with 6 dots.

Types of Chemical Bonds

Ionic and Covalent Bonds

Chemical bonds are formed by the transfer or sharing of electrons to achieve stable electron configurations.

• Ionic bond: Transfer of electrons from a metal to a nonmetal, resulting in cations and anions that attract each other.

• Covalent bond: Sharing of electrons between two nonmetal atoms.

• Bonding atoms generally attain stable electron configurations, usually with eight electrons in their outermost shell (octet rule).

Lewis Symbols and Practice

Lewis Symbols for Elements

Lewis symbols provide a simple way to visualize the number of valence electrons in a main-group atom.

• Example: Silicon (Si) has four valence electrons, so its Lewis symbol is Si with four dots.

Ionic Bonding and Formula Units

Nature of Ionic Compounds

Ionic compounds are composed of cations (usually metals) and anions (usually nonmetals). The basic unit is the formula unit, which is not a molecule but the smallest electrically neutral collection of ions.

• Formula units exist as part of a large lattice structure.

• Example: Table salt (NaCl) is composed of Na+ and Cl- ions in a lattice.

Lewis Model and Bonding in Ionic Compounds

The Lewis model can be used to illustrate the transfer of electrons from metals to nonmetals, resulting in ionic bonds.

• 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+: $1s^2 2s^2 2p^6 3s^2 3p^6$

Predicting Ionic Compound Formulas Using Lewis Symbols

Lewis theory predicts the number of electrons a metal atom should lose or a nonmetal atom should gain to attain a stable electron arrangement.

• The octet rule helps predict the formulas of ionic compounds and the relative strengths of ionic bonds (Coulomb's law).

• Example: Calcium (Ca) and chlorine (Cl): Ca loses two electrons, each Cl gains one electron, so the formula is CaCl2.

Properties of Ionic Solids

Physical Properties Predicted by Lewis Theory

Lewis theory implies several key properties of ionic solids:

• Hardness: Ionic solids are relatively hard due to the strong electrostatic forces in the crystal lattice.

• Brittleness: Displacement of ions leads to repulsive forces, causing the crystal to break apart.

• Electrical Conductivity: Ionic solids do not conduct electricity in the solid state because ions are locked in place. However, when melted or dissolved in water, ions can move and conduct electricity.

Crystal Lattice and Lattice Energy

The crystal lattice is a structure in which every cation is surrounded by anions and vice versa, maximizing electrostatic attractions and stability.

• Lattice energy is the energy released when the solid crystal forms from separate ions in the gas state. It is always exothermic.

• Lattice energy depends directly on the size of the charges and inversely on the distance between ions.

• Example: Formation of NaCl: $\text{Na}(s) \rightarrow \text{Na}^+(g) + e^- \quad \Delta H^\circ = +496\ \text{kJ/mol}$ $\text{Cl}_2(g) + e^- \rightarrow \text{Cl}^-(g) \quad \Delta H^\circ = -349\ \text{kJ/mol}$ $\text{Na}(s) + \frac{1}{2}\text{Cl}_2(g) \rightarrow \text{NaCl}(s) \quad \Delta H_f = -411\ \text{kJ/mol}$

Writing and Naming Ionic Compounds

Writing Formulas for Ionic Compounds

The formula of an ionic compound reflects the smallest whole-number ratio of ions, ensuring charge neutrality.

• Sum of positive charges (cations) must equal sum of negative charges (anions).

• Procedure:

  1. Write the symbol for the metal cation and its charge, followed by the symbol for the nonmetal anion and its charge.

  2. Adjust subscripts to balance the total charge.

  3. Check that the sum of charges equals zero.

• Example: Calcium (Ca2+) and oxygen (O2-): Formula is CaO.

Types of Metals in Ionic Compounds

Metals can be categorized based on whether their charge is invariant or variable in compounds.

• Type I: Metals with invariant charge (alkali metals, alkaline earth metals, Al, Zn, Cd, Ag).

• Type II: Metals with variable charge (transition metals, some p-block metals like Pb, Tl, Sn).

Naming Binary Ionic Compounds

Binary ionic compounds contain only two different elements. Their names follow a specific pattern:

• Name of the cation (metal) followed by the base name of the anion (nonmetal) with the ending -ide.

• Example: KCl is potassium chloride; CaO is calcium oxide.

Common Monoatomic Anions

Element	Symbol	Base Name	Anion Name
Phosphorus	P3-	Phosph-	Phosphide
Chlorine	Cl-	Chlor-	Chloride
Oxygen	O2-	Ox-	Oxide
Bromine	Br-	Brom-	Bromide
Sulfur	S2-	Sulf-	Sulfide
Additional info:		Other common anions include nitride (N3-), fluoride (F-).

Naming Ionic Compounds with Multivalent Metals (Type II)

For metals that can form more than one type of cation, the charge is indicated by a Roman numeral in parentheses after the metal name.

• Example: Cu2O is copper(I) oxide; CuO is copper(II) oxide.

• Older naming system uses suffixes: -ous for lower charge, -ic for higher charge (e.g., ferrous vs. ferric), but the Roman numeral system is preferred.

Practice: Determining Charges and Naming Compounds

• To name CrBr3: Chromium must have a 3+ charge, so the name is chromium(III) bromide.

• To name SnCl2: Tin must have a 2+ charge, so the name is tin(II) chloride.

• To name PbCl4: Lead must have a 4+ charge, so the name is lead(IV) chloride.

Comparing Melting Points of Ionic Compounds

Effect of Ionic Charge on Melting Point

The melting point of ionic compounds depends on the magnitude of the charges of the ions.

• Compounds with higher charges (e.g., MgO with Mg2+ and O2-) have higher melting points than those with lower charges (e.g., NaCl with Na+ and Cl-).

• Reason: Greater charge leads to stronger electrostatic attraction and higher lattice energy.

Summary Table: Writing and Naming Ionic Compounds

Step	Description
1	Write the symbol and charge for each ion.
2	Adjust subscripts to balance charges.
3	Check that the total charge is zero.
4	Name the compound: cation name + anion base name + -ide (or Roman numeral for multivalent metals).

Additional info: These notes cover the essential concepts of chemical bonding, Lewis structures, ionic compounds, and their properties, suitable for General Chemistry exam preparation.