Study Notes: Chemical Bonding I – The Lewis Model

Chemical Bonding I: The Lewis Model

Lewis Dot Symbols

The Lewis Dot Symbol (or Electron Dot Diagram) is a visual representation of the valence electrons in an atom or ion. These diagrams are essential for understanding chemical bonding and molecular structure.

• Valence Electrons: Electrons in the outermost shell of an atom, involved in chemical bonding.

• For main group elements, the number of valence electrons equals the group number.

• For transition metals, the number of valence electrons varies.

Example: Which element will possess the most valence electrons? (Answer: Group 8A elements, e.g., Neon)

Drawing Lewis Dot Symbols

• Write the element symbol.

• Place dots around the symbol to represent valence electrons.

• Maximum of two dots per side (up to four sides).

Example: Draw the Lewis Dot Symbol for Tellurium (Te).

Chemical Bonds

Chemical bonds are the attractive forces that hold atoms together in compounds. There are three main types:

• Ionic Bonds: Formed by the transfer of electrons from one atom to another, resulting in oppositely charged ions.

• Covalent Bonds: Formed by the sharing of valence electrons between nonmetals.

• Metallic Bonds: Involve the attraction between free-flowing electrons and positively charged ions in metals.

Ionic Bonding

• Occurs between metals and nonmetals.

• Electrons are transferred from the metal to the nonmetal.

Example: Which of the following species has bonds with the most ionic character? (Answer: NaCl)

Covalent Bonding

• Occurs between nonmetals.

• Electrons are shared between atoms.

Example: Which of these elements is unlikely to form covalent bonds? (Answer: Na)

Metallic Bonding

• Occurs in metals.

• Electrons are delocalized and move freely throughout the metal lattice.

• Responsible for properties such as conductivity, malleability, and luster.

Example: Which of the following is a physical property attributed to metallic bonding? (Answer: Conductivity)

Electronegativity and Dipole Moment

Electronegativity (EN) is a measure of an atom's ability to attract electrons in a bond. Differences in EN between atoms lead to bond polarity.

• EN increases across a period and decreases down a group.

• Dipole Moment: Occurs when there is a significant difference in EN between bonded atoms, resulting in a polar bond.

Example: Calculate the difference in EN between carbon and fluorine. (Answer: 1.5)

Bond Classification by Electronegativity Difference

Bond Type	EN Difference	Example
Nonpolar Covalent	0–0.4	Cl–Cl
Polar Covalent	0.5–1.9	C–F
Ionic	≥2.0	Na–Cl

The Octet Rule

The Octet Rule states that atoms tend to gain, lose, or share electrons to achieve eight valence electrons, similar to noble gases.

• Valence Electrons: Electrons available for bonding.

• Shared Electrons: Electrons shared between atoms in a covalent bond.

Example: How many shared electrons are around the oxygen atom in H2O? (Answer: 4)

Incomplete vs. Expanded Octet

• Incomplete Octet: Some elements (e.g., H, Be, B) are stable with fewer than eight electrons.

• Expanded Octet: Elements in period 3 or higher can have more than eight electrons.

Example: How many total electrons are around the phosphorus atom in PF5? (Answer: 10)

Formal Charge

Formal Charge is a bookkeeping tool used to determine the distribution of electrons in a molecule.

• Formula:

$\text{Formal Charge} = \text{Valence Electrons} - \text{Nonbonding Electrons} - \frac{1}{2} \times \text{Bonding Electrons}$

• Helps identify the most stable Lewis structure.

Example: Determine the formal charge of nitrogen in NH3. (Answer: 0)

Lewis Dot Structures for Neutral Compounds

Lewis structures show how atoms are connected and how electrons are distributed in molecules.

• Count total valence electrons.

• Arrange atoms (central atom is usually the least electronegative).

• Distribute electrons to satisfy the octet rule.

• Assign formal charges to check stability.

Example: Draw the Lewis Dot Structure for CH2O (formaldehyde).

Lone Pairs

• Lone Pair: A pair of nonbonding electrons on an atom.

Example: How many lone pairs does sulfur have in H2S? (Answer: 2)

Lewis Dot Structures for Ions

For ions, adjust the total number of electrons to account for the charge.

• Cations: Fewer electrons than the neutral atom.

• Anions: More electrons than the neutral atom.

Example: Draw the Lewis Dot Structure for NaCl.

Lewis Dot Structures: Exceptions

• Some elements have incomplete or expanded octets.

• Free Radicals: Molecules with an unpaired electron.

Example: Draw the Lewis Dot Structure for NO (nitric oxide).

Lewis Dot Structures: Acids

• Acids are covalent compounds that release H+ ions in solution.

• Common acids: HBr, HNO3, H2SO4, H3PO4, CH3COOH

Example: Draw the Lewis Dot Structure for HNO3.

Resonance Structures

Resonance Structures are multiple valid Lewis structures for a molecule or ion, differing only in the placement of electrons.

• Use double-headed arrows to indicate resonance.

• The true structure is a hybrid of all resonance forms.

Example: Draw all possible resonance structures for CO32–.

Average Charge in Resonance Structures

• Average charge is calculated by dividing the total charge by the number of atoms sharing it.

Example: Determine the average charge of oxygen atoms in PO43–.

Bond Order

Bond Order is the average number of chemical bonds between a pair of atoms in a molecule.

• Bond order = (Total number of bonds) / (Number of bond locations)

Example: What is the bond order of the S–O bonds in SO3? (Answer: 2)

Bond Energy and Enthalpy of Reaction

Bond Energy is the amount of energy required to break a bond in a molecule. Enthalpy of Reaction ($\Delta H$) can be estimated using bond energies.

• $\Delta H = \sum \text{Bond Energies of Reactants} - \sum \text{Bond Energies of Products}$

Example: Calculate $\Delta H$ for the formation of NH3 from N2 and H2.

Standard Bond Energies Table

Bond	$\Delta H$ (kJ/mol)
S–H	347
F–H	565
F–F	159
C–H	413
C–C	347
C=O	799
C–O	358
O=O	498
O–H	467

Lattice Energy

Lattice Energy is the energy change when gaseous ions form an ionic solid. It is a measure of the strength of ionic bonds.

• Exothermic Reaction: Formation of ionic solid releases energy.

• Endothermic Reaction: Dissociation of ionic solid absorbs energy.

Formula:

$\text{Lattice Energy} = \dfrac{Q_1 Q_2}{r}$

• Where $Q_1$ and $Q_2$ are the charges of the ions, and $r$ is the distance between them.

Example: Which compound possesses the strongest ionic bond: MgBr2 or KCl? (Answer: MgBr2)

Physical Properties Related to Lattice Energy

• Higher lattice energy correlates with higher melting and boiling points.

Born-Haber Cycle

The Born-Haber Cycle is a series of steps used to calculate the lattice energy of an ionic solid from its constituent elements.

• Includes enthalpy changes for atomization, ionization, electron affinity, and formation.

Example: How many ionization energies and electron affinities are involved in forming K2O? (Answer: 2 IE, 1 EA)

Born-Haber Cycle Table

Step	$\Delta H$ (kJ/mol)
Atomization	76
Ionization	376
Electron Affinity	122
Formation	-349
Lattice Energy	-861

Example: Calculate the lattice energy for BaBr2 using the Born-Haber cycle values.

Additional info: These notes are based on the "Tro - Chemistry: A Molecular Approach" Ch.10 - Chemical Bonding I: The Lewis Model, and cover all major concepts, examples, and tables presented in the provided study prep material.