Chemical Bonding

Lewis Dot Symbols

Lewis Dot Symbols (Electron Dot Diagrams) are diagrams that represent the valence electrons of an atom or ion. These diagrams are essential for visualizing how atoms bond and interact in molecules and compounds.

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

• For Main Group Elements: Number of valence electrons = group number.

• For Transition Metals: Number of valence electrons varies.

Example: Which element will possess the most valence electrons? (Answer: Noble gases in Group 8A have the most, with 8 valence electrons.)

Drawing Lewis Dot Symbols:

1. Write the element symbol.

2. Place dots around the symbol to represent valence electrons (one per side before pairing).

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

Chemical Bonds

Chemical bonds are the attractive forces that hold atoms or ions together in a chemical compound. The three main types are ionic, covalent, and metallic bonds.

Ionic Bonding

• Occurs between metals and nonmetals.

• Involves the transfer of electrons from one atom (usually a metal) to another (usually a nonmetal), forming cations and anions.

• Ionic Bond Formation: The electrostatic attraction between oppositely charged ions.

Example: Na + Cl → Na+ + Cl-

Covalent Bonding

• Occurs between nonmetals.

• Involves the sharing of valence electrons between atoms.

Example: H2O, CO2

Metallic Bonding

• Occurs in metals.

• Involves a 'sea' of delocalized electrons moving freely among positively charged metal ions.

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

Electronegativity and Dipole Moment

Electronegativity (EN): A measure of an atom's ability to attract electrons in a bond. Increases across a period and decreases down a group.

• Dipole Moment: Occurs when a bond has a significant difference in electronegativity, resulting in a separation of charge.

Example: Calculate the difference in EN values between C and F.

Bond Classification by Electronegativity Difference

Octet Rule

The tendency of most main group elements to achieve eight valence electrons (an octet) through chemical bonding. Exceptions include hydrogen (2 electrons), boron (6 electrons), and expanded octets for elements in period 3 and beyond.

• Valence Electrons: Electrons in the outermost shell.

• Shared Electrons: Electrons shared in a covalent bond.

Formal Charge

Formal charge is used to determine the most stable Lewis structure for a molecule or ion.

• Formula:

Example: Calculate the formal charge of nitrogen in NH3.

Lewis Dot Structures (Neutral Compounds and Ions)

Lewis structures show how valence electrons are arranged among atoms in a molecule. They help predict molecular shape, reactivity, and properties.

1. Count total valence electrons.

2. Arrange atoms (least electronegative in center, except H).

3. Connect atoms with single bonds, complete octets with lone pairs.

4. Use double/triple bonds if necessary to satisfy octet rule.

5. Assign formal charges to check stability.

Example: Draw the Lewis structure for CH2O (formaldehyde).

Lone Pairs

• Lone pairs are pairs of valence electrons not involved in bonding.

Sigma and Pi Bonds

• Sigma (σ) Bond: The strongest type of covalent bond, formed by head-on overlap of orbitals.

• Pi (π) Bond: Weaker than sigma bonds, formed by side-to-side overlap of p orbitals.

• Single bond = 1 σ bond; Double bond = 1 σ + 1 π; Triple bond = 1 σ + 2 π

Lewis Structures for Ions

• Cations lose valence electrons; anions gain valence electrons.

• Draw brackets around ions and indicate charge.

Example: Draw the Lewis structure for NO3-.

Exceptions to the Octet Rule

• Incomplete Octet: Some elements (e.g., B, Be) are stable with fewer than 8 electrons.

• Expanded Octet: Elements in period 3 or higher can have more than 8 electrons.

• Free Radicals: Molecules with an odd number of electrons.

Acids and Lewis Structures

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

• Draw the Lewis structure for the acid and its conjugate base.

Resonance Structures

• Some molecules/ions have more than one valid Lewis structure (resonance).

• Resonance hybrid is the average of all resonance structures.

• Use double-headed arrows to indicate resonance.

Example: Draw all resonance structures for CO32-.

Average Charge in Resonance Structures

• Average charge = (sum of formal charges on an atom in all resonance structures) / (number of resonance structures)

Bond Energy and Enthalpy of Reaction

• Bond Energy (Bond Enthalpy): The energy required to break one mole of a bond in a molecule.

• Enthalpy of Reaction (ΔH): Calculated using bond energies:

Example: Calculate ΔH for the reaction: N2 + 3 H2 → 2 NH3

Lattice Energy

• Lattice Energy: The energy change when separated gaseous ions combine to form an ionic solid.

• Calculated using Coulomb's Law:

• Q = charge of ions, r = distance between ion centers.

• Lattice energy increases with higher charges and smaller ionic radii.

Physical Properties Related to Lattice Energy

• Higher lattice energy leads to higher melting points, boiling points, and lower solubility.

Summary Table: Bond Types and Properties

Additional info: These notes include foundational concepts for understanding chemical bonding, Lewis structures, resonance, bond energies, and lattice energies, all of which are essential for success in General Chemistry.