Covalent Bonding Notes

Covalent Bonding

Valence Electrons

• Valence electrons are the outer shell electrons of an atom.
• These electrons participate in chemical bonding.

Lewis Structure

• Shows how valence electrons are arranged among atoms in a molecule.
• Reflects the idea that compound stability relates to noble gas electron configuration.

Octet Rule

• Atoms share electrons to become more stable, achieving a noble-gas electron configuration ($ns^2np^6$).
• Elements form stable molecules when surrounded by eight electrons.

Duet Rule

• Hydrogen follows the duet rule, forming stable molecules by sharing two electrons ($1s^2$).

Covalent Bond

• A chemical bond where two or more electrons are shared by two atoms.
• Atoms share electrons to achieve a stable, noble-gas electron configuration ($ns^2np^6$).

Localized Electron Model

• Electron pairs are localized on a particular atom or in the space between two atoms.
• Lone pairs (lp): pairs of electrons localized on an atom.
• Bonding pairs (bp): pairs of electrons found in the space between the atoms.

Predicting Number of Bonds

• Atoms gain access to shared electrons through bonding.
• Period 2 elements: the number of electrons needed to have a full valence shell predicts the number of bonds it is most likely to make.
• Period 3 and beyond elements can form exceptions to the octet rule.

Single, Double, and Triple Bonds

• One shared pair: single bond
• Two shared pairs: double bond
• Three shared pairs: triple bond

Lewis Structures

• Used for molecules and polyatomic ions following the octet rule.

Drawing Lewis Structures: Basic Rules

• Draw the atomic symbol.
• Count valence shell electrons.
• Fill boxes (sides) with electrons, pairing when necessary.
• Placement of dots is arbitrary for individual Lewis symbols.

Valid Lewis Structures

• The number of electrons equals the total valence electrons in the molecule's atoms.
• Atoms obey the octet/duet rule.

Multiple Bonds

• Common in C, O, N, S, P.

Guidelines for Drawing Lewis Structures

1. Sum valence electrons for all atoms.
2. Select a central atom (usually the one needing more bonds to become stable, but never hydrogen).
3. Connect the central atom to outer atoms with single bonds.
4. Distribute remaining valence electrons to outer atoms as lone pairs to complete octets (H is an exception).
5. Assign any remaining electrons to the central atom.
6. If the central atom lacks an octet, move lone pairs from outer atoms to form double or triple bonds as needed.
7. Verify the structure fulfills the duet rule for hydrogen, the octet rule (with exceptions), and contains the correct number of valence electrons.

Lewis Structures of Polyatomic Ions

• Follow the same guidelines as neutral molecules with minor additions.
• Consider the charge when summing valence electrons.
• Add an electron for each negative charge (anion) and subtract for each positive charge (cation).
• Place the structure within brackets, indicating the charge outside.

Resonance and Formal Charge

• Resonance structures occur when multiple valid Lewis structures can be drawn for a molecule, differing only in the placement of multiple bonds and lone pairs.

Formal Charge

• Used to determine the best Lewis structure when multiple valid structures exist.
• Formal charge = # valence e - (# shared e / 2 + # unshared e)
• Most plausible Lewis structure: small or zero formal charges, negative formal charges on more electronegative atoms, and positive formal charges on less electronegative elements.

Exceptions to the Octet Rule

• Incomplete Octet: Be (4 electrons), B (6 electrons).
• Odd-Electron Molecules (Radicals): Molecules with unpaired electrons (e.g., NO, $NO_2$).
• Expanded Octet: Elements in the third period or beyond can have more than eight electrons.

Polar Bonds and Electronegativity

• Electronegativity: ability of a bonded atom to attract shared electrons.
• Increases across a row and up a group in the periodic table.

Electronegativity and Bond Polarity

• Uneven sharing of electrons forms polar covalent bonds.
• The greater the electronegativity difference, the more polar the bond.

Bonding Continuum

• Bonds range from nonpolar covalent to polar covalent to ionic, based on electron sharing.

Dipole Moment

• $μ = q × r$
• Polar covalent bonds are dipoles with separated areas of opposite charge.

Bond Enthalpy

• Bond strength measured as bond enthalpy, the enthalpy change when breaking a specific bond in 1 mol of gaseous molecules.
• $ΔH_{rxn} = Σ(ΔH \text{ of bonds broken}) – Σ(ΔH \text{ of bonds formed})$

Bond Length and Bond Enthalpy

• As bond length increases, bond enthalpy generally decreases.
• Triple bonds are shortest and strongest; single bonds are longest and weakest.