Structure and Bonding

Electronic Configuration of 1st and 2nd Row Elements

The electronic configuration describes the arrangement of electrons in atomic orbitals. For the first and second row elements, electrons fill the 1s, 2s, and 2p orbitals according to the Aufbau principle.

• Hydrogen (H): 1s1

• Helium (He): 1s2

• Lithium (Li): 1s2 2s1

• Beryllium (Be): 1s2 2s2

• Boron (B): 1s2 2s2 2p1

• Carbon (C): 1s2 2s2 2p2

• Nitrogen (N): 1s2 2s2 2p3

• Oxygen (O): 1s2 2s2 2p4

• Fluorine (F): 1s2 2s2 2p5

• Neon (Ne): 1s2 2s2 2p6

Example: The electronic configuration of Carbon is 1s2 2s2 2p2.

Octet Rule

The octet rule states that atoms tend to gain, lose, or share electrons to achieve a valence shell with eight electrons, similar to the noble gases.

• Applies mainly to second-row elements (C, N, O, F).

• Hydrogen follows the duet rule (2 electrons).

Example: In methane (CH4), carbon shares four pairs of electrons to complete its octet.

Lewis Structures

Lewis structures represent molecules showing all valence electrons as dots or lines for bonds.

• Count total valence electrons.

• Arrange atoms and connect with single bonds.

• Complete octets with lone pairs.

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

Example: The Lewis structure of water (H2O) shows two lone pairs on oxygen and two single bonds to hydrogen.

Condensed and Line Structures

Condensed structures write atoms in sequence, omitting some bonds and lone pairs. Line structures (skeletal structures) use lines to represent bonds, with carbon atoms at line ends or vertices and hydrogens on carbons implied.

• Example: Ethanol: Condensed: CH3CH2OH; Line: a zig-zag line with an -OH group.

Calculations

• Empirical and Molecular Formula: Given % composition and molecular weight, calculate the simplest ratio (empirical) and actual formula (molecular).

• pKa and pKb: ;

• pH Calculation: For strong acids: ; calculate moles of acid, divide by volume (L) to get [H+].

• Dipole Moment and Bond Polarity: (where is charge, is distance). Larger indicates greater polarity.

• Formal Charge:

Example: For O in H2O: 6 - 4 (lone pairs) - 2 (from two bonds) = 0.

Resonance Structures

Drawing and Evaluating Resonance Structures

Resonance structures are alternative Lewis structures for the same molecule, differing only in electron placement.

• Move only electrons, not atoms.

• All resonance forms must have the same number of electrons and obey the octet rule.

• Eliminate impossible structures (e.g., those with more than an octet on second-row elements).

Example: The acetate ion (CH3COO-) has two resonance forms with the negative charge on either oxygen.

Resonance Effect as a Stabilizing Factor

Resonance delocalizes charge, stabilizing conjugate bases and acids.

• Greater resonance stabilization leads to stronger acids (more stable conjugate base).

Example: Carboxylic acids are more acidic than alcohols due to resonance stabilization of the carboxylate anion.

Wave Functions and Molecular Orbitals

Sigma and Pi Bonds, Antibonding Orbitals

Sigma (σ) bonds are formed by head-on overlap of orbitals; pi (π) bonds by side-on overlap. Antibonding orbitals have higher energy and a node between nuclei.

• In-phase overlap: Increases electron density between nuclei (bonding).

• Out-of-phase overlap: Creates a node (antibonding).

Example: In ethylene (C2H4), the double bond consists of one σ and one π bond.

Nodes, Electron Density, and Distance

• Node: Region of zero electron density in a molecular orbital.

• Electron Density: Probability of finding an electron in a region; higher in bonding orbitals.

• Distance: Refers to internuclear separation; affects orbital overlap and bond strength.

Intermolecular Forces

Types of Intermolecular Forces

• Hydrogen Bonding: Strong dipole-dipole interaction between H bonded to N, O, or F and a lone pair on N, O, or F.

• London Dispersion Forces: Weak, temporary attractions due to instantaneous dipoles in all molecules, especially nonpolar ones.

• Dipole-Dipole Interactions: Attractions between permanent dipoles in polar molecules.

Example: Water exhibits hydrogen bonding, leading to high boiling point.

Boiling Point Trends

Boiling points increase with stronger intermolecular forces and larger molecular size.

• Hydrogen bonding > dipole-dipole > London forces

• Branching lowers boiling point by reducing surface area.

Acids and Bases

Definitions of Acids and Bases

• Arrhenius: Acids produce H+ in water; bases produce OH-.

• Brønsted-Lowry: Acids donate protons (H+); bases accept protons.

• Lewis: Acids accept electron pairs; bases donate electron pairs.

Example: NH3 is a Brønsted-Lowry base (accepts H+) and a Lewis base (donates a lone pair).

Factors Affecting Acidity

• Inductive Effects: Electronegative atoms withdraw electron density, stabilizing negative charge and increasing acidity.

• Atom Size: Larger atoms stabilize negative charge better (e.g., HI > HBr > HCl > HF in acidity).

• Resonance Stabilization: Delocalization of charge increases acidity.

• Hybridization: Greater s-character (sp > sp2 > sp3) stabilizes negative charge, increasing acidity.

Example: Acetylene (HC≡CH, sp) is more acidic than ethylene (H2C=CH2, sp2).

Polarity and Solubility

Polarity affects solubility: "like dissolves like." Polar compounds dissolve in polar solvents; nonpolar in nonpolar solvents.

• London forces: Nonpolar solutes in nonpolar solvents.

• Dipole-dipole: Polar solutes in polar solvents.

• Ionic: Ionic compounds in highly polar solvents (e.g., water).

Example: NaCl dissolves in water due to ion-dipole interactions.

Summary Table: Intermolecular Forces and Boiling Points