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 filled valence shell with eight electrons, resembling the noble gas configuration.

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

• Hydrogen follows the "duet rule" (2 electrons).

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

Lewis Structures

Lewis structures are diagrams showing the bonding between atoms and the lone pairs of electrons in a molecule.

• Each line represents a pair of shared electrons (a bond).

• Lone pairs are shown as dots.

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

Condensed and Line Structures

Chemical structures can be represented in different ways:

• Condensed structure: Atoms are written in sequence, e.g., CH3CH2OH for ethanol.

• Line structure (skeletal): Lines represent bonds between carbons; hydrogens on carbons are implied.

Example: The line structure for butane is a zig-zag line with four carbon atoms.

Empirical and Molecular Formula Calculations

Given percent composition and molecular weight, you can determine the empirical and molecular formulas.

• Empirical formula: Simplest whole-number ratio of atoms.

• Molecular formula: Actual number of atoms in a molecule.

Steps:

1. Convert % composition to grams (assume 100 g sample).

2. Convert grams to moles for each element.

3. Divide by the smallest number of moles to get ratios.

4. Multiply to get whole numbers if necessary.

5. Calculate molecular formula using molecular weight.

Example: A compound with 40% C, 6.7% H, and 53.3% O (by mass) and MW = 180 g/mol has empirical formula CH2O; molecular formula is C6H12O6.

Formal Charge Calculation

Formal charge is used to determine the charge on an atom in a molecule.

• Formula:

Example: In O3, the central oxygen has a formal charge of +1.

Resonance Structures

Drawing and Evaluating Resonance Structures

Resonance structures are different Lewis structures for the same molecule, showing delocalization of electrons.

• Only electrons move; atom positions do not change.

• Resonance structures must have the same number of electrons and obey the octet rule.

• Impossible structures violate octet or have incorrect charges.

Example: The acetate ion (CH3COO-) has two resonance structures with the negative charge on different oxygens.

Resonance Effect as a Stabilizing Factor

Resonance stabilizes molecules by delocalizing charge or electron density, especially in 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

Molecular orbitals are formed by the combination of atomic orbitals.

• Sigma (σ) bond: Formed by head-on overlap; electron density along the internuclear axis.

• Pi (π) bond: Formed by side-on overlap; electron density above and below the plane.

• Antibonding orbitals (σ*, π*): Higher energy, formed by out-of-phase overlap.

Wavefunction graphs show regions of electron density (in-phase) and nodes (out-of-phase, zero density).

Intermolecular Forces

Types of Intermolecular Forces

• Hydrogen bonding: Strong dipole-dipole interaction between H and N, O, or F.

• London dispersion forces: Weak, temporary attractions due to instantaneous dipoles; present in all molecules.

• Dipole-dipole interactions: Attractions between polar molecules.

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

Boiling Point Trends

Boiling points depend on molecular structure and intermolecular forces.

• Stronger intermolecular forces = higher boiling point.

• Branching lowers boiling point (less surface area).

Example: n-pentane boils at a higher temperature than neopentane due to less branching.

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 and a Lewis base.

pKa, pKb, and pH Calculations

• pKa:

• pKb:

• pH:

Example: Calculate pH for 0.01 M HCl:

Effects on Acidity

• Inductive effects: Electronegative atoms withdraw electron density, stabilizing conjugate base 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 increases acidity (sp > sp2 > sp3).

Example: Acetylene (C2H2, sp-hybridized) is more acidic than ethylene (sp2) or ethane (sp3).

Solubility and Polarity

Effects of Polarity on Solubility

Polarity and intermolecular forces determine solubility.

• "Like dissolves like": Polar solutes dissolve in polar solvents; nonpolar in nonpolar.

• Forces: London dispersion (all molecules), dipole-dipole (polar), ionic (ions in polar solvents).

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

Sample Table: Intermolecular Forces and Boiling Points

Additional info: Academic context and examples have been added to clarify and expand upon the original study list, ensuring the notes are self-contained and suitable for exam preparation.