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 electrons with four hydrogens to complete its octet.

Lewis Structures

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

• Count total valence electrons.

• Arrange atoms and connect with single bonds.

• Distribute remaining electrons to satisfy the octet rule.

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: Lines represent bonds between carbon atoms; hydrogens on carbons are implied.

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

Calculations in Organic Chemistry

Empirical and Molecular Formula from % Composition and Molecular Weight

To determine the empirical and molecular formula:

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 the ratio.

4. Empirical formula: simplest whole-number ratio.

5. Molecular formula: , where

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

pKa, pKb, and pH Calculations

• pKa:

• pKb:

• Relationship: (at 25°C)

• pH calculation for strong acid:

• To find [H+], use moles of acid and volume:

Example: 0.01 mol HCl in 1 L:

Charge Separation and Bond Polarity

• Dipole moment (μ):

• Where is the magnitude of charge separation (in coulombs), is the distance (in meters).

• Bond polarity is predicted by the difference in electronegativity between atoms.

Example: If μ = 1.85 D and r = 1.0 Å,

Formal Charge Calculation

• Formal charge:

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

Resonance Structures

Drawing and Evaluating Resonance Structures

Resonance structures are alternative Lewis structures for a molecule that differ only in the placement of electrons.

• Move electrons, not atoms.

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

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

Example: The acetate ion has two resonance structures with the negative charge on either oxygen.

Resonance Stabilization

• Resonance delocalizes charge, stabilizing the molecule or ion.

• Conjugate bases stabilized by resonance are more acidic.

Example: The conjugate base of acetic acid is stabilized by resonance, making acetic acid more acidic than ethanol.

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 is along the internuclear axis.

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

• Antibonding orbitals (σ*, π*): Formed by out-of-phase overlap; have a node between nuclei.

• In-phase: Constructive interference, increases electron density between nuclei.

• Out-of-phase: Destructive interference, creates a node (region of zero electron density).

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

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.

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

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

Boiling Point Trends

Boiling points depend on molecular structure and intermolecular forces:

• Stronger intermolecular forces → higher boiling point.

• Hydrogen bonding > dipole-dipole > London forces.

• Branching lowers boiling point (less surface area).

Example: n-pentane has a higher boiling point 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.

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 stabilizes negative charge (sp > sp2 > sp3).

Example: Acetic acid is more acidic than ethanol due to resonance stabilization of its conjugate base.

Solubility and Polarity

Effects of Polarity on Solubility

Solubility depends on the ability of molecules to interact via intermolecular forces:

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

• London forces dominate in nonpolar compounds.

• Dipole-dipole and ionic interactions dominate in polar/ionic compounds.

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

Summary Table: Intermolecular Forces and Boiling Points

Additional info: This table summarizes the main types of intermolecular forces, their relative strengths, and their effects on boiling points.