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, 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

Organic compounds can be represented in different ways:

• Condensed structure: Shows all atoms, e.g., CH3CH2OH for ethanol.

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

Example: Butane can be written as CH3CH2CH2CH3 (condensed) or as a zig-zag line (line structure).

Empirical and Molecular Formula Calculations

Given percent composition and molecular weight, you can determine:

• 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 smallest number of moles to get ratio.

4. Multiply to get whole numbers if necessary.

5. Use molecular weight to find the molecular formula.

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 the nitrate ion (NO3-), the formal charge on the central N is +1.

Bond Polarity and Dipole Moment

Bond polarity arises from differences in electronegativity. The dipole moment () quantifies the separation of charge:

• Formula:

• Where is the magnitude of charge and is the distance between charges.

Example: The dipole moment of HCl can be calculated if and are known.

Resonance Structures

Drawing and Evaluating Resonance Structures

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

• Move only electrons, not atoms.

• All resonance forms must have the same arrangement of atoms.

• Best resonance structures have full octets and minimal formal charges.

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

Resonance Stabilization

Resonance stabilizes molecules by delocalizing charge, 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 ion.

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.

• Bonding orbitals: In-phase overlap increases electron density between nuclei.

• Antibonding orbitals: Out-of-phase overlap creates a node (region of zero electron density).

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

Wavefunction Graphs

• In-phase: Constructive interference, bonding orbital.

• Out-of-phase: Destructive interference, antibonding orbital (node present).

• Node: Region where probability of finding an electron is zero.

Example: The σ* (sigma antibonding) orbital has a node between nuclei.

Intermolecular Forces

Types of Intermolecular Forces

• Hydrogen bonding: Strong dipole-dipole interaction involving H bonded to N, O, or F.

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

• 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 depend on molecular structure and intermolecular forces:

• Stronger forces (e.g., hydrogen bonding) = higher boiling point.

• Branching lowers boiling point (less surface area).

Example: n-Butanol (hydrogen bonding) has a higher boiling point than diethyl ether (no H-bonding).

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 electron pair).

pKa, pKb, and pH Calculations

• pKa:

• pKb:

• pH:

Example: To calculate pH, convert grams of acid to moles, divide by volume (L) to get [H+], then use the formula above.

Factors Affecting 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 (sp > sp2 > sp3) stabilizes negative charge, increasing acidity.

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: Important for nonpolar solutes/solvents.

• Dipole-dipole and ionic interactions: Important for polar and ionic compounds.

Example: NaCl dissolves in water due to ion-dipole interactions; hexane does not.

Summary Table: Intermolecular Forces and Properties

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