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

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 carbons; hydrogens on carbons are implied.

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

Calculations in Organic Chemistry

Empirical and Molecular Formula from % Composition and Molecular Weight

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

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

• Convert grams to moles for each element.

• Divide by the smallest number of moles to get the empirical formula.

• Calculate the empirical formula mass and divide the molecular weight by this value to find the molecular formula.

Example: A compound with 40% C, 6.7% H, and 53.3% O and a molecular weight of 180 g/mol has an empirical formula of CH2O and a molecular formula of C6H12O6.

pKa, pKb, and pH Calculations

• pKa:

• pKb:

• pH:

• For strong acids: Calculate moles of acid, divide by volume (L) to get [H+], then use the pH formula.

Example: 0.01 mol HCl in 1 L solution: [H+] = 0.01 M, so .

Charge Separation and Bond Polarity

• Dipole moment (μ):

• q = charge separation (in coulombs), r = distance (in meters).

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

Example: For a bond with μ = 1.85 D and r = 1.0 Å, calculate q.

Formal Charge Calculation

• Formal charge:

• Calculate for each non-hydrogen atom in a molecule.

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

Resonance Structures

Drawing and Evaluating Resonance Structures

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

• Move only electrons, not atoms.

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

• 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 different oxygens.

Resonance Stabilization

• 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 ion.

Wave Functions and Molecular Orbitals

Sigma and Pi Bonds, Antibonding Orbitals

Molecular orbitals are formed by the combination of atomic orbitals. The wave function describes the probability of finding an electron in a region of space.

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

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

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

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

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 and 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 are influenced by molecular structure and intermolecular forces.

• Stronger intermolecular forces = higher boiling point.

• Branching lowers boiling point (less surface area for London forces).

• Hydrogen bonding increases boiling point significantly.

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

Acids and Bases

Definitions of Acids and Bases

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.

Polarity and Solubility

Polarity affects solubility in solvents:

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

• London forces, dipole-dipole, and ionic interactions all contribute to solubility.

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

Additional info: Academic context and examples have been added to clarify and expand upon the original study list.