Electron Configuration, Intermolecular Forces, and Chemical Reactivity Review

Electron Configuration

The distribution of electrons within atoms or ions.

Subshell Capacities

The maximum electron counts for types s, p, d, and f, which are $2\text{ e}^-$, $6\text{ e}^-$, $10\text{ e}^-$, and $14\text{ e}^-$ respectively.

Hund's Rule

Electrons occupy orbitals singly before pairing up.

Intermolecular Forces (IMFs)

Non-bonding attractive forces that occur between separate molecules and dictate physical properties like phase changes and solubility.

Intramolecular Forces

Chemical bonds, such as covalent, ionic, or metallic bonds, that hold atoms together within a single molecule.

Dipole-Dipole Interactions

Electrostatic attractions between polar molecules with a permanent, asymmetrical distribution of electron density.

Hydrogen Bonding

A specialized, strong dipole-dipole interaction occurring when a hydrogen atom is directly bonded to Nitrogen (N), Oxygen (O), or Fluorine (F).

London Dispersion Forces (LDFs)

Intermolecular forces that increase with molecular size and surface area as molecules with more electrons become more polarizable.

Amphiphilic Balance

The dual nature of molecules like alcohols, which possess a polar hydrophilic head ($-OH$) and a nonpolar hydrophobic hydrocarbon tail.

First Ionisation Energy

The energy required to remove one mole of electrons from one mole of atoms in their gaseous state to form one mole of $1+$ ions in their gaseous state.

Successive Ionisation Energies

The energy values associated with removing sequential electrons one by one from an atom or ion.

Nitrate

$$NO_3^-$$

Phosphate

$$PO_4^{3-}$$

Dichromate (VI)

$$Cr_2O_7^{2-}$$

Manganate (VII)

$$MnO_4^-$$

Ideal Gas Law

$$pV = nRT$$

Arrhenius Equation (Logarithmic Form)

$$\text{ln}(k) = -\frac{E_a}{RT} + \text{ln}(A)$$

Equilibrium Constant (Kc)

$$K_c = \frac{[C]^c[D]^d}{[A]^a[B]^b}$$

pH

$$\text{pH} = -\text{log}_{10}([H^+])$$

Water Constant (Kw)

$K_w = [H^+][OH^-]$, which equals $1.00 \times 10^{-14} \text{ mol}^2 \text{ dm}^{-6}$ at $298 \text{ K}$.

Gibbs Free Energy

$$\triangle G = \triangle H - T\triangle S$$

Hess's Law (Enthalpy of Formation)

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