Weakest IMF; present in all molecules

• Caused by temporary electron cloud distortions (instantaneous dipoles)
• Strength increases with molecular size/polarizability
• Only IMF for nonpolar molecules (e.g., $N_2$, $CH_4$)

Attraction between positive end of one polar molecule and negative end of another

• Stronger than LDF for similar-sized molecules
• Only in polar molecules (permanent dipole)
• Strength depends on magnitude of dipole moment

Force between ion and dipole molecule

• Strongest IMF
• Important in solutions of ionic compounds (e.g., $NaCl$ in $H_2O$)
• Ions attract partial charges on polar molecules

Special dipole-dipole interaction when H bonded to N, O, or F

• Very strong for IMF (but weaker than covalent/ionic bonds)
• Responsible for unique properties of water
• Present in $H_2O$, $NH_3$, HF, DNA, proteins

Crystalline solids: Highly ordered, repeating structure

• Sharp melting points
• Long-range order

Amorphous solids: Disordered, random structure

• No sharp melting point (soften gradually)
• Short-range order only

Solid: Fixed shape and volume; particles vibrate in fixed positions Liquid: Fixed volume, variable shape; particles flow past each other Gas: Variable shape and volume; particles move freely, far apart

Ideal gas law: $PV = nRT$

• P = pressure (atm)
• V = volume (L)
• n = moles
• R = 0.08206 L - atm/mol - K (ideal gas constant)
• T = temperature (K)

Applications:

• Calculate unknown variable when others known
• Compare two sets of conditions (combined gas law)

Dalton's Law of Partial Pressures: $P_{total} = P_1 + P_2 + P_3 + \cdots$

Mole fraction and partial pressure: $P_A = \chi_A \times P_{total}$

Distribution of molecular speeds at a given temperature

• Peak: most probable speed
• Average speed: moves right as T increases
• Curve spreads: wider distribution at higher T

$v_{rms} = \sqrt{\frac{3RT}{M}}$

Ideal gas assumptions:

• No volume occupied by gas particles
• No intermolecular forces between particles

Real gases deviate at:

• High pressure (particle volume significant)
• Low temperature (IMF significant)

van der Waals equation: $\left(P + \frac{an^2}{V^2}\right)(V - nb) = nRT$

Molarity (M): $M = \frac{\text{moles of solute}}{\text{liters of solution}}$

Other concentration units:

• Molality (m) = moles solute / kg solvent
• Mole fraction ($\chi$) = $n_A / n_{total}$
• ppm = (mass solute / mass solution) × 106

Chromatography: Separates based on differential affinity to stationary/mobile phase

• Paper chromatography
• Column chromatography
• Gas chromatography

Distillation: Separates based on boiling point differences

• Simple distillation (large BP difference)
• Fractional distillation (small BP difference)

Common soluble compounds (exceptions noted):

• Group 1 ions: all soluble
• $NH_4^+$: all soluble
• Nitrates ($NO_3^-$): all soluble
• Acetates ($CH_3COO^-$): all soluble
• Cl-, Br-, I-: soluble (except with Ag+, Pb2+, Hg22+)
• SO42-: soluble (except with Pb2+, Ba2+, Ca2+, Sr2+, Ra2+)

Common insoluble compounds:

• Carbonates ($CO_3^{2-}$), phosphates ($PO_4^{3-}$), sulfides ($S^{2-}$), hydroxides ($OH^-$)

Photoelectric effect: Emission of electrons when light shines on metal surface

• Each photon has energy $E = hf = hc/\lambda$
• Electrons emitted only if photon energy > binding energy
• Kinetic energy of ejected electron: $KE = hf - \Phi$ (work function)

$A = \varepsilon \cdot l \cdot c$

• A = absorbance (dimensionless)
• epsilon = molar absorptivity (L/mol - cm)
• l = path length (cm)
• c = concentration (mol/L)

Relationship: Absorbance directly proportional to concentration