AP Chemistry Unit 3

Intermolecular Forces (IMFs)

Forces between molecules that determine physical properties like boiling point, melting point, and solubility. Types include London dispersion, dipole–dipole, hydrogen bonding, and ion–dipole forces.

London Dispersion Forces (LDFs)

Temporary attractions from instantaneous dipoles. Present in all molecules. Strength increases with more electrons and greater surface area (e.g., He → Xe).

Why do LDFs increase down the noble gas group (He → Xe)?

Larger atoms have more electrons and more polarizable electron clouds, leading to stronger temporary dipoles.

Linear vs. Spherical Nonpolar Molecules

The linear molecule has a higher boiling point due to greater surface area → stronger London dispersion forces → more energy required to separate molecules.

Polarizability

The ease with which an electron cloud can be distorted. Increases with more electrons or delocalized π systems.

Bond Dipole vs. Molecular Dipole

Bond dipole = polarity within one bond. Molecular dipole = vector sum of all bond dipoles. Example: CO₂ has polar bonds but no net dipole (nonpolar overall).

Why is H₂O polar but CCl₄ nonpolar?

H₂O has a bent shape so dipoles add up → polar. CCl₄ is tetrahedral, dipoles symmetrically cancel → nonpolar.

Dipole–Dipole vs. Dipole-Induced Dipole

Dipole–dipole: attraction between permanent dipoles (e.g., HCl–HCl). Dipole-induced dipole: a polar molecule induces a temporary dipole in a nonpolar one (e.g., O₂ dissolved in water).

Rank IMFs from weakest to strongest

London Dispersion < Dipole–Dipole < Hydrogen Bond < Ion–Dipole < Ionic/Covalent Bonds.

Orientation of Polar Molecules for Maximum Attraction

δ⁺ end of one molecule aligns with δ⁻ end of another to maximize electrostatic attraction.

Ion–Dipole Forces

Attractions between ions and polar molecules; critical for dissolving ionic compounds like NaCl in water.

Hydrogen Bonding

Strong dipole interaction between H and N, O, or F. These are strong due to large electronegativity differences and small H size allowing close approach.

Ionic, Molecular, Metallic, and Covalent Network Solids Comparison –

Ionic: cations + anions, ionic bonds, high melting point.

Molecular: molecules, IMFs, low melting point.

Metallic: metal atoms, metallic bonds, variable melting point.

Covalent Network: atoms, covalent bonds, very high melting point.

Why Ionic Solids are Brittle but High MP

High MP due to strong electrostatic attractions; brittle because shifting layers align like charges → repulsion → fracture.

Diamond vs. Graphite (Both Carbon)

Diamond has a 3D covalent network → very hard, nonconductive. Graphite has 2D sheets → soft, conducts electricity along planes.

Molecular Solids: Low MP and Poor Conductivity

Held by weak IMFs → low MP/BP. Localized electrons → poor conductors.

Metallic Conductivity and Malleability (Sea of Electrons Model)

Delocalized electrons move freely → good conductors, malleable. Interstitial alloys add smaller atoms, reducing conductivity and malleability.

Crystalline vs. Amorphous Solids

Crystalline = ordered lattice, sharp melting point. Amorphous = disordered, gradual softening.

Liquids Have Definite Volume but No Fixed Shape

Particles are close (IMFs strong) → definite volume, but can move past each other → no fixed shape.

Why Solids and Liquids Have Similar Molar Volumes

Both dense with little empty space; gases are much less dense due to large spacing.

Hydrogen-Bonding Liquids (H₂O) vs. Non-H-Bonding (C₃H₈)

Strong H-bonds → lower vapor pressure, higher viscosity, and surface tension.

Ideal Gas Law

PV = nRT; P (atm), V (L), n (mol), R (0.0821 L·atm/mol·K), T (K). Best obeyed at low pressure and high temperature.

Kinetic Molecular Theory and Gas Pressure

Gas pressure results from collisions with container walls. Higher T → faster motion → higher pressure.

Boyle’s Law (P–V Relationship)

P ∝ 1/V; inverse relationship gives a hyperbolic curve at constant T and n.

Maxwell–Boltzmann Distribution (Temperature Effect)

Higher T → curve broadens and flattens, peak shifts right (higher average speed).

He vs. Xe Speed Distribution

At same T, He moves faster (same KE, smaller mass).

Kinetic Energy Formula

KEavg = ½mv²; average kinetic energy ∝ absolute temperature (Kelvin).

Why Gases Expand When Heated

Higher T → faster motion → greater volume at constant P. Increasing volume (constant T) → fewer collisions → lower P.

Causes of Non-Ideal Gas Behavior

1) Molecular volume (high P). 2) Intermolecular forces (low T). Deviations increase at high P, low T.

PV/RT vs. P Plot Interpretation

PV/RT < 1 → attractive forces dominate; PV/RT > 1 → repulsive/volume effects dominate.

Van der Waals Equation

(P + a(n/V)²)(V

a = IMF correction (large for polar gases), b = volume correction (large for big molecules).

Example of Ideal and Non-Ideal Gases

Nearly ideal: He (small, nonpolar). Deviates strongly: NH₃ (H-bonding, polar).

Solutions, Solutes, and Solvents

Solution = homogeneous mixture of solute (minor) and solvent (major). Homogeneous = uniform; heterogeneous = visibly distinct phases.

Molarity (M)

M = n/V; n = MV; V = n/M. Units: mol/L.

Beer–Lambert Law Variables

A = εbc. A = absorbance, ε = molar absorptivity (L·mol⁻¹·cm⁻¹), b = path length (cm), c = concentration (mol/L).

Why Absorbance Changes with Concentration

A = εbc; increasing c increases A, while ε and b stay constant.

NaCl(aq) Particulate Diagram

Na⁺ surrounded by δ⁻ O ends of water; Cl⁻ surrounded by δ⁺ H ends (ion–dipole interactions).

Dilute vs. Concentrated Solutions

Same solute-solvent interactions; only ratio changes (particle spacing).

Unsaturated, Saturated, Supersaturated Solutions

Unsat: more can dissolve. Sat: dynamic equilibrium (dissolve = crystallize). Supersat: unstable, extra solute temporarily dissolved.

Filtration vs. Chromatography

Filtration separates by size; chromatography separates true solutions by IMF differences.

Chromatography and Polarity

Polar stationary phase retains polar compounds (move less); nonpolar compounds travel farther.

Rf Value

Rf = distance solute / distance solvent front. Lower Rf = more polar.

Improving Chromatography Separation

Adjust mobile phase polarity or use 2D chromatography for better resolution.

“Like Dissolves Like” Explanation

Solution forms if solute–solvent attractions compensate for breaking solute–solute and solvent–solvent forces.

NaCl in Water vs. Hexane

Water forms strong ion–dipole attractions → dissolves. Hexane cannot stabilize ions → no dissolution.

Ethanol vs. Hexane in Water

Ethanol (H-bonding) → miscible. Hexane (nonpolar) → immiscible.

Solubility of I₂ in CCl₄ vs. H₂O

Dissolves better in nonpolar CCl₄ (LDF compatibility); water too polar.

Microwave, IR, and UV/Visible Light

Microwave = rotational transitions; IR = vibrational; UV/Visible = electronic transitions. Energy increases from microwave → UV.

Photon Relationships

c = λν; E = hν = hc/λ. Shorter λ → higher frequency → higher energy.

Absorption vs. Emission

Absorption: electron moves to higher energy level. Emission: electron falls to lower level, releasing photon (ΔE = hν).

Why Atomic Emission Spectra Have Lines

Atoms have quantized energy levels → discrete photon energies → line spectra.

Beer–Lambert Law (Equation)

A = εbc; A = absorbance (unitless), ε = molar absorptivity, b = path length (cm), c = concentration (mol/L). A is unitless because it’s a log of intensity ratio.

Beer–Lambert Law Graph Behavior

A ∝ c → linear through origin. Slope = εb (for b = 1 cm, slope = ε). Using λmax gives strongest absorption and best sensitivity.

End of Unit 3 Summary

Intermolecular forces govern structure, solubility, and phase properties. Gas laws describe ideal and real behavior. Spectroscopy connects light and matter through quantized energy transitions.