Intermolecular Forces and Properties — AP Chemistry
1. Molecular Interpretation of States of Matter ★★☆☆☆ ⏱ 3 min
All three common states of matter are defined by the relative strength of intermolecular forces (IMFs) between particles and the average kinetic energy (KE) of the particles themselves:
- **Solids**: IMFs are strong enough to lock particles into fixed, ordered lattice positions. Particles only vibrate, so solids have fixed shape/volume and negligible compressibility.
- **Liquids**: IMFs keep particles close but do not lock them in place. Particles move freely, so liquids have fixed volume but take container shape, with very low compressibility.
- **Gases**: IMFs are negligible relative to particle KE. Particles move freely and are far apart, so gases have no fixed shape/volume and are highly compressible.
Exam tip: Examiners require explicit reference to specific IMF types (hydrogen bonding, dipole-dipole, London dispersion) when explaining phase change temperatures for full credit, so avoid generic references to "strong forces".
2. Ideal and Real Gases ★★★☆☆ ⏱ 4 min
An ideal gas is a hypothetical model that simplifies gas behavior calculations, based on four core assumptions:
- Gas particles have negligible volume relative to the total volume of the container
- No IMFs exist between gas particles
- All collisions between particles and container walls are perfectly elastic (no kinetic energy loss)
- Average particle KE is directly proportional to absolute temperature (Kelvin)
PV = nRT
Where $P$ = pressure (atm or kPa), $V$ = volume (L), $n$ = moles of gas, $R$ = gas constant ($0.0821 \text{ L·atm/(mol·K)}$ or $8.314 \text{ J/(mol·K)}$), and $T$ = absolute temperature (Kelvin).
Real gases deviate from ideal behavior at two extreme conditions:
- **High pressure**: Particles are packed close enough that their own volume is no longer negligible, and IMFs become significant, reducing the pressure exerted on container walls
- **Low temperature**: Particles move slowly enough that IMFs alter collision trajectories, further reducing measured pressure relative to ideal predictions
\left(P + a\frac{n^2}{V^2}\right)(V - nb) = nRT
Constant $a$ accounts for IMF strength (higher $a$ = stronger IMFs) and constant $b$ accounts for particle volume (higher $b$ = larger molecules). You do not need to memorize this equation for the AP exam.
Exam tip: You do not need to memorize the van der Waals equation, but you must know what $a$ and $b$ represent, and the conditions that cause real gas deviations.
3. Solutions and Solubility ★★★☆☆ ⏱ 3 min
A solution is a homogeneous mixture of a solute (minor component) dissolved in a solvent (major component). Solubility is the maximum amount of solute that can dissolve in a fixed amount of solvent at a given temperature, governed by the "like dissolves like" rule: polar solutes dissolve in polar solvents, and nonpolar solutes dissolve in nonpolar solvents. This rule arises from energy requirements for dissolution: the process is favorable if energy released from forming new solute-solvent IMFs offsets the energy required to break existing solute-solute and solvent-solvent IMFs.
- For solid solutes: Solubility generally increases with increasing temperature (most dissolution reactions are endothermic)
- For gaseous solutes: Solubility decreases with increasing temperature, and increases with increasing partial pressure of the gas above the solution, described by Henry's Law:
C_g = k_H P_g
Where $C_g$ = dissolved gas concentration, $k_H$ = Henry's constant for the specific gas-solvent pair, and $P_g$ = partial pressure of the gas above the solution.
Exam tip: Always reference solute-solute, solvent-solvent, and solute-solvent IMFs in solubility explanations, rather than only stating "like dissolves like" to earn full marks.
4. Spectroscopy Basics (UV-Vis and IR) ★★★☆☆ ⏱ 3 min
Spectroscopy studies how matter interacts with electromagnetic radiation, and is used to identify molecular structure and measure solute concentration. Two common types you need to know for AP Chemistry are:
- **Infrared (IR) spectroscopy**: Uses infrared radiation to excite molecular vibrations (bond stretching and bending). Each bond type has a characteristic absorption frequency (wavenumbers, cm⁻¹): - O-H (hydroxyl): 3200-3600 cm⁻¹, broad peak (from hydrogen bonding) - C=O (carbonyl): ~1700 cm⁻¹, sharp, strong peak - C-H (alkyl): 2800-3000 cm⁻¹, sharp peaks
- **UV-Vis spectroscopy**: Uses ultraviolet and visible light to excite electrons from ground to excited state in molecules with conjugated pi systems or transition metal complexes. Absorbance follows the Beer-Lambert Law:
A = \varepsilon bc
Where $A$ = absorbance (unitless), $\varepsilon$ = molar absorptivity (L/(mol·cm), constant for a given substance at a specific wavelength), $b$ = path length of the sample cuvette (cm), and $c$ = solute concentration (mol/L).
5. Separation Techniques: Chromatography and Distillation ★★☆☆☆ ⏱ 3 min
Both chromatography and distillation separate mixture components based on differences in intermolecular force strength between components.
**Chromatography** separates components based on their relative affinity for a stationary phase (solid or liquid supported on a solid) and mobile phase (liquid or gas that moves through the stationary phase). The retention factor ($R_f$) for each component is calculated as:
R_f = \frac{\text{distance moved by solute}}{\text{distance moved by solvent front}}
Polar solutes have lower $R_f$ values when using a polar stationary phase (e.g., silica gel), because they form stronger IMFs with the stationary phase and move slower with the mobile phase.
**Distillation** separates liquid mixtures based on differences in boiling point, which is directly related to IMF strength. Simple distillation is used for liquids with boiling point differences >30°C, while fractional distillation uses a fractionating column to separate liquids with smaller boiling point differences. Lower boiling point liquids (with weaker IMFs) distill first.
Common Pitfalls
Why: Both involve attractive forces, so the terms are often mixed up.
Why: Celsius is more commonly used in daily life, so students forget the ideal gas law requires absolute temperature.
Why: STP is taught as the standard condition for ideal gas calculations, leading to the misconception that real gases are perfect at STP.
Why: The rule is short and easy to remember, but examiners require evidence of understanding the underlying mechanism.
Why: Both are spectral techniques, so students mix up their applications.