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Unit 3
18–22% of examThis is the biggest unit on the exam. You'll use the forces between particles to explain boiling points, solubility and the properties of solids, model gases with the ideal gas law and kinetic molecular theory, and work out solution concentrations. It ends with how light interacts with matter, including how a spectrophotometer measures concentration.
Longer videos that cover the whole unit. Good for a first pass or a final review.
Intermolecular forces are attractions between separate particles, and they are much weaker than the bonds inside molecules. London dispersion forces act between all molecules and get stronger with more electrons and bigger, more easily distorted (polarizable) electron clouds; polar molecules add dipole-dipole forces, molecules with H bonded to N, O or F can hydrogen bond, and ions attract polar molecules through ion-dipole forces.
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The particles in a substance and the forces between them explain its properties. Ionic solids have high melting points and conduct only when melted or dissolved, covalent network solids like diamond and SiO₂ are hard with very high melting points, molecular solids melt easily and don't conduct, and metals conduct and can be bent into shape. In liquids, stronger intermolecular forces mean a lower vapor pressure and a higher boiling point.
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In a solid, particles stay in place and only vibrate; they can be neatly ordered (crystalline) or disordered (amorphous). Liquid particles stay in contact but slide past each other, so a solid and its liquid have similar volumes, while gas particles are far apart and move freely, filling whatever container they are in.
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The ideal gas law, PV = nRT, links the pressure, volume, moles and Kelvin temperature of a gas. In a mixture, each gas has its own partial pressure, the total pressure is the sum of them (Dalton's law), and each gas's partial pressure equals its mole fraction times the total pressure.
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Kinetic molecular theory explains gas behavior with particles in constant, random motion. The Kelvin temperature is proportional to the particles' average kinetic energy (KE = ½mv²), so at the same temperature lighter particles move faster on average; a Maxwell-Boltzmann distribution shows the spread of particle speeds, which flattens and shifts to higher speeds as temperature rises.
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Real gases stop behaving ideally when their particles attract each other, especially at low temperatures close to condensing, and when they are squeezed to very high pressures so the particles' own volume matters. Attractions make the measured pressure lower than PV = nRT predicts, while particle volume makes the real volume larger than predicted.
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A solution is a homogeneous mixture, the same all the way through, while a heterogeneous mixture varies from place to place. You'll mostly describe concentration with molarity, M = moles of solute ÷ liters of solution, and use it to find moles, volume or number of particles.
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A particle diagram of a solution should show the right ratio of particles and how they interact, such as water molecules pointing their partly negative oxygen toward cations and their partly positive hydrogens toward anions. For example, a drawing of CaCl₂ dissolved in water should show twice as many Cl⁻ ions as Ca²⁺ ions.
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You can't filter the parts of a solution apart, but you can separate them using differences in intermolecular forces. In chromatography, a component that is more strongly attracted to the stationary phase moves more slowly and travels a shorter distance; in distillation, the component with weaker intermolecular forces has a higher vapor pressure and boils off first.
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Substances with similar intermolecular forces tend to mix: ionic and polar substances generally dissolve in polar solvents like water, and nonpolar substances dissolve in nonpolar solvents ('like dissolves like'). Oil and water don't mix because water molecules attract each other through hydrogen bonds far more than they attract nonpolar oil molecules.
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Different kinds of light cause different changes in molecules: microwaves make molecules rotate, infrared light makes bonds vibrate, and ultraviolet or visible light moves electrons to higher energy levels. That is why each region of the spectrum tells chemists something different about a substance.
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Light comes in packets of energy called photons. Wavelength and frequency are linked by c = λν, and a photon's energy is E = hν, so shorter wavelengths carry more energy; when an atom or molecule absorbs or emits a photon, its energy goes up or down by exactly the photon's energy.
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The Beer-Lambert law, A = εbc, says a solution's absorbance depends on how strongly the substance absorbs that wavelength (molar absorptivity, ε), the path length (b) and the concentration (c). With path length and wavelength held constant, absorbance is directly proportional to concentration, so a graph of absorbance against known concentrations lets you find an unknown one.
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