Difficulty: Intermediate to Advanced | Prerequisites: Chapters 8-9 (bonding, polarity), Chapter 11-12 (mole, stoichiometry)
Source: Chem 1 Final Study Guide, General Chemistry (Purdue University, Chhabra)
Tags: intermolecular forces, dispersion forces, dipole-dipole, hydrogen bonding, viscosity, surface tension, phase changes, phase diagram, triple point, critical point, Boyle's law, gas laws, Dalton's law, solutions, molarity, molality, colligative properties, boiling point elevation, freezing point depression, Tyndall effect, calorimetry, specific heat, enthalpy, Hess's law, entropy, Gibbs free energy, spontaneity
These chapters move from the forces between molecules (which determine whether something is a solid, liquid, or gas) through the quantitative behaviour of gases and solutions, and finally into the energy changes that accompany chemical reactions. Chapter 13 covers intermolecular forces and phase changes. Chapter 14 is the gas laws. Chapter 15 deals with solutions and their properties. Chapter 16 introduces thermochemistry, including enthalpy, Hess's law, entropy, and Gibbs free energy. Together, these chapters teach you to predict physical properties, calculate concentrations, and determine whether a reaction will release or absorb energy.
Intermolecular forces (dispersion, dipole-dipole, hydrogen bonds) control whether a substance is solid, liquid, or gas, plus properties like viscosity and surface tension. Gas behaviour follows Boyle's law and other relationships. Solutions have colligative properties that depend on solute particle count, not identity. Energy changes in reactions are measured as enthalpy, and spontaneity depends on the balance between enthalpy and entropy (Gibbs free energy).
Intermolecular forces (IMFs)
Forces of attraction between molecules. Weaker than intramolecular (bonding) forces. The three types are dispersion forces, dipole-dipole forces, and hydrogen bonds.
Dispersion forces (London forces)
Weak, temporary attractions caused by momentary shifts in electron density. Present in all molecules. Dominant in nonpolar molecules. Strength increases with more electrons (larger molecules).
Dipole-dipole forces
Attractions between the permanent partial charges on polar molecules. Stronger than dispersion forces for molecules of similar size.
Hydrogen bond
A particularly strong dipole-dipole force that occurs when hydrogen is bonded to fluorine, oxygen, or nitrogen (small, highly electronegative atoms with lone pairs). This is why water has unusually high boiling point and surface tension.
Viscosity
A measure of a liquid's resistance to flow. Higher intermolecular forces, larger particles, and lower temperatures all increase viscosity.
Surface tension
The energy required to increase the surface area of a liquid. Caused by the inward pull on surface molecules. Water's high surface tension comes from its extensive hydrogen bonding.
Cohesion
Attraction between identical molecules (water to water).
Adhesion
Attraction between different molecules (water to glass).
Capillary action
The ability of a liquid to flow in narrow spaces against gravity, caused by adhesion being stronger than cohesion. Creates the concave meniscus in a glass tube.
Crystalline solid
A solid with atoms, ions, or molecules arranged in an orderly, repeating 3D structure.
Amorphous solid
A solid without a regular repeating pattern (e.g. glass, rubber).
Melting point
The temperature at which the forces holding a crystal lattice together are overcome, converting solid to liquid.
Boiling point
The temperature at which a liquid's vapour pressure equals the external (atmospheric) pressure.
Vapour pressure
The pressure exerted by a vapour above its liquid in a closed container. Higher temperature = higher vapour pressure.
Sublimation
The direct change from solid to gas, skipping the liquid phase.
Deposition
The direct change from gas to solid, skipping the liquid phase.
Phase diagram
A graph of pressure vs temperature showing the conditions under which each phase (solid, liquid, gas) is stable.
Triple point
The specific temperature and pressure where all three phases coexist in equilibrium.
Critical point
The temperature and pressure above which a distinct liquid phase cannot exist.
Kinetic-molecular theory
A model describing gas behaviour: gas particles are tiny, far apart, in constant random motion, and their collisions are perfectly elastic (no kinetic energy lost).
Elastic collision
A collision in which kinetic energy may transfer between particles, but the total kinetic energy is conserved.
Diffusion
The movement of one substance through another. Lighter particles diffuse faster.
Effusion
The escape of gas particles through a tiny opening.
Graham's law of effusion
Rate A / Rate B = √(molar mass B / molar mass A). Lighter gases effuse faster.
Boyle's law
At constant temperature, pressure and volume are inversely proportional: P₁V₁ = P₂V₂.
Dalton's law of partial pressures
The total pressure of a gas mixture equals the sum of the partial pressures of each gas.
Heat of solution
The energy change during the solvation process.
Saturated solution
A solution at dynamic equilibrium where the rate of dissolving equals the rate of crystallisation.
Supersaturated solution
A solution containing more dissolved solute than a saturated solution at that temperature. Adding a seed crystal or disturbing it can trigger rapid crystallisation.
Henry's law
S₁/P₁ = S₂/P₂. The solubility of a gas in a liquid is directly proportional to the pressure of that gas above the liquid.
Molarity (M)
Moles of solute per litre of solution.
Molality (m)
Moles of solute per kilogram of solvent.
Dilution equation
M₁V₁ = M₂V₂. Used when diluting a solution to a lower concentration.
Colligative properties
Properties that depend on the number of dissolved solute particles, not their identity: vapour pressure lowering, boiling point elevation, freezing point depression, and osmotic pressure.
Nonvolatile solute
A solute with no tendency to evaporate. Adding it to a solvent lowers the solvent's vapour pressure.
Osmotic pressure
The extra pressure caused by water molecules moving across a semi-permeable membrane into a more concentrated solution.
Tyndall effect
The scattering of light by colloidal particles. Solutions show no Tyndall effect; colloids and suspensions do.
Specific heat (c)
The amount of heat required to raise 1 gram of a substance by 1°C. Water's specific heat is 4.184 J/g°C.
Calorimeter
An insulated device used to measure heat absorbed or released during a chemical or physical process.
Enthalpy (H)
The heat content of a system at constant pressure. You cannot measure absolute enthalpy, only the change (ΔH).
Enthalpy of reaction (ΔH_rxn)
ΔH = H(products) – H(reactants). Negative = exothermic; positive = endothermic.
Exothermic reaction
Releases heat to the surroundings. ΔH < 0.
Endothermic reaction
Absorbs heat from the surroundings. ΔH > 0.
Thermochemical equation
A balanced equation that includes the physical states of all substances and the enthalpy change.
Hess's law
If a reaction can be expressed as the sum of two or more simpler reactions, ΔH for the overall reaction equals the sum of the ΔH values for those steps. Used when direct measurement is not feasible.
Standard enthalpy of formation (ΔH°f)
The enthalpy change for forming one mole of a compound from its elements in their standard states. Elements in their standard state have ΔH°f = 0.
Standard state
298 K (25°C) and 1 atm. Different from STP (0°C and 1 atm).
Entropy (S)
A measure of the disorder or randomness of particles in a system. Higher entropy = more disorder.
Gibbs free energy (ΔG)
ΔG = ΔH – TΔS. Determines spontaneity. ΔG < 0: spontaneous. ΔG > 0: nonspontaneous.
Intermolecular forces (weakest to strongest for molecules of similar size)
Dispersion forces: present in all molecules; dominant in nonpolar substances; strength grows with molecular size
Dipole-dipole forces: between polar molecules; stronger than dispersion at similar molecular weight
Hydrogen bonds: between molecules where H is bonded to F, O, or N; strongest IMF type among small molecules
Liquid properties
Density: much greater than gases, compressible to a very limited extent
Viscosity: resistance to flow; increases with stronger IMFs, larger molecules, and lower temperature
Surface tension: inward pull from interior molecules; water is high because of hydrogen bonding
Capillary action: adhesion > cohesion = liquid climbs (concave meniscus for water in glass)
Water's unusual properties
Maximum density at 4°C; ice is less dense than liquid water (ice floats)
This is why lakes freeze from the top down, keeping aquatic life alive underneath
Solids
Crystalline: ordered 3D structure (ionic crystals, metallic, molecular, covalent network)
Amorphous: disordered (glass, rubber)
Molecular solids are held together by IMFs (dispersion, dipole-dipole, hydrogen bonds)
Phase changes
Melting/freezing, vaporisation/condensation, sublimation/deposition
Evaporation: vaporisation at the surface only
Boiling: vaporisation throughout the liquid, occurs when vapour pressure = external pressure
Phase diagrams
Triple point: all three phases coexist; all six phase changes can occur
Critical point: above this temperature and pressure, liquid and gas become indistinguishable
Kinetic-molecular theory assumptions
Gas particles do not attract or repel each other
Particles are much smaller than the distances between them
Constant, random motion; elastic collisions
All gases have the same average KE at a given temperature
KE = ½mv²
Key relationships
Boyle's law: P₁V₁ = P₂V₂ (T constant, inverse relationship)
Dalton's law: P(total) = P₁ + P₂ + P₃ + ...
Graham's law
Rate A / Rate B = √(molar mass B / molar mass A)
Lighter gases effuse and diffuse faster
Solubility and saturation
Saturated: dynamic equilibrium between dissolving and crystallising
Supersaturated: excess solute remains dissolved; a seed crystal or disturbance triggers crystallisation
Henry's law (gases): S₁/P₁ = S₂/P₂ (higher pressure = more gas dissolved)
Concentration units
Molarity (M) = moles solute / litres solution
Molality (m) = moles solute / kg solvent
Dilution: M₁V₁ = M₂V₂
Colligative properties
Depend on the number of dissolved particles, not their chemical identity
Adding a nonvolatile solute lowers vapour pressure (solute particles block surface evaporation)
Boiling point elevation: solute raises the boiling point
Freezing point depression: solute lowers the freezing point
Osmotic pressure: water moves toward higher solute concentration through a membrane
Electrolytes dissociate into ions, producing more particles per formula unit than nonelectrolytes
Mixtures, colloids, and suspensions
Suspensions: settle out, filterable, show Tyndall effect
Thixotropic: solid-like when settled, flows when stirred (e.g. paint)
Colloids: intermediate particle size, do not settle, show Tyndall effect
Stabilised by Brownian motion and electrostatic layering
Heating or adding electrolyte can break the colloid
Solutions: smallest particles, no settling, no Tyndall effect
Energy basics
Energy = ability to do work or produce heat
Chemical potential energy: energy stored in a substance due to its composition
Law of conservation of energy: energy is neither created nor destroyed
Heat (q) flows from hot to cold
Units and specific heat
1 calorie = heat to raise 1 g of water by 1°C; 1 cal = 4.184 J
Specific heat of water: 4.184 J/g°C
q = c x m x ΔT (heat = specific heat x mass x temperature change)
Calorimetry
Measure q by tracking temperature change in a known mass of water
q(water) = –q(substance) in a closed calorimeter
System = the reaction; surroundings = everything else
Enthalpy
ΔH(rxn) = H(products) – H(reactants)
Exothermic: ΔH < 0 (energy appears on the products side)
Endothermic: ΔH > 0 (energy appears on the reactants side)
ΔH(vap) = –ΔH(cond); ΔH(fusion) = –ΔH(solidification)
Hess's law
ΔH for a reaction is the same regardless of the pathway, as long as the initial and final states are the same
Sum the ΔH values of individual steps (reversing a reaction flips the sign; multiplying the equation multiplies ΔH)
Standard enthalpy of formation (ΔH°f)
Enthalpy change for forming 1 mol of a compound from its elements in their standard states
Free elements in standard state: ΔH°f = 0
ΔH°(rxn) = Σ ΔH°f(products) – Σ ΔH°f(reactants)
Entropy (S)
Measure of disorder/randomness
Entropy increases: solid → liquid → gas; dissolving a solid in a solvent; increasing number of gas particles; increasing temperature
Entropy decreases: dissolving a gas in a liquid or solid solvent
Gibbs free energy
ΔG = ΔH – TΔS (T in kelvin)
ΔG < 0: spontaneous
ΔG > 0: nonspontaneous
Spontaneity depends on both ΔH and ΔS
–ΔH and +ΔS: always spontaneous
+ΔH and –ΔS: never spontaneous
Mixed signs: spontaneity depends on temperature
Important correction from the source material
The claim that "all exothermic processes are spontaneous and all endothermic processes are nonspontaneous" is false. Ice melting at room temperature is a spontaneous endothermic process.
q = cmΔT (heat = specific heat x mass x temperature change)
Boyle's law: P₁V₁ = P₂V₂
Graham's law: Rate A / Rate B = √(molar mass B / molar mass A)
Dalton's law: P(total) = P₁ + P₂ + ...
Henry's law: S₁/P₁ = S₂/P₂
Dilution: M₁V₁ = M₂V₂
Percent yield: (actual / theoretical) x 100
Gibbs free energy: ΔG = ΔH – TΔS
Hess's law: ΔH(rxn) = Σ ΔH°f(products) – Σ ΔH°f(reactants)
Hydrogen bonding in water is responsible for ice floating, which insulates lakes in winter and keeps aquatic ecosystems alive. This is also why water has an unusually high specific heat, making it an excellent coolant in car radiators and industrial processes.
Colligative properties explain why salt is spread on icy roads (freezing point depression) and why antifreeze works in car engines (both boiling point elevation and freezing point depression).
Hess's law is used industrially to calculate energy requirements for reactions that cannot be safely run in a calorimeter.
Students often assume hydrogen bonding involves a bond to hydrogen in general. It specifically requires hydrogen bonded to F, O, or N. H bonded to carbon, for example, does not form hydrogen bonds.
A frequent error: thinking that exothermic = spontaneous and endothermic = nonspontaneous. Spontaneity depends on ΔG, which accounts for both enthalpy and entropy. Ice melts spontaneously at room temperature despite being endothermic.
Students sometimes confuse molarity and molality. Molarity uses litres of solution; molality uses kilograms of solvent. They are not interchangeable.
In Boyle's law problems, students occasionally forget that temperature must be held constant for P₁V₁ = P₂V₂ to apply.
⚠️ Identifying the type and relative strength of IMFs for a given molecule is a common exam question. Practise classifying molecules by their dominant intermolecular force.
⚠️ Phase diagram questions (identifying the triple point, reading which phase exists at given conditions) appear frequently.
⚠️ q = cmΔT calculations are tested both on their own and embedded in calorimetry problems.
⚠️ Gibbs free energy: know the sign conventions and be able to determine spontaneity from ΔH and ΔS at a given temperature.
⚠️ Hess's law problems (summing thermochemical equations) are a staple of the final. Practise reversing equations and adjusting ΔH.
⚠️ Know the difference between colligative and non-colligative properties, and why electrolytes produce a greater effect than nonelectrolytes.
True or False: Hydrogen bonds can form between any molecule that contains hydrogen.
Fill in the blank: At 4°C, liquid water has its __________ density.
True or False: According to Boyle's law, doubling the pressure on a gas at constant temperature will double its volume.
Fill in the blank: A reaction with ΔG < 0 is __________.
True or False: Molality uses litres of solution as its denominator.
Q: Rank the following in order of increasing boiling point: CH₄, H₂O, HCl.
A: CH₄ < HCl < H₂O. CH₄ is nonpolar (dispersion only). HCl is polar (dipole-dipole). H₂O has hydrogen bonding, the strongest IMF here.
Q: A gas occupies 2.0 L at 1.0 atm. What volume will it occupy at 4.0 atm (constant temperature)?
A: P₁V₁ = P₂V₂. (1.0)(2.0) = (4.0)(V₂). V₂ = 0.50 L.
Q: Calculate the heat absorbed when 50.0 g of water is heated from 25.0°C to 75.0°C.
A: q = cmΔT = (4.184)(50.0)(50.0) = 10,460 J = 10.5 kJ.
Q: Is a reaction with ΔH = +100 kJ and ΔS = +200 J/K spontaneous at 600 K?
A: ΔG = ΔH – TΔS = 100,000 – (600)(200) = 100,000 – 120,000 = –20,000 J = –20 kJ. Yes, spontaneous at 600 K.
Q: What is the molarity of a solution made by dissolving 5.85 g of NaCl (molar mass 58.44 g/mol) in enough water to make 500.0 mL of solution?
A: Moles NaCl = 5.85 / 58.44 = 0.1001 mol. M = 0.1001 / 0.5000 = 0.200 M.
Intermolecular forces determine solubility, which is central to solution chemistry (Ch. 15), acids and bases (Ch. 19), and equilibrium (Ch. 18).
Thermochemistry connects directly to equilibrium (Le Chatelier's principle uses ΔH to predict shifts) and electrochemistry (cell potentials relate to ΔG).
Gas laws reappear in stoichiometry problems involving gaseous reactants or products.
intermolecular forces, dispersion forces London forces, dipole-dipole forces, hydrogen bonding, viscosity, surface tension, cohesion adhesion, capillary action meniscus, crystalline solid amorphous, phase changes melting boiling, sublimation deposition, vapour pressure, evaporation, phase diagram triple point critical point, kinetic molecular theory, elastic collision, diffusion effusion, Graham's law, pressure, Boyle's law P1V1 P2V2, Dalton's law partial pressures, solution saturated supersaturated, Henry's law gas solubility pressure, molarity molality, M1V1 M2V2 dilution, colligative properties, boiling point elevation, freezing point depression, osmotic pressure, electrolyte nonelectrolyte, nonvolatile solute, Tyndall effect, colloid suspension, Brownian motion, thixotropic, specific heat, calorimeter, enthalpy, exothermic endothermic, thermochemical equation, Hess's law, standard enthalpy of formation, entropy disorder, Gibbs free energy spontaneity, coupled reactions