South 고등학교 3학년 Chemistry II
Chapters: 4
1. States of matter and solutions
Gas variables · Ideal gas equation · Mole fractions and partial pressure · Intermolecular forces · Special properties of water · Vapour pressure and boiling · Types of solids · Concentration units · Colligative properties · Osmosis
- Perfect Gas Equation PV = nRT and Work in Compressing a Gas – A gas is made of countless tiny molecules flying about. Their hits on the walls make pressure. For a low-density gas, three simple laws hold: at fixed temperature, P × V stays constant (Boyle); at fixed pressure, V grows in step with kelvin temperature (Charles); at the same P and T, equal volumes hold equal numbers of molecules (Avogadro). Put together they give the perfect gas equation PV = nRT = N k T. One mole holds Avogadro's number, 6.022 × 10²³, of particles. Pushing a piston in does work on the gas; the work equals the area under the P–V graph, and at fixed temperature W = nRT ln(V₁/V₂).
- Gas Laws: How Pressure, Volume and Temperature Are Linked – Gas pressure comes from particles hitting the walls. For a fixed amount of gas: Boyle's law P₁V₁ = P₂V₂ (constant T); Charles's law V₁/T₁ = V₂/T₂ (constant P); pressure (Gay-Lussac's) law P₁/T₁ = P₂/T₂ (constant V). Temperature must be in kelvin. Combined: P₁V₁/T₁ = P₂V₂/T₂. Dalton's law: in a mixture, total pressure = sum of partial pressures, and each partial pressure = mole fraction × total pressure.
- Intermolecular Forces – Intermolecular forces are weak attractions BETWEEN molecules. From weakest to strongest (for similar-size molecules): London dispersion < dipole–dipole < hydrogen bond. Stronger forces mean higher melting and boiling points, higher viscosity and surface tension, and they decide what dissolves in what.
- Water: The Small Molecule That Makes Life Possible – A water molecule (H₂O) is bent and polar: oxygen is slightly negative and the hydrogens slightly positive. Because of this, water molecules stick together with hydrogen bonds. These bonds explain why water dissolves many substances, stores a lot of heat, takes a lot of energy to evaporate, sticks together (cohesion) and to surfaces (adhesion), and why ice floats. Living things also use water as a reactant, for transport and to keep cells firm. Mineral salts dissolved in water supply ions like Na⁺, K⁺, Ca²⁺, Fe²⁺ and phosphate.
- Phase Changes: Latent Heat, Vapour Pressure and Phase Diagrams – A phase change is when matter moves between solid, liquid and gas. During a change the temperature stays the same; the heat used is latent heat (Q = m × L). A liquid boils when its vapour pressure equals the pressure above it, so boiling point falls with lower pressure. A phase diagram maps the state at every temperature and pressure; its lines meet at the triple point and the liquid–gas line ends at the critical point.
- The Solid State – In a solid the particles sit close together and only vibrate in place, so a solid keeps its shape. In crystalline solids the particles repeat in a regular pattern; in amorphous solids (like glass) they do not. The smallest repeating box is the unit cell. Simple cubic holds 1 particle, body-centred cubic 2 and face-centred cubic 4. By the bonds holding them, crystals are ionic, covalent network, molecular or metallic.
- Solutions: Types, Concentration and Henry's Law – A solution is an even mix of two or more substances. The part in bigger amount is the solvent, the smaller part is the solute. We tell 'how strong' a solution is with concentration terms: mass %, volume %, ppm, mole fraction, molarity (per litre of solution) and molality (per kg of solvent). Solids usually dissolve more when hot. Gases dissolve more when their pressure is high (Henry's law, p = KH·x) and less when it is hot.
- Colligative Properties: Counting Particles, Not Their Type – Colligative properties depend only on HOW MANY solute particles are in the solution, not on what they are. There are four: relative lowering of vapour pressure (Δp/p° = x₂), elevation of boiling point (ΔTb = Kb·m), depression of freezing point (ΔTf = Kf·m) and osmotic pressure (π = CRT). We use them to find molar mass. Salts that split into ions (NaCl) or molecules that pair up (acetic acid in benzene) give abnormal molar masses, fixed with the van 't Hoff factor i.
2. Reaction enthalpy and equilibrium
Thermochemical equations · Bond energies and Hess's law · Equilibrium constant · Le Chatelier · Phase diagrams · Ionisation constants · Buffers
- Enthalpy, Calorimetry and Hess's Law – Enthalpy H = U + pV. At constant pressure the heat taken in or given out equals ΔH; at constant volume it equals ΔU. We measure ΔU in a sealed bomb calorimeter and ΔH in an open cup calorimeter, using q = C ΔT or q = m c ΔT. For gases, ΔH = ΔU + Δn_g RT. The reaction enthalpy ΔrH is negative for exothermic and positive for endothermic reactions, and ΔrH° = ΣΔfH°(products) − ΣΔfH°(reactants). Hess's law: the total enthalpy change is the same whether a reaction happens in one step or many, so thermochemical equations can be added like algebra.
- Chemical Equilibrium: Kc, Kp, Q and Gibbs Energy – In a closed container a reversible reaction goes both ways. After some time the forward and backward rates become equal, so amounts stop changing, but the reaction does not stop. This is dynamic equilibrium. At equilibrium the ratio of products to reactants (each raised to its coefficient) is a fixed number, the equilibrium constant K. Kc uses concentrations, Kp uses partial pressures, and Kp = Kc(RT)^Δn. Pure solids and liquids are left out of K. The reaction quotient Q tells the direction: Q < K goes forward, Q > K goes backward, Q = K is equilibrium. K and Gibbs energy are linked: ΔG = ΔG° + RT ln Q and ΔG° = −RT ln K.
- Le Chatelier's Principle: Factors Affecting Equilibrium – Le Chatelier's principle says: if you disturb a system at equilibrium, it shifts in the direction that reduces the disturbance. Add a reactant or remove a product → shifts forward. Increase pressure (smaller volume) → shifts to the side with fewer gas moles. Raise temperature → shifts in the heat-absorbing (endothermic) direction, and K changes. A catalyst only helps reach equilibrium faster; it does not move it. Inert gas at constant volume does nothing; at constant pressure it acts like lowering pressure. Only temperature changes the value of K.
- Phase Changes: Latent Heat, Vapour Pressure and Phase Diagrams – A phase change is when matter moves between solid, liquid and gas. During a change the temperature stays the same; the heat used is latent heat (Q = m × L). A liquid boils when its vapour pressure equals the pressure above it, so boiling point falls with lower pressure. A phase diagram maps the state at every temperature and pressure; its lines meet at the triple point and the liquid–gas line ends at the critical point.
- Ionic Equilibrium: Acids, Bases, Ionisation and pH – Acids and bases can be defined in three ways: Arrhenius (give H⁺ or OH⁻ in water), Brønsted–Lowry (proton donor or acceptor) and Lewis (electron-pair acceptor or donor). Strong acids and bases ionise fully; weak ones ionise only a little and set up an equilibrium with constant Ka or Kb. Water itself ionises: Kw = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴ at 25 °C. pH = −log[H⁺] and pH + pOH = 14. For a weak electrolyte, α ≈ √(K/C), so dilution increases α (Ostwald's dilution law). For a conjugate pair, Ka × Kb = Kw. Salts of weak acids or weak bases react with water (hydrolysis) and give non-neutral solutions.
- Buffer Solutions and Solubility Product – A buffer is a solution that keeps its pH nearly the same when a little acid or base is added. An acidic buffer is a weak acid with its salt (CH₃COOH + CH₃COONa); a basic buffer is a weak base with its salt (NH₄OH + NH₄Cl). Its pH is given by the Henderson equation: pH = pKa + log([salt]/[acid]). A sparingly soluble salt in its saturated solution sets up an equilibrium with its ions; the product of ion concentrations (each raised to its coefficient) is the solubility product Ksp. For AB, Ksp = s²; for AB₂ or A₂B, Ksp = 4s³. If the ionic product Q exceeds Ksp, a precipitate forms. A common ion lowers solubility.
3. Reaction rates and catalysts
Rates vary · Rate laws from data · First-order half-life · Activation energy · Concentration effects · Temperature effects · Catalysts · Catalysts in life and industry
- Rate of a Chemical Reaction – The rate of a reaction tells how fast a reactant is used up or a product is made, per unit time. Rate = −Δ[R]/Δt = +Δ[P]/Δt (unit mol L⁻¹ s⁻¹). The rate law, rate = k[A]^x[B]^y, is found by experiment; x + y is the order. Molecularity is the number of particles that collide in one elementary step.
- Integrated Rate Equations: Zero and First Order – Integrated rate equations link concentration with time. Zero order: [R] = [R]₀ − kt (straight line), t½ = [R]₀/2k. First order: k = (2.303/t) log([R]₀/[R]), ln[R] falls in a straight line with slope −k, and t½ = 0.693/k, which does not depend on the starting amount.
- Temperature, Activation Energy, Catalysts and Collision Theory – Reactant molecules must climb an energy hill, the activation energy Ea, to become products. The Arrhenius equation k = A e^(−Ea/RT) shows that a small rise in temperature lets many more molecules cross, so k rises fast. A catalyst gives a lower hill without changing ΔH. Collision theory: molecules react only when they collide with enough energy and the right orientation.
- Catalysts: Speeding Up Reactions Without Being Used Up – A catalyst is a substance that makes a chemical reaction faster but is not used up; its mass and chemical nature are the same at the end. It works by giving the reaction a different route with a lower activation energy. Manganese dioxide speeds up the breakdown of hydrogen peroxide into water and oxygen. Enzymes are catalysts in living things.
- Enzymes – Enzymes are proteins that act as biological catalysts: they speed up reactions in cells without being used up. Each enzyme has an active site that fits only one kind of substrate (specificity). Enzymes work by lowering the activation energy. They work best at an optimum temperature and pH; too much heat or the wrong pH changes the active site's shape and denatures them.
4. Electrochemistry
Galvanic cells · Electrolysis · Hydrogen fuel cells
- Galvanic Cells and the Nernst Equation – A galvanic cell turns the energy of a redox reaction into electricity. Oxidation happens at the anode (−) and reduction at the cathode (+). Each electrode has a potential measured against the standard hydrogen electrode (0 V). E°cell = E°cathode − E°anode. The Nernst equation, E = E° − (0.059/n) log Q at 298 K, gives the cell voltage at any concentration, and ΔG = −nFE links voltage to energy.
- Electrolysis, Batteries, Fuel Cells and Corrosion – In an electrolytic cell an outside source of electricity forces a non-spontaneous reaction: cations are reduced at the cathode (−) and anions oxidised at the anode (+). Faraday's laws link the mass changed to the charge: m = (M/nF) × I × t. Which product forms depends on electrode potentials and overpotential. Batteries are galvanic cells: primary (dry cell, mercury cell) cannot be recharged; secondary (lead storage, Ni–Cd) can. Fuel cells burn H₂ with O₂ to give electricity directly. Corrosion (rusting) is an unwanted galvanic cell on the metal surface.