China 高二 Chemistry
Chapters: 7
1. Selective 1 Ch.1 Thermal effects
Reaction heat; enthalpy · Calculating reaction heat; Hess's law
- 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.
2. Selective 1 Ch.2 Rate and equilibrium
Reaction rate · Chemical equilibrium; K · Direction of reactions (enthalpy, entropy) · Controlling reactions
- 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.
- 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.
- Spontaneity, Entropy and Gibbs Energy – A spontaneous process happens by itself, without outside help (it may be fast or slow). ΔH alone cannot predict it: ice melts on its own although it takes in heat. Entropy S measures how spread out energy and matter are; ΔS = q_rev/T. Second law: for any spontaneous change, ΔS_total = ΔS_system + ΔS_surroundings > 0; at equilibrium it is zero. Third law: a perfect crystal at 0 K has zero entropy. Gibbs energy G = H − TS combines both: at constant T and p, ΔG = ΔH − TΔS, and ΔG < 0 means spontaneous, ΔG = 0 means equilibrium. It links to the equilibrium constant by ΔG° = −RT ln K = −2.303 RT log 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.
3. Selective 1 Ch.3 Ionic equilibria
Ionisation equilibrium · Water ionisation; pH · Salt hydrolysis · Precipitation–dissolution equilibrium
- 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.
4. Selective 1 Ch.4 Reactions and electricity
Galvanic cells; batteries · Electrolytic cells · Corrosion and protection
- 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.
- Corrosion, Its Prevention and Alloys – Corrosion is the slow eating away of a metal by air, moisture or chemicals around it. Iron rusts only when both air (oxygen) and water are present, forming hydrated iron oxide. It can be prevented by painting, oiling, galvanising, chrome plating, anodising or making alloys. An alloy is a uniform mixture of a metal with other metals or a non-metal; alloys are usually harder, stronger and more corrosion-resistant than the pure metal.
5. Selective 2 Ch.1 Atomic structure and properties
Energy levels, orbitals, configuration rules · Radius, ionisation energy, electronegativity
- Quantum Mechanical Model: Quantum Numbers, Orbitals and Configuration – Moving electrons also behave like waves: λ = h/mv (de Broglie). Because of this we cannot know exact position and momentum together: Δx·Δp ≥ h/4π (Heisenberg). So we drop sharp orbits and use orbitals: regions where the electron is likely to be, from the Schrödinger equation (ψ² gives probability). Four quantum numbers describe each electron: n (shell, size, energy), l (subshell, shape: 0 to n−1), mₗ (orientation: −l to +l) and mₛ (spin: +½ or −½). s is spherical, p is dumbbell, d is mostly four-lobed. Electrons fill by the Aufbau (n + l) rule, Pauli principle (max 2, opposite spins) and Hund's rule (single first). Half-filled and fully filled subshells are extra stable, so Cr is 3d⁵4s¹ and Cu is 3d¹⁰4s¹.
- Periodic Trends in Properties – Two forces decide almost every trend: the pull of the nucleus (effective nuclear charge) and the distance of the outer shell. Across a period the nuclear pull grows while the shell stays the same, so atoms shrink, ionisation enthalpy rises, electron gain enthalpy becomes more negative and electronegativity rises. Down a group a new shell is added, so atoms grow and these values fall. Cations are smaller and anions bigger than their atoms. Valence follows the outer electrons; metallic reactivity is highest at the bottom left and non-metallic reactivity at the top right.
6. Selective 2 Ch.2 Molecular structure
Covalent bonds (σ, π) · Molecular shapes; VSEPR; hybridisation · Polarity; intermolecular forces; hydrogen bonds
- Valence Bond Theory and Hybridisation – Valence bond theory says a covalent bond forms when half-filled orbitals of two atoms overlap and the two electrons pair up with opposite spins. Head-on overlap makes a strong σ (sigma) bond; side-on overlap of p orbitals makes a weaker π (pi) bond. To explain real shapes, an atom first mixes its orbitals into equal hybrid orbitals: sp (linear, 180°), sp² (trigonal planar, 120°), sp³ (tetrahedral, 109.5°), sp³d (trigonal bipyramidal) and sp³d² (octahedral).
- VSEPR Theory: Shapes of Molecules – VSEPR stands for Valence Shell Electron Pair Repulsion. The electron pairs around a central atom repel each other and move as far apart as possible. The number of pairs sets the basic arrangement: 2 linear, 3 trigonal planar, 4 tetrahedral, 5 trigonal bipyramidal, 6 octahedral. Lone pairs take more room than bond pairs (lp–lp > lp–bp > bp–bp), so they squeeze bond angles and change the shape we see, for example NH₃ (pyramidal, 107°) and H₂O (bent, 104.5°).
- Hydrogen Bonding – A hydrogen bond is a weak attraction between an H atom that is joined to a small, very electronegative atom (F, O or N) and a lone pair on another such atom. It is shown with a dotted line: X–H···Y. It is much weaker than a covalent bond (about 10–40 kJ mol⁻¹) but strong enough to raise boiling points, make ice float and hold DNA together. Intermolecular H-bonds join different molecules; intramolecular H-bonds form inside one molecule.
7. Selective 2 Ch.3 Crystals
States of aggregation; unit cells · Molecular and covalent crystals · Metallic and ionic crystals · Complexes and supramolecules
- 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.
- Coordination Compounds: Werner's Theory, Terms and IUPAC Names – A coordination compound has a central metal atom or ion joined to a fixed number of ions or molecules called ligands. Each ligand gives one pair of electrons to the metal. The metal and its ligands sit together inside square brackets, the coordination sphere. The number of donor atoms bonded to the metal is the coordination number. Werner said a metal shows two kinds of valence: primary (ionisable, outside the bracket) and secondary (fixed, inside the bracket).