Russia 11 класс Chemistry (advanced)

1. Theoretical foundations of chemistry

Atomic structure and quantum numbers · Periodic law · Chemical bonding and complex compounds · Solutions · Reactions and thermochemistry · Kinetics and catalysis · Equilibrium · Electrolytes, pH, hydrolysis · Redox and electrolysis

• 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.
• Kossel-Lewis Approach and the Ionic Bond – Atoms join so that each gets a stable outer shell of 8 electrons (an octet), like a noble gas. Kossel said atoms can give or take electrons to make ions (ionic bond). Lewis said atoms can also share pairs of electrons (covalent bond). Lewis structures show these electrons as dots and lines. Formal charge (V − L − B/2) helps pick the best Lewis structure. Ions pack into a crystal, and the energy released is linked to the lattice enthalpy.
• 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.
• 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.
• 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.
• 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.
• 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.

2. Non-metals

Hydrogen · Halogens · Oxygen and ozone · Sulfur · Nitrogen · Phosphorus · Carbon and silicon

• Hydrogen: The Lightest Element – Hydrogen (H) is element 1: one proton and one electron. It exists as H₂ molecules, a colourless, odourless gas that is the least dense of all gases and hardly dissolves in water. In the lab it is made from a reactive metal and a dilute acid (Zn + 2HCl → ZnCl₂ + H₂) and collected over water. It burns with a squeaky pop to form only water. It reacts with non-metals to give hydrides such as HCl, NH₃ and H₂S, and it reduces some metal oxides. It has three isotopes: protium, deuterium and tritium. Uses: ammonia and fertilisers, margarine, rocket fuel and clean fuel cells.
• Halogens: The Group 17 Elements – The halogens are fluorine, chlorine, bromine, iodine and astatine, in group 17 of the periodic table. Each atom has 7 outer electrons, so it gains 1 electron to form a halide ion (X⁻). Down the group, atoms get bigger, colours get darker, melting and boiling points rise, and reactivity falls. A more reactive halogen displaces a less reactive one from its salt.
• Oxygen – Oxygen (O, atomic number 8) is a colourless, odourless gas that makes up about 21% of air by volume. It exists as O₂ molecules. In the lab it is made by breaking down hydrogen peroxide with a manganese dioxide catalyst (2H₂O₂ → 2H₂O + O₂) or by heating potassium manganate(VII); in industry it comes from fractional distillation of liquid air. It relights a glowing splint. Oxygen supports burning and respiration and reacts with most elements to form oxides: metal oxides are basic, non-metal oxides are mostly acidic. Its other form, ozone (O₃), shields Earth from UV rays.
• Sulfur and Its Compounds – Sulfur (S, atomic number 16, electrons 2, 8, 6) is a yellow non-metal in group 16 with oxygen. Solid sulfur is made of S₈ rings and has rhombic and monoclinic forms. It burns to sulfur dioxide (SO₂), a gas that bleaches, kills microbes and causes acid rain. In the contact process SO₂ is oxidised to SO₃ over a V₂O₅ catalyst and turned into sulfuric acid (H₂SO₄), a strong acid that also dehydrates and oxidises. Hydrogen sulfide (H₂S) smells of bad eggs. Sulfate ions are found with barium chloride: a white precipitate of BaSO₄ that does not dissolve in dilute HCl.
• Nitrogen: the Element, Its Compounds and the Nitrogen Cycle – Nitrogen (N₂) makes up 78% of air, but its triple bond is so strong that most living things cannot use it. Fixation (by root-nodule bacteria, lightning and the Haber process) turns it into ammonia. From ammonia we make nitric acid and fertilisers. In nature, nitrogen moves in a cycle: fixation → nitrification → absorption by plants → animals → decay (ammonification) → denitrification back to air.
• Carbon and Its Compounds: Bonding, Hydrocarbons and Naming – Carbon has 4 outer electrons, so it shares electrons (covalent bonds) instead of gaining or losing them. Because it bonds to itself (catenation) and always makes 4 bonds (tetravalency), it forms millions of compounds: chains, branches and rings, saturated or unsaturated. Compounds with the same functional group form a homologous series that differs by –CH₂–, and IUPAC names are built from the number of carbons + a suffix or prefix for the functional group.

3. Metals

General properties of metals · Groups IA, IIA and aluminium · Transition metals

• Metals and Non-metals: Properties and the Reactivity Series – Metals are usually shiny, hard, malleable, ductile and good conductors; non-metals are usually dull, brittle and poor conductors. Metals form basic oxides, react with water and dilute acids (giving hydrogen) and a more reactive metal pushes a less reactive one out of its salt solution. Listing metals from most to least reactive gives the reactivity series.
• Transition Elements (d-Block) – Transition elements are the d-block metals of groups 3 to 12. Their last electron goes into an inner (n−1)d orbital. Because the d-orbitals are only partly filled, they show many oxidation states, coloured ions, magnetism (μ = √(n(n+2)) BM), good catalytic power and easy alloy formation.

4. Chemistry and life

Chemical industry · Chemistry and health, food, home · Chemistry in building and agriculture

Coming soon