Russia 9 класс Chemistry (advanced)

1. Substance and chemical reaction

Atom properties · Structure of substance · Chemical thermodynamics · Kinetics · Chemical equilibrium · Redox reactions · Electrolytic dissociation · Ion exchange and hydrolysis

• Structure of the Atom – An atom has a tiny, heavy, positive nucleus made of protons and neutrons. Electrons move around it in fixed shells K, L, M, N. The number of protons (Z) tells the element; protons + neutrons give the mass number (A). Outer electrons decide valency. Isotopes share Z; isobars share A.
• Ionic Compounds: How Metals and Non-metals React – A metal atom gives its outer electrons to a non-metal atom so that both get a full outer shell (octet). The metal becomes a positive ion (cation), the non-metal a negative ion (anion), and the strong pull between them is an ionic (electrovalent) bond. Ionic compounds are hard crystalline solids with high melting points, dissolve in water, and conduct electricity only when molten or dissolved.
• 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.
• Rates of Reaction – The rate of a reaction is how fast reactants are used up or products are made. Rate = change in amount ÷ time. Particles must collide with at least the activation energy to react. More frequent, more energetic collisions mean a faster rate. Raising concentration (or gas pressure), temperature or surface area, or adding a catalyst, makes reactions faster. A catalyst gives a path with lower activation energy and is not used up.
• 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.
• Oxidation and Reduction: Corrosion and Rancidity – Oxidation is gain of oxygen (or loss of hydrogen); reduction is loss of oxygen (or gain of hydrogen). They always happen together, so these are called redox reactions. The substance that gives oxygen is the oxidising agent; the one that takes it is the reducing agent. In daily life, oxidation causes corrosion (like rusting of iron) and rancidity (fats and oils going stale).
• Electrolytes and Electrolytic Dissociation – An electrolyte is a substance that splits into free-moving ions when it dissolves in water (or melts), so its solution carries electric current. Non-electrolytes such as sugar stay as neutral molecules and do not conduct. Strong electrolytes split almost fully; weak ones split only a little. The degree of dissociation α = ions-forming units ÷ total units.
• Ionic Equations and Ion-Exchange Reactions – Salts, strong acids and alkalis split into ions in water: they are electrolytes. When two such solutions mix, a reaction happens only if some ions are removed from the solution as a precipitate, a gas or water. A full ionic equation shows every ion; crossing out the ions that do not change (spectator ions) gives the net ionic equation, for example Ag⁺ + Cl⁻ → AgCl↓. Charges and atoms must balance on both sides.

2. Non-metals and their compounds

General properties of non-metals · Halogens · Oxygen and sulfur · Nitrogen and phosphorus · Carbon, silicon, boron · First organic substances

• 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.
• 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 and their compounds

General properties of metals · Metals of A-groups · Metals of B-groups

• 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.
• Occurrence and Extraction of Metals – Metals are found in the earth's crust, mostly as compounds called minerals; a mineral from which a metal can be taken out profitably is an ore. Extraction has three stages: enrichment (removing gangue), getting the crude metal (roasting or calcination, then reduction or electrolysis, chosen by the metal's place in the reactivity series), and refining (usually electrolytic).
• Ionic Compounds: How Metals and Non-metals React – A metal atom gives its outer electrons to a non-metal atom so that both get a full outer shell (octet). The metal becomes a positive ion (cation), the non-metal a negative ion (anion), and the strong pull between them is an ionic (electrovalent) bond. Ionic compounds are hard crystalline solids with high melting points, dissolve in water, and conduct electricity only when molten or dissolved.
• 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 Metals: Basic Properties – Transition metals are the block of elements in the middle of the periodic table (the d-block), such as iron, copper, nickel, chromium, manganese, cobalt, titanium, silver and gold. Compared with Group 1 metals like sodium, they have much higher melting points, higher densities, and are harder and stronger; they react far more slowly with water and oxygen. They form ions with different charges (Fe²⁺ and Fe³⁺, Cu⁺ and Cu²⁺), make coloured compounds, and are useful catalysts. Their atoms fill an inner d sub-shell; chromium and copper 'borrow' an electron from 4s to get a half-full or full 3d. Many of their ions also form complex ions with water or ammonia.

4. Chemistry and the environment

Materials and green chemistry

• Green Chemistry – Green chemistry means designing chemical products and processes that make less waste, use safer substances, save energy and use raw materials that can be renewed. Key tools are high atom economy, catalysts, safer solvents such as water, renewable feedstocks, and products that break down safely. A life cycle assessment checks the impact of a product from raw material to disposal.

5. Review of grades 8–9

Systematisation

Coming soon