National Year 10 Chemistry
Chapters: 5
1. 4.1 Atomic structure and the periodic table
4.1.1 A simple model of the atom, symbols, relative atomic mass, electronic charge and isotopes · 4.1.2 The periodic table · 4.1.3 Properties of transition metals (Chemistry only)
- 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.
- The Periodic Table: Groups, Periods and Trends – The periodic table lists elements in order of atomic number (number of protons). Rows are periods; columns are groups. Elements in the same group have the same number of outer-shell electrons, so they react in similar ways. Mendeleev built an early table, left gaps for unknown elements and predicted their properties. Metals are on the left and bottom, non-metals on the right and top. Group 0 noble gases have full outer shells and are unreactive. Group 1 alkali metals get more reactive going down. Group 7 halogens get less reactive going down.
- 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.
2. 4.2 Bonding, structure, and the properties of matter
4.2.1 Chemical bonds, ionic, covalent and metallic · 4.2.2 How bonding and structure are related to the properties of substances · 4.2.3 Structure and bonding of carbon · 4.2.4 Bulk and surface properties of matter including nanoparticles (Chemistry only)
- Chemical Bonding: Ionic, Covalent and Metallic Bonds – Atoms join together (bond) to become more stable. Only their outer electrons take part. Most atoms are most stable with 8 outer electrons: the octet rule. There are three main ways to reach it. In an ionic bond, a metal gives electrons to a non-metal, making oppositely charged ions that attract. In a covalent bond, two non-metals share pairs of electrons to make molecules. In a metallic bond, metal atoms release outer electrons into a shared 'sea' that holds positive ions together. The difference in electronegativity (how strongly an atom pulls shared electrons) tells us which kind of bond forms. The type of bond explains melting points, whether a substance conducts electricity, and whether it dissolves in water.
- How Structure Decides the Properties of Substances – The way particles are held together decides how a substance behaves. Ionic and giant covalent substances have strong bonds all through, so they melt very high. Small molecules have weak forces between molecules, so they melt and boil low. Metals have free electrons, so they conduct.
- Allotropes of Carbon: Diamond, Graphite, Graphene and Fullerenes – Allotropes are different forms of the same element in the same state. Carbon has several. In diamond each atom makes 4 strong covalent bonds in a giant 3D network, so it is very hard, has a very high melting point and does not conduct electricity. In graphite each atom bonds to 3 others in flat layers of hexagons; the layers slide (soft, slippery) and one spare electron per atom is free to move, so graphite conducts. Graphene is a single graphite layer: strong, light and an excellent conductor. Fullerenes such as C₆₀ are hollow cages; nanotubes are rolled-up tubes. All of them burn in oxygen to make carbon dioxide.
- Nanoparticles: Why Tiny Particles Behave Differently – Nanoparticles are 1 to 100 nanometres (nm) across; 1 nm = 1 × 10⁻⁹ m. They are only a few hundred atoms wide, so a large share of their atoms sit on the surface. Their very high surface area to volume ratio makes them react, block light and kill germs differently from the same material in bulk.
3. 4.3 Quantitative chemistry
4.3.1 Chemical measurements, conservation of mass and quantitative interpretation of equations · 4.3.2 Use of amount of substance in relation to masses of pure substances · 4.3.3 Yield and atom economy of chemical reactions (Chemistry only) · 4.3.4 Concentrations of solutions in mol/dm3 (Chemistry only) (HT) · 4.3.5 Amount of substance and volumes of gases (Chemistry only) (HT)
- Atoms and Molecules – In a chemical reaction mass is neither created nor destroyed (conservation of mass), and a compound always has its elements in the same ratio by mass (constant proportions). Dalton explained both: matter is made of tiny atoms that join in small whole numbers. Atoms join to form molecules; charged atoms or groups are ions. Formulae are written by crossing valencies. Molecular mass (or formula unit mass for ionic compounds) is the sum of the atomic masses in the formula, in u.
- Moles: Counting Particles by Weighing – A mole is a fixed number of particles: 6.02 × 10²³ (the Avogadro constant). One mole of a substance has a mass in grams equal to its relative formula mass (Mr). So moles = mass ÷ Mr. Balanced equations tell you the ratio of moles that react, and the reactant that runs out first is the limiting reactant.
- Percentage Yield and Atom Economy – Percentage yield compares the product you actually got with the most you could get: actual ÷ theoretical × 100. Atom economy tells you what share of the reactant atoms end up in the product you want: Mr of wanted product ÷ total Mr of reactants × 100. Both help make chemistry cheaper and less wasteful.
- Titration: Finding an Unknown Concentration – Titration finds the unknown concentration of a solution. A solution of known concentration (the titrant) is added slowly from a burette to a measured volume of the unknown (the analyte) until the reaction is just complete. An indicator or a meter shows this point. At the equivalence point the moles react in the ratio of the equation, so for a 1 : 1 reaction C₁V₁ = C₂V₂.
- Avogadro's Law: Equal Volumes, Equal Numbers of Molecules – Avogadro's law: at the same temperature and pressure, equal volumes of all gases hold the same number of molecules. So volume is proportional to the number of moles: V ∝ n, or V₁/n₁ = V₂/n₂. At STP (0 °C, 101.3 kPa) one mole of any gas fills 22.4 L (molar volume, Vm). Useful links: n = V ÷ Vm, N = n × 6.02 × 10²³, m = n × M. Equal volumes have different masses, so relative density D = M₁ ÷ M₂ (D by hydrogen = M ÷ 2; by air = M ÷ 29). In gas reactions, volumes are in the same ratio as the equation coefficients (2H₂ + O₂ → 2H₂O: 2 : 1 : 2).
4. 4.4 Chemical changes
4.4.1 Reactivity of metals · 4.4.2 Reactions of acids · 4.4.3 Electrolysis
- 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.
- Acids and Bases: Properties, Indicators and Reactions – An acid gives H⁺ ions in water and a base gives OH⁻ ions. Indicators show which one is present by a colour or smell change, and when H⁺ meets OH⁻ they make water, leaving a salt behind.
- Electrolysis: Splitting Compounds with Electricity – Electrolysis uses a direct current to break down an ionic compound that is melted or dissolved (the electrolyte). Positive ions move to the cathode (−) and gain electrons (reduction). Negative ions move to the anode (+) and lose electrons (oxidation). Molten compounds give the metal and the non-metal. In solutions, water also gives H⁺ and OH⁻: hydrogen forms at the cathode unless the metal is less reactive than hydrogen, and oxygen forms at the anode unless a halide is present. Reactive metals like aluminium are extracted by electrolysis.
5. 4.5 Energy changes
4.5.1 Exothermic and endothermic reactions · 4.5.2 Chemical cells and fuel cells (Chemistry only)
- Exothermic and Endothermic Reactions – Energy is conserved in every reaction; it only moves between the chemicals and the surroundings. An exothermic reaction transfers energy to the surroundings, so the temperature rises (burning, neutralisation, respiration, hand warmers). An endothermic reaction takes energy in, so the temperature falls (thermal decomposition, citric acid + sodium hydrogencarbonate, photosynthesis, cold packs). A reaction profile shows reactants, products and the activation energy (the minimum energy needed to react). Breaking bonds takes energy in; making bonds gives energy out. ΔH = energy to break bonds − energy released making bonds; negative means exothermic.
- Chemical Cells and Fuel Cells – A chemical cell turns chemical energy into electrical energy. It has two different metals (electrodes) in an electrolyte. The more reactive metal gives away electrons, which flow through the wire to the less reactive metal. The bigger the difference in reactivity, the bigger the voltage; the electrolyte matters too. A battery is two or more cells in series, so their voltages add. Non-rechargeable cells stop when a reactant is used up; rechargeable cells can be reversed by an external current. A hydrogen fuel cell combines hydrogen and oxygen to make water and electricity, as long as fuel is supplied.