National Year 12 Chemistry
Chapters: 3
1. 3.1 Physical chemistry
3.1.1 Atomic structure · 3.1.2 Amount of substance · 3.1.3 Bonding · 3.1.4 Energetics · 3.1.5 Kinetics · 3.1.6 Chemical equilibria, Le Chatelier's principle and Kc · 3.1.7 Oxidation, reduction and redox equations
- Subatomic Particles and Early Atomic Models – An atom is made of three main particles. Electrons (charge −1, very light) were found with cathode ray tubes. Protons (charge +1, about 1836 times heavier) were found in canal rays. Neutrons (no charge, mass about a proton) were found by Chadwick in 1932. Thomson pictured the atom as a ball of positive charge with electrons stuck in it. Rutherford fired alpha particles at gold foil and found a tiny, heavy, positive nucleus with electrons moving outside. Atomic number Z = number of protons; mass number A = protons + neutrons. Rutherford's model could not explain why atoms are stable or why they give line spectra, which led to Bohr's model.
- Mole Concept, Molar Mass and Chemical Formulas – Atoms are far too small to count one by one, so chemists count them by weighing. Atomic masses are given in u, where 1 u is one-twelfth the mass of a carbon-12 atom. One mole is 6.022 × 10²³ particles, and its mass in grams equals the formula mass in u. With moles we can find the percentage of each element in a compound, and work back from percentages to the empirical and molecular formulas.
- 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.
- 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.
- Redox Reactions: From Oxygen to Electron Transfer – Oxidation first meant adding oxygen or removing hydrogen. Reduction meant the opposite. Today we use a bigger idea: oxidation is losing electrons and reduction is gaining electrons. Both always happen together, so we call them redox reactions. A more active metal gives electrons to the ion of a less active metal.
2. 3.2 Inorganic chemistry
3.2.1 Periodicity · 3.2.2 Group 2, the alkaline earth metals · 3.2.3 Group 7(17), the halogens
- 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.
- Group 2, the Alkaline Earth Metals: Trends, Reactions and Uses – Group 2 metals (Be, Mg, Ca, Sr, Ba) have two outer s electrons and form M²⁺ ions. Down the group atomic radius increases, first ionisation energy decreases and melting point generally decreases (Mg is an exception). Reactivity with water increases: M + 2H₂O → M(OH)₂ + H₂, while Mg reacts with steam to give MgO + H₂. Hydroxides become more soluble down the group; sulfates become less soluble, so BaSO₄ is insoluble. Uses: Mg(OH)₂ antacid, Ca(OH)₂ for acidic soil, CaO/CaCO₃ to remove SO₂, BaSO₄ for X-ray meals, BaCl₂ to test for sulfate, Mg to extract titanium.
- 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.
3. 3.3 Organic chemistry
3.3.1 Introduction to organic chemistry · 3.3.2 Alkanes · 3.3.3 Halogenoalkanes · 3.3.4 Alkenes · 3.3.5 Alcohols · 3.3.6 Organic analysis
- IUPAC Nomenclature and Isomerism of Organic Compounds – Millions of organic compounds exist, so every one needs a clear, unique name. The IUPAC system builds a name from three parts: the word root (number of carbons in the main chain), the suffix (type of bond and main functional group) and the prefixes (side groups). You pick the longest chain that holds the main group, number it so the main group and branches get the lowest numbers, and write side groups in alphabetical order. Isomers are compounds with the same molecular formula but different arrangement. Structural isomers differ in how atoms are joined (chain, position, functional group, metamerism). Stereoisomers are joined the same way but differ in 3D arrangement (geometrical cis/trans and optical isomers).
- Alkanes – Alkanes are hydrocarbons with only single C–C bonds. Their general formula is CnH2n+2. We make them by adding hydrogen to alkenes, by the Wurtz reaction, by reducing alkyl halides and by decarboxylation. They are non-polar, burn with a lot of heat and swap H for Cl in light. Around the C–C bond the molecule can twist into staggered and eclipsed shapes called conformations.
- Haloalkanes – A haloalkane is an alkane in which a hydrogen is replaced by a halogen (R–X). The C–X bond is polar: carbon is slightly positive, so nucleophiles attack it. Haloalkanes react by SN2 (one step, backside attack, inversion; best for 1°), SN1 (two steps, flat carbocation, racemic product; best for 3°) or elimination (strong base in alcohol gives an alkene).
- Alkenes – Alkenes are hydrocarbons with one C=C double bond and the formula CnH2n. The double bond is one σ bond plus one π bond, so the molecule is flat around it and cannot twist. That gives cis and trans isomers. We make alkenes by removing small molecules (H₂O, HX, X₂) or by partly adding H₂ to alkynes. Alkenes mainly do addition reactions: H₂, X₂, HX (Markovnikov rule, or anti-Markovnikov with peroxide), water, ozone, KMnO₄ and polymerisation.
- Alcohols – An alcohol is a carbon compound with an –OH group on an sp³ carbon (R–OH). Count the carbons joined to the C–OH carbon to classify it as primary (1°), secondary (2°) or tertiary (3°). Name it as alkan-n-ol, giving –OH the lowest number. Make alcohols from alkenes (acid-catalysed hydration or hydroboration–oxidation), by reducing aldehydes, ketones and acids, or from Grignard reagents. H-bonding gives them high boiling points and water solubility. Reactions break either the O–H bond (acidic: Na, esterification) or the C–O bond (HX, dehydration), and 1°/2° alcohols can be oxidised. Methanol and ethanol are the key commercial alcohols.
- Organic Analysis: Identifying Organic Compounds – Chemists identify an unknown organic compound in two ways. Quick test-tube reactions show which functional group is present: bromine water goes colourless with alkenes, acidified dichromate turns green with primary and secondary alcohols and aldehydes, Tollens' reagent gives a silver mirror with aldehydes, and sodium carbonate fizzes with carboxylic acids. Instruments then confirm it: a mass spectrum gives the relative molecular mass from the molecular ion peak (M⁺), and an infrared spectrum shows which bonds are present.