United Grade 11 Chemistry
Chapters: 3
1. Structure and properties of matter
Periodic table patterns from electron structure · Bulk properties and electrical forces between particles · Nuclear processes: fission, fusion, decay · Designing materials from molecular structure
- 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.
- Intermolecular Forces – Intermolecular forces are weak attractions BETWEEN molecules. From weakest to strongest (for similar-size molecules): London dispersion < dipole–dipole < hydrogen bond. Stronger forces mean higher melting and boiling points, higher viscosity and surface tension, and they decide what dissolves in what.
- Nuclei: Size, Nuclear Force, Binding Energy, Fission and Fusion – A nucleus is made of Z protons and N neutrons (A = Z + N nucleons). Its radius is R = R₀A^(1/3) with R₀ ≈ 1.2 fm, so every nucleus has almost the same huge density. A very strong, short-range nuclear force holds the nucleons together. The nucleus weighs a little less than its loose parts; this mass defect Δm is the binding energy, E = Δm c² (1 u = 931.5 MeV). Binding energy per nucleon is highest near iron (A ≈ 56), so heavy nuclei give energy when they split (fission) and light nuclei give energy when they join (fusion).
- 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.
- 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.
2. Chemical reactions
Predicting simple reactions from outer electrons · Energy change in reactions (bond energy) · Reaction rates and collisions · Equilibrium and Le Chatelier · Conservation of mass and stoichiometry
- Chemical Reactions: Equations, Types, Prediction and Coupled Reactions – A chemical reaction rearranges atoms into new substances; atoms and mass are conserved, so equations must balance. The same reaction can be written as a word equation, a balanced formula equation with state symbols, a full ionic equation or a net ionic equation that leaves out spectator ions. Reactions are sorted into types: synthesis, decomposition, single displacement, double displacement (precipitation, neutralisation, gas-forming), combustion, and by the particle that moves: electrons (redox) or protons (acid–base). Outer (valence) electrons let us predict what forms: metals lose electrons, non-metals gain them, and the numbers lost and gained must match. A reaction that cannot happen alone (ΔG > 0) can be driven by coupling it to a strongly favourable one so that the total ΔG is negative.
- Enthalpies of Different Types of Reactions – Each kind of change has its own named enthalpy, always for 1 mol. Phase changes: fusion (melting), vaporisation (boiling) and sublimation (ΔsubH = ΔfusH + ΔvapH); all are endothermic. Combustion: 1 mol burns completely in oxygen; always exothermic. Atomisation: 1 mol of a substance breaks fully into gaseous atoms. Bond dissociation enthalpy: 1 mol of a given bond is broken in the gas phase; for polyatomic molecules we use mean bond enthalpy. ΔrH ≈ Σ(bonds broken) − Σ(bonds formed). Lattice enthalpy: 1 mol of an ionic solid is split into gaseous ions (found by a Born–Haber cycle). Enthalpy of solution = lattice enthalpy + hydration enthalpy.
- 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.
- 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.
- Stoichiometry, Limiting Reagent and Concentration Terms – Stoichiometry means measuring the amounts of substances in a reaction. A balanced equation acts like a recipe: its numbers give the mole ratio of reactants and products. With it you can change any mass into moles, use the ratio, and change back to the mass of any other substance. The reactant that runs out first is the limiting reagent; it decides how much product forms. For solutions, concentration is given as mass percent, mole fraction, molarity or molality.
3. Common course additions (beyond NGSS PEs)
Gas laws · Solutions and concentration · Acids, bases and pH · Molecular shape
- Perfect Gas Equation PV = nRT and Work in Compressing a Gas – A gas is made of countless tiny molecules flying about. Their hits on the walls make pressure. For a low-density gas, three simple laws hold: at fixed temperature, P × V stays constant (Boyle); at fixed pressure, V grows in step with kelvin temperature (Charles); at the same P and T, equal volumes hold equal numbers of molecules (Avogadro). Put together they give the perfect gas equation PV = nRT = N k T. One mole holds Avogadro's number, 6.022 × 10²³, of particles. Pushing a piston in does work on the gas; the work equals the area under the P–V graph, and at fixed temperature W = nRT ln(V₁/V₂).
- 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.
- 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.
- 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°).