IB Chemistry HL
Comprehensive flashcards for IB Diploma Programme Chemistry at Higher Level (current guide, first exams 2025). Covers the full SL core plus all Additional Higher Level (AHL) extensions across the organising concepts Structure and Reactivity.
Ämne: Kemi · Nivå: Gymnasium (16–19) · 483 kort
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- The mole (mol) is the SI unit for amount of substance. One mole contains exactly 6.02 × 10²³ specified entities (the Avogadro constant, Nₐ).
- The Avogadro constant Nₐ = 6.02 × 10²³ mol⁻¹. Amount in moles n = N / Nₐ, where N is the number of particles.
- Molar mass (M) is the mass of one mole of a substance, in g mol⁻¹. It is numerically equal to the relative atomic or molecular mass. Amount n = m / M.
- Relative atomic mass (Ar) is the weighted average mass of an element's atoms relative to 1/12 the mass of a carbon-12 atom. It has no units.
- Relative molecular mass (Mr) is the sum of the relative atomic masses of all atoms in a molecule. For ionic/giant compounds the equivalent term is relative formula mass.
- The empirical formula is the simplest whole-number ratio of atoms of each element in a compound. The molecular formula is a whole-number multiple of the empirical formula.
- Avogadro's law: equal volumes of all gases, at the same temperature and pressure, contain equal numbers of molecules. Hence mole ratio equals volume ratio for gases.
- The molar volume of an ideal gas at STP (0 °C, 100 kPa) is 22.7 dm³ mol⁻¹. Amount of gas n = V / 22.7 when V is in dm³ at STP.
- The ideal gas equation is PV = nRT, where R = 8.31 J K⁻¹ mol⁻¹. Pressure in Pa, volume in m³, temperature in kelvin (K = °C + 273).
- Real gases deviate most from ideal behaviour at high pressure and low temperature, where particle volume and intermolecular forces become significant.
- Concentration (mol dm⁻³) = amount of solute (mol) ÷ volume of solution (dm³). Square brackets [X] denote concentration in mol dm⁻³.
- Percentage yield = (actual yield ÷ theoretical yield) × 100%. Atom economy = (mass of desired product ÷ total mass of reactants) × 100%.
- An atom has a dense central nucleus of protons (charge +1) and neutrons (charge 0), surrounded by electrons (charge −1) in energy levels. The nucleus holds nearly all the mass.
- Relative masses/charges: proton mass 1, charge +1; neutron mass 1, charge 0; electron mass ~1/1836 (negligible), charge −1.
- Atomic number (Z) = number of protons; it defines the element. Mass number (A) = protons + neutrons. Neutrons = A − Z.
- Isotopes are atoms of the same element with the same number of protons but different numbers of neutrons. They have identical chemical properties but different masses.
- Relative atomic mass is calculated from isotopic abundances: Ar = Σ(isotope mass × fractional abundance). Mass spectra give the abundances.
- In a mass spectrometer, atoms are ionised, accelerated, and separated by mass-to-charge ratio (m/z). Peaks show each isotope's m/z and relative abundance.
- Chlorine's Ar of 35.45 reflects roughly 75% ³⁵Cl and 25% ³⁷Cl. The non-integer value arises from the weighted average of isotope masses.
- Radioisotopes have unstable nuclei that decay, emitting radiation. Examples: carbon-14 for radiocarbon dating, cobalt-60 in radiotherapy, iodine-131 for thyroid imaging.
- Electrons occupy discrete energy levels (shells). Emission spectra are produced when excited electrons fall to lower levels, releasing photons of specific energy (E = hf).
- An emission spectrum is a line spectrum, not continuous. The discrete lines are direct evidence that electron energy levels are quantised (discrete, not continuous).
- In the hydrogen emission spectrum, the lines converge at higher frequency. The Lyman series (UV) ends at n=1, the Balmer series (visible) at n=2, the Paschen (IR) at n=3.
- Within a main energy level are sublevels: s (1 orbital, 2 e⁻), p (3 orbitals, 6 e⁻), d (5 orbitals, 10 e⁻), f (7 orbitals, 14 e⁻). Each orbital holds a maximum of 2 electrons.
- The Aufbau principle: electrons fill the lowest-energy sublevels first. Filling order: 1s 2s 2p 3s 3p 4s 3d 4p... (4s fills before 3d because it is lower in energy).
- Hund's rule: electrons occupy degenerate (equal-energy) orbitals singly with parallel spins before pairing. The Pauli exclusion principle: an orbital holds at most 2 electrons, with opposite spins.
- Chromium and copper are exceptions to Aufbau filling: Cr is [Ar]3d⁵ 4s¹ and Cu is [Ar]3d¹⁰ 4s¹, because half-full and full d-sublevels are extra stable.
- Orbitals have shapes: an s orbital is spherical; a p orbital is dumbbell-shaped with three orientations (px, py, pz) along the x, y and z axes.
- First ionisation energy is the energy needed to remove one mole of electrons from one mole of gaseous atoms: X(g) → X⁺(g) + e⁻. It is always endothermic (positive).
- First ionisation energy decreases down a group (electron is further from nucleus, more shielding) and generally increases across a period (greater nuclear charge, similar shielding).
- Successive ionisation energies of one element always increase: removing an electron from an increasingly positive ion requires more energy. Large jumps mark the start of a new inner shell.
- Successive ionisation energies are evidence for electron shells: a graph of log(IE) vs electron number shows distinct jumps that match the 2,8,8 shell pattern.
- Small dips in first ionisation energy across period 2 (B below Be; O below N) are HL evidence for sublevels: the 2p electron is higher in energy than 2s, and pairing in one 2p orbital adds repulsion.
- The convergence limit of the hydrogen Lyman series corresponds to ionisation (n=1 → n=∞). The frequency at convergence allows the ionisation energy of hydrogen to be calculated (HL).
- Continuous spectrum of white light spans all visible wavelengths; an absorption spectrum shows dark lines where atoms absorbed specific photons. Both connect to quantised electron transitions.
- Limiting (limiting) reagent is the reactant that is completely consumed and determines the maximum amount of product. The other reactant(s) are in excess.
- Ionic bonding is the electrostatic attraction between oppositely charged ions, formed when metals transfer electrons to non-metals. It typically occurs at large electronegativity differences (> ~1.8).
- Ionic compounds form giant 3D lattices held by strong electrostatic forces. They have high melting points, are brittle, and conduct electricity only when molten or in aqueous solution.
- Lattice enthalpy increases with greater ionic charge and smaller ionic radius (Coulomb's law: force ∝ q₁q₂/r²). MgO has a far higher lattice enthalpy than NaCl.
- Covalent bonding is the electrostatic attraction between a shared pair of electrons and the two nuclei. A single bond shares one pair, a double bond two pairs, a triple bond three pairs.
- Bond length decreases and bond strength increases from single to double to triple bonds between the same atoms (e.g. C–C > C=C > C≡C in length; reverse in strength).
- A coordinate (dative) covalent bond is a covalent bond in which both shared electrons come from the same atom, e.g. the bond from N to H⁺ in the ammonium ion NH₄⁺.
- Electronegativity is the ability of an atom to attract a shared pair of electrons in a covalent bond. It increases across a period and decreases down a group; fluorine is the most electronegative element.
- A polar covalent bond has unequal electron sharing due to an electronegativity difference, creating a dipole (δ+ and δ−). A nonpolar bond shares electrons equally (zero or near-zero difference).
- VSEPR theory: electron domains around a central atom repel and arrange to minimise repulsion. Lone pairs repel more strongly than bonding pairs, compressing bond angles.
- Two electron domains give a linear shape (180°), e.g. CO₂ and BeCl₂. Three bonding domains give trigonal planar (120°), e.g. BF₃.
- Four electron domains give a tetrahedral arrangement (109.5°). CH₄ is tetrahedral; NH₃ with one lone pair is trigonal pyramidal (107°); H₂O with two lone pairs is bent (104.5°).
- Five electron domains give trigonal bipyramidal (90° and 120°), e.g. PCl₅. Six domains give octahedral (90°), e.g. SF₆. These expanded octets are HL extensions.
- A molecule is polar if it has polar bonds AND an asymmetric shape so the dipoles do not cancel. CO₂ is nonpolar (linear, dipoles cancel); H₂O is polar (bent, dipoles add).
- London (dispersion) forces are temporary induced dipole attractions present between all molecules. They strengthen with more electrons (greater molar mass) and larger surface area.