Chemistry (UK A-Level)
Comprehensive A-Level Chemistry covering physical, inorganic and organic chemistry to AQA/OCR/Edexcel standard. Includes atomic structure, bonding, energetics, kinetics, equilibria, periodicity, transition metals, organic synthesis and spectroscopy.
Ämne: Kemi · Nivå: Gymnasium (16–19) · 415 kort
Innehåll
- The three sub-atomic particles are the proton (relative mass 1, charge +1), neutron (relative mass 1, charge 0) and electron (relative mass 1/1836, charge -1).
- Mass number (A) = number of protons + number of neutrons. Atomic number (Z) = number of protons, which defines the element.
- Isotopes are atoms of the same element with the same number of protons but different numbers of neutrons. They have identical chemical properties but slightly different physical properties (e.g. density).
- Relative atomic mass (Ar) is the weighted mean mass of an atom of an element relative to 1/12 the mass of an atom of carbon-12.
- Relative isotopic mass is the mass of an atom of an isotope relative to 1/12 the mass of an atom of carbon-12. Isotopic masses are very close to whole numbers.
- A time-of-flight (TOF) mass spectrometer has five stages: ionisation, acceleration, ion drift, detection and analysis. It can determine relative atomic mass and identify isotopes.
- Two ionisation methods in TOF mass spectrometry: electrospray ionisation (sample dissolved in volatile solvent, injected through a fine needle at high voltage, forms MH⁺) and electron impact (sample vaporised, hit by high-energy electrons, forms M⁺).
- In a TOF mass spectrometer, all ions are accelerated to the same kinetic energy, so lighter ions travel faster and reach the detector first. Time of flight depends on the mass/charge ratio (m/z).
- Electrons occupy energy levels (shells) divided into sub-shells (s, p, d, f) made of orbitals. An orbital holds a maximum of 2 electrons with opposite spins. s holds 2, p holds 6, d holds 10, f holds 14 electrons.
- Electron configuration of chromium is [Ar] 3d⁵ 4s¹ and copper is [Ar] 3d¹⁰ 4s¹ — these are exceptions because a half-full or full 3d sub-shell gives extra stability.
- The 4s sub-shell fills before 3d (lower energy when empty), but when transition metals form ions, electrons are removed from 4s first. E.g. Fe²⁺ is [Ar] 3d⁶, not [Ar] 3d⁴ 4s².
- First ionisation energy is the energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1+ ions. Equation: X(g) → X⁺(g) + e⁻.
- Three factors affecting ionisation energy: nuclear charge (more protons = stronger attraction), atomic radius (greater distance = weaker attraction) and shielding (more inner shells = weaker attraction on outer electrons).
- Successive ionisation energies always increase because each electron is removed from an increasingly positive ion. Large jumps indicate a change of shell, giving evidence for electron shell structure.
- The dip in first ionisation energy from Group 2 to Group 3 (e.g. Mg to Al) is because the outer electron in Group 3 is in a higher-energy p sub-shell, which is easier to remove.
- The dip in first ionisation energy from Group 5 to Group 6 (e.g. N to O) occurs because in Group 6 two electrons pair up in one p orbital, and electron-electron repulsion makes one easier to remove.
- The Avogadro constant is 6.022 × 10²³ mol⁻¹ — the number of particles in one mole of a substance.
- Number of moles = mass (g) / molar mass (g mol⁻¹). This is one of the most-used equations in chemistry: n = m/M.
- For a solution, number of moles = concentration (mol dm⁻³) × volume (dm³). Note 1 dm³ = 1000 cm³, so divide a cm³ volume by 1000.
- The ideal gas equation is pV = nRT, where p is pressure (Pa), V is volume (m³), n is moles, R is the gas constant (8.31 J K⁻¹ mol⁻¹) and T is temperature (K).
- Empirical formula is the simplest whole-number ratio of atoms of each element in a compound. Molecular formula gives the actual number of atoms of each element in a molecule.
- Percentage yield = (actual mass of product / theoretical mass of product) × 100. It is always ≤ 100% because of losses, side reactions and incomplete reactions.
- Atom economy = (molar mass of desired product / sum of molar masses of all products) × 100. High atom economy means less waste and more sustainable processes.
- The molar gas volume at room temperature and pressure (RTP, 25°C and 100 kPa) is approximately 24 dm³ mol⁻¹. So moles of gas = volume (dm³) / 24.
- The limiting reagent is the reactant completely used up in a reaction; it determines the maximum amount of product. The other reactant(s) are in excess.
- Ionic bonding is the electrostatic attraction between oppositely charged ions, formed by transfer of electrons from a metal to a non-metal. It produces giant ionic lattices.
- Covalent bonding is a shared pair of electrons between two non-metal atoms. The bond is the electrostatic attraction between the shared electrons and the two nuclei.
- A dative covalent (coordinate) bond is a covalent bond where both shared electrons come from the same atom. Examples: NH₄⁺ and H₃O⁺. Shown by an arrow from donor to acceptor.
- Metallic bonding is the electrostatic attraction between positive metal ions and a sea of delocalised electrons. It explains conductivity, malleability and high melting points.
- Electronegativity is the ability of an atom to attract the bonding electrons in a covalent bond. It increases across a period and decreases down a group. Fluorine is the most electronegative element (Pauling 4.0).
- A polar covalent bond has an unequal sharing of electrons due to an electronegativity difference, creating a permanent dipole (δ+ and δ- charges). A large difference (>1.7) gives ionic character.
- VSEPR theory: electron pairs around a central atom repel and arrange to be as far apart as possible. Lone pairs repel more strongly than bonding pairs, reducing bond angles by about 2.5° per lone pair.
- Common molecular shapes and angles: linear (180°), trigonal planar (120°), tetrahedral (109.5°), trigonal bipyramidal (120° and 90°), octahedral (90°).
- Water (H₂O) is bent/V-shaped with a bond angle of 104.5° due to two lone pairs. Ammonia (NH₃) is trigonal pyramidal with a bond angle of 107° due to one lone pair.
- The three types of intermolecular force, in order of increasing strength: van der Waals (London/induced dipole-dipole), permanent dipole-dipole, and hydrogen bonding.
- Van der Waals forces (London forces) arise from instantaneous dipoles caused by random electron movement, which induce dipoles in neighbouring molecules. They increase with more electrons and larger surface area.
- Hydrogen bonding occurs when H is bonded to N, O or F (highly electronegative) and is attracted to a lone pair on N, O or F of another molecule. It is the strongest intermolecular force.
- Ice is less dense than liquid water because hydrogen bonds hold the molecules in an open tetrahedral lattice with gaps. This anomaly means ice floats.
- Enthalpy change (ΔH) is the heat energy transferred at constant pressure. Exothermic reactions release heat (ΔH negative); endothermic reactions absorb heat (ΔH positive). Units: kJ mol⁻¹.
- Standard conditions for enthalpy measurements: 100 kPa pressure, a stated temperature (usually 298 K), and solutions at 1 mol dm⁻³. Denoted by the symbol ⊕ (standard state).
- Standard enthalpy of formation (ΔfH⊕) is the enthalpy change when one mole of a compound is formed from its elements in their standard states. ΔfH⊕ of any element in its standard state is zero.
- Standard enthalpy of combustion (ΔcH⊕) is the enthalpy change when one mole of a substance is completely burned in excess oxygen under standard conditions. Always exothermic (negative).
- Hess's law states the total enthalpy change of a reaction is independent of the route taken, provided the initial and final conditions are the same. It follows from conservation of energy.
- The calorimetry equation is q = mcΔT, where q is heat energy (J), m is mass of solution/water (g), c is specific heat capacity (4.18 J g⁻¹ K⁻¹ for water) and ΔT is temperature change.
- Mean bond enthalpy is the average energy needed to break one mole of a covalent bond in the gaseous state, averaged over different molecules. ΔH = Σ(bonds broken) - Σ(bonds formed).
- Collision theory: for a reaction to occur, particles must collide with energy greater than or equal to the activation energy AND with the correct orientation.
- Activation energy (Ea) is the minimum energy required for a reaction to occur — the energy needed to break bonds and start the reaction.
- A Maxwell-Boltzmann distribution shows the spread of molecular energies in a gas. The area under the curve equals the total number of molecules; only molecules to the right of Ea can react.
- A catalyst increases the rate of reaction by providing an alternative reaction pathway with a lower activation energy. It is not used up and does not change ΔH or the position of equilibrium.
- The rate equation has the form rate = k[A]ᵐ[B]ⁿ, where k is the rate constant, and m and n are the orders with respect to each reactant. Orders are found experimentally, not from the equation.