S1.1Introduction to the particulate nature of matter
Key terms
Element
A substance made of only one type of atom; it cannot be broken down chemically.
Compound
Two or more elements chemically combined in a fixed ratio, with different properties from its elements.
Mixture
Two or more substances physically combined in no fixed ratio; each keeps its own properties.
Homogeneous mixture
Uniform composition throughout, in a single phase (e.g. salt solution, air).
Heterogeneous mixture
Non-uniform composition, with visibly different phases (e.g. sand in water).
Kinetic molecular theory
Matter is made of particles in constant random motion; their average kinetic energy is proportional to absolute temperature.
\[T(\text{K})=\theta(^{\circ}\text{C})+273\qquad R_F=\frac{\text{distance moved by spot}}{\text{distance moved by solvent front}}\]
Separation techniques exploit differences in physical properties: filtration (insoluble solid from liquid), evaporation and crystallization (soluble solid from solution), distillation (liquids with different boiling points; fractional distillation for close boiling points), solvation (dissolving one component), paper chromatography (different affinities for the stationary and mobile phases).
States of matter: solids have fixed positions and vibrate; liquids are close together but move past each other; gases are far apart and move fast and randomly.
Changes of state: melting, freezing, vaporization (boiling or evaporation), condensation, sublimation (solid → gas) and deposition (gas → solid).
During a change of state the temperature stays constant: energy is used to overcome forces between particles, not to raise their kinetic energy.
The Kelvin scale starts at absolute zero (0 K), where particles have minimum kinetic energy. A change of 1 K equals a change of 1 °C.
Worked example
In paper chromatography a spot moves 3.6 cm while the solvent front moves 8.0 cm. Find the \(R_F\) value.
\(R_F=\dfrac{3.6}{8.0}=\mathbf{0.45}\).
\(R_F\) has no units and is always between 0 and 1. Comparing it with known values identifies the substance (same solvent and paper).
Exam tip: Name the physical property each separation technique relies on (solubility, boiling point, particle size): that is usually where the marks are.
S1.2The nuclear atom
Key terms
Atomic number (Z)
Number of protons in the nucleus; it defines the element.
Mass number (A)
Number of protons plus neutrons.
Isotopes
Atoms of the same element with different numbers of neutrons.
Relative atomic mass (Ar)
Weighted mean mass of an element's atoms compared with 1/12 of a carbon-12 atom.
Ion
A charged particle formed when an atom gains or loses electrons.
Protons: relative mass 1, charge +1. Neutrons: relative mass 1, charge 0. Electrons: relative mass about 1/1836 (≈ 0.0005), charge −1.
Almost all the mass of an atom is in the tiny nucleus; most of the atom's volume is the space occupied by electrons.
Isotopes have the same chemical properties (same electron arrangement) but different physical properties such as mass, density and rate of diffusion.
In an ion, electrons = protons − charge: \({}^{32}_{16}\text{S}^{2-}\) has 16 protons, 16 neutrons and 18 electrons.
HL A mass spectrometer separates isotopes by their mass-to-charge ratio (m/z). The heights of the peaks give relative abundances, from which \(A_r\) is calculated. Fragments in the spectra of compounds are used to identify molecules (Structure 3.2).
Worked example
Chlorine is 75.8% chlorine-35 (mass 34.97) and 24.2% chlorine-37 (mass 36.97). Calculate its relative atomic mass.
A continuous spectrum has all wavelengths; a line spectrum has only certain ones, showing that electron energy levels are discrete.
In hydrogen, transitions to n = 1 give ultraviolet lines (Lyman series) and to n = 2 visible lines (Balmer series). Lines get closer together at higher energy and converge because the energy levels get closer.
Sublevels: s (1 orbital, 2 e⁻), p (3 orbitals, 6 e⁻), d (5 orbitals, 10 e⁻), f (7 orbitals, 14 e⁻). Filling order: 1s 2s 2p 3s 3p 4s 3d 4p.
Exceptions: chromium is [Ar]3d⁵4s¹ and copper is [Ar]3d¹⁰4s¹ (half-filled and full d sublevels are more stable).
Transition metal atoms lose their 4s electrons before their 3d electrons: Fe is [Ar]3d⁶4s², Fe²⁺ is [Ar]3d⁶ and Fe³⁺ is [Ar]3d⁵.
HL The ionization energy can be found from the convergence limit of the Lyman series: \(IE=\dfrac{hc}{\lambda}\times N_A\). Successive ionization energies jump when an electron is removed from a new, inner shell, revealing the group of an element. First ionization energies across a period show dips at groups 13 (new p sublevel) and 16 (first paired p electron, extra repulsion).
Worked example
Write the full electron configuration of iron (Z = 26) and the configuration of Fe³⁺.
Fe: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁶, or [Ar] 3d⁶ 4s².
Fe³⁺ loses the two 4s electrons first, then one 3d electron: [Ar] 3d⁵.
Exam tip: In orbital diagrams, fill each box of a sublevel with one up-arrow before adding any down-arrows (Hund's rule).
\(M_r\) is the sum of the \(A_r\) values in a formula; molar mass has the same number with units g mol⁻¹.
Empirical formula: divide each mass (or %) by \(A_r\), divide by the smallest, then multiply to whole numbers if needed.
Molecular formula = (empirical formula) × \(\dfrac{M_r}{M_r\text{ of empirical formula}}\).
Avogadro's law: equal volumes of gases at the same temperature and pressure contain equal numbers of particles, so gas volumes react in the same ratio as moles.
A standard solution has an accurately known concentration, made by dissolving a weighed solid in a volumetric flask.
Worked example
A compound is 40.0% carbon, 6.7% hydrogen and 53.3% oxygen by mass. Its \(M_r\) is 180. Find its empirical and molecular formulas.
In \(pV=nRT\), use SI units: \(p\) in Pa, \(V\) in m³ (1 dm³ = 10⁻³ m³), \(T\) in K.
Real gases deviate from ideal behaviour at high pressure (particle volume becomes significant) and low temperature (intermolecular forces become significant).