Comprehensive IB Chemistry SL Study Notes

Subatomic Particles & Atomic Structure

• Subatomic Particles:

  • Proton: Relative charge $+1$, relative mass $1$, actual charge $+1.602189 \times 10^{-19}\,\text{C}$.

  • Neutron: Relative charge $0$, relative mass $1$.

  • Electron: Relative charge $-1$, relative mass is negligible ($1836$ times lighter than a proton or neutron), actual charge $-1.602189 \times 10^{-19}\,\text{C}$.

• Atomic Notation:     

  

  • Mass Number ($A$): Total number of nucleons (protons + neutrons).

  • Atomic Number ($Z$): Total number of protons in the nucleus.

  • $\text{Number of neutrons} = A - Z$

• Ions:

  • Cations (positive ions) form when an atom loses electrons.

  • Anions (negative ions) form when an atom gains electrons.

• Isotopes:

  • Atoms of the same element with the same number of protons ($Z$) but different numbers of neutrons (different mass number $A$).

  • Chemical properties remain identical due to having the same electron configuration.

  • Physical properties (density, mass, melting/boiling point, rate of diffusion) differ.

• Relative Atomic Mass ($A_r$):

  • Weighted average mass of an atom relative to $\frac{1}{12}$th the mass of a carbon-12 atom:         $A_r = \frac{\text{weighted average mass of one atom of an element}}{\frac{1}{12}\text{ mass of one atom of carbon-12}}$

  • Calculated using percentage abundances:         $A_r = \frac{\sum (\%\,\text{abundance} \times \text{mass})}{100}$

Electromagnetic Spectrum & Emission Spectra

• Electromagnetic Spectrum:     

  

  • All electromagnetic waves travel at the speed of light in a vacuum ($c = 3.00 \times 10^8\,\text{m\,s}^{-1}$).

  • Wavelength ($\lambda$) and frequency ($f$) are inversely related:         $c = f\lambda$

  • High frequency / short wavelength radiation (Gamma rays, X-rays, UV) carries high energy.

• Continuous vs. Line Spectra:

  • Continuous Spectrum: Shows all frequencies/wavelengths of visible light without gaps.

  • Line Spectrum: Shows only discrete frequencies, proving that atomic energy levels are quantised.

• Hydrogen Emission Spectrum:     

  

  • Emitted light originates from excited electrons releasing energy as they transition from higher to lower principal energy levels.

  • Spectral lines converge at higher energies (shorter wavelengths / blue end) towards the ionisation energy limit.

  • Lyman Series: Transitions to $n = 1$ (Ultraviolet region, high energy).

  • Balmer Series: Transitions to $n = 2$ (Visible light region, medium energy; named after Johannes Balmer).

  • Paschen / Ritz-Paschen Series: Transitions to $n = 3$ (Infrared region, low energy).

Subshells & Electron Configurations

• Principal Energy Levels & Subshells:

  • Principal quantum shells ($n$) hold a maximum of $2n^2$ electrons ($n=1: 2$, $n=2: 8$, $n=3: 18$, $n=4: 32$).

  • Subshell capacities: $s$ (1 orbital, 2 electrons), $p$ (3 orbitals, 6 electrons), $d$ (5 orbitals, 10 electrons), $f$ (7 orbitals, 14 electrons).

• Orbital Shapes:     

  

  • $s$ Orbitals: Spherical in shape, increasing in size with principal quantum shell number.

  • $p$ Orbitals: Dumbbell-shaped, arranged perpendicularly along the $x$, $y$, and $z$ axes ($p_x, p_y, p_z$).

• Rules for Writing Configurations:

  • Aufbau Principle: Electrons fill subshells in order of increasing energy ($1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p$).

  • Pauli Exclusion Principle: An atomic orbital holds up to two electrons with opposite spins.

  • Hund's Rule: Degenerate orbitals fill singly with parallel spins before pairing up to minimize spin-pair repulsion.

• Exceptions & Transition Metals:

  • Chromium ($\text{Cr}$): $[\text{Ar}]\,3d^5\,4s^1$

  • Copper ($\text{Cu}$): $[\text{Ar}]\,3d^{10}\,4s^1$

  • Transition metal ions lose electrons from the $4s$ subshell before the $3d$ subshell.

Stoichiometry & The Mole Concept

• The Mole Concept:     

  

  • One mole ($\text{mol}$) contains $L = 6.02 \times 10^{23}\,\text{mol}^{-1}$ particles (Avogadro constant).

  • $\text{Number of particles} = n \times L$

  • $\text{Mass } (m) = n \times M$

• Formulae:

  • Empirical Formula: The simplest whole-number ratio of atoms of each element in a compound.

  • Molecular Formula: The actual number of atoms of each element in a molecule.         $\text{Whole-number multiplier} = \frac{M_r}{\text{empirical formula mass}}$

• Molar Concentration:

  • Concentration in moles: $c\,(\text{mol\,dm}^{-3}) = \frac{n\,(\text{mol})}{V\,(\text{dm}^3)}$

  • Concentration in mass: $c\,(\text{g\,dm}^{-3}) = \frac{m\,(\text{g})}{V\,(\text{dm}^3)}$

  • $\text{Parts per million (ppm)} = 1\,\text{mg\,dm}^{-3} = 1\,\text{mg\,kg}^{-1}$

• Avogadro's Law for Gases:

  • Equal volumes of gases under identical temperature and pressure contain equal numbers of particles.

  • Molar volume of an ideal gas at STP ($0\,^\circ\text{C} / 273.15\,\text{K}$ and $100\,\text{kPa}$) is $22.7\,\text{dm}^3\,\text{mol}^{-1}$.

Particulate Nature of Matter & Separation Techniques

• Classification of Matter:

  • Elements: Pure substances made of one type of atom.

  • Compounds: Two or more elements chemically bonded in fixed stoichiometric ratios.

  • Mixtures: Substances physically combined in non-fixed ratios (Homogeneous = uniform composition; Heterogeneous = non-uniform composition).

• Separation Methods:

  • Filtration: Separates an insoluble solid residue from a liquid filtrate.

  • Crystallisation & Recrystallisation: Purifies solids based on differential solubility in hot vs. cold solvents.

  • Distillation:

    • Simple Distillation: Separates liquid solvent from a solution with widely different boiling points.

    • Fractional Distillation: Separates miscible liquids with close boiling points using a fractionating column.

  • Paper Chromatography: Separates dissolved substances based on relative solubility in the mobile phase versus adsorption onto the stationary phase.

• States of Matter & Phase Transitions:     

  

  • Endothermic transitions (energy absorbed): Melting, vaporisation/boiling, sublimation.

  • Exothermic transitions (energy released): Freezing, condensation, deposition.

  • During state changes, temperature remains constant because thermal energy breaks intermolecular forces rather than increasing particle kinetic energy.

• Temperature & Kinetic Energy:

  • Absolute temperature in Kelvin ($\text{K}$) is directly proportional to the average kinetic energy ($E_k$) of particles.

  • Absolute zero ($0\,\text{K}$ or $-273.15\,^\circ\text{C}$) is the point of zero kinetic energy.

  • Conversion formula:         $T\,(\text{K}) = \theta\,(^\circ\text{C}) + 273.15$