Overview
This lecture provides a comprehensive review of atomic structure, covering subatomic particles, ions, isotopes, historical models, mass spectrometry, electron configuration, and trends in ionization energy, specifically for the AQA Chemistry curriculum.
Atomic Structure and Particles
- Atoms consist of a nucleus (protons and neutrons) surrounded by electrons in shells.
- Protons have a +1 charge, neutrons are neutral, electrons have a -1 charge.
- Relative masses: proton = 1, neutron = 1, electron ≈ 1/2000.
- Atomic number = number of protons; mass number = protons + neutrons.
- Atoms are neutral: number of protons equals number of electrons.
Ions and Isotopes
- Ions: charged particles with unequal numbers of protons and electrons.
- Negative ions gain electrons, positive ions lose electrons.
- Isotopes: atoms of the same element with the same number of protons but different numbers of neutrons; they have different masses but identical chemical properties.
Historical Models of the Atom
- Dalton: atoms are solid spheres.
- J.J. Thomson: discovered electrons, proposed the "plum pudding" model.
- Rutherford: discovered the nucleus, most of the atom is empty space (gold foil experiment).
- Bohr: electrons occupy fixed energy levels (shells).
- Modern model includes subshells (s, p, d, f).
Mass Spectrometry and Relative Atomic Mass
- Time-of-flight mass spectrometer: vaporizes, ionizes, accelerates, and detects ions based on mass-to-charge (m/z) ratio.
- Relative atomic mass = Σ (isotope abundance × m/z) ÷ total abundance.
- Mass spectra show isotopic composition and help calculate relative atomic mass.
Electron Configuration and Subshells
- Electrons fill subshells in order of increasing energy: 1s < 2s < 2p < 3s < 3p < 4s < 3d, etc.
- Subshell capacities: s=2, p=6, d=10, f=14 electrons.
- Write configurations as: shell number, subshell, electron count (e.g., 1s² 2s² 2p⁶).
- For ions, remove electrons from the highest energy shell first; transition metals may have exceptions (e.g., Cr, Cu).
Ionization Energy and Trends
- First ionization energy: energy needed to remove one mole of electrons from one mole of gaseous atoms.
- Ionization energy decreases down a group due to increased atomic radius and electron shielding.
- Ionization energy increases across a period due to increased nuclear charge and similar shielding.
- Successive ionization energies increase, with large jumps after removing all outer-shell electrons.
- Exceptions in trends: lower ionization at Al (due to 3p subshell) and S (due to electron repulsion in paired orbitals).
Key Terms & Definitions
- Atom — Smallest unit of matter, made of protons, neutrons, and electrons.
- Ion — Atom or molecule with a net electric charge (gained or lost electrons).
- Isotope — Atoms of the same element with different numbers of neutrons.
- Relative Atomic Mass (Ar) — Average mass of an atom relative to 1/12th the mass of carbon-12.
- Mass Spectrometry — Technique to determine masses of particles and their relative abundance.
- Electron Configuration — Distribution of electrons among atomic orbitals.
- Ionization Energy — Energy required to remove an electron from an atom in gaseous state.
- Shielding — Reduction in effective nuclear attraction due to inner-shell electrons.
Action Items / Next Steps
- Review and memorize definitions, especially those related to atomic structure and mass spectrometry.
- Practice calculating relative atomic mass from isotopic data.
- Complete any assigned reading on electron configuration and periodic trends.