Topic 1 Atomic structure and the periodic table

Overview

This lecture gives a comprehensive summary of Edexcel Chemistry Topic 1: Atomic Structure and the Periodic Table, covering atom structure, isotopes, mass spectrometry, electron configuration, periodic trends, and bonding.

Atomic Structure

• An atom has a nucleus containing protons (charge +1, mass 1) and neutrons (charge 0, mass 1).
• Electrons (charge -1, mass 1/2000) orbit the nucleus in shells.
• The mass number = protons + neutrons; atomic number = protons.
• In a neutral atom, number of protons = number of electrons.

Ions and Isotopes

• Ions have unequal protons and electrons; negative ions gain electrons, positive ions lose electrons.
• Isotopes are atoms with the same number of protons but different numbers of neutrons.

Relative Masses and Mass Spectrometry

• Relative atomic mass is the weighted mean mass of an atom compared to 1/12 of carbon-12.
• Mass spectrometry displays isotopes by mass/charge (m/z) ratio and their relative abundance.
• Relative atomic mass can be calculated as:
  (Σ (isotope abundance × isotope mass)) ÷ total abundance.

Electron Configuration and Orbitals

• Electrons fill shells and subshells in the order of increasing energy: s (2), p (6), d (10), f (14).
• Electron configuration is written as shell number + subshell letter + electron count (e.g., 1s²).
• The s orbital is spherical; p orbitals are dumbbell-shaped and oriented at right angles.
• Hund's Rule: Electrons fill orbitals singly before pairing due to repulsion.
• For transition metals, electrons may move between 4s and 3d for stability (e.g., chromium, copper).

Periodic Table Structure

• Elements are arranged by increasing proton number (atomic number).
• Groups (columns) share the same number of outer electrons; periods (rows) share the same number of shells.
• Blocks: s-block (groups 1, 2), p-block (groups 13-18), d-block (transition metals), f-block (lanthanides, actinides).

Evidence for Quantum Shells

• Emission spectra show discrete lines, proving electrons exist in quantized energy levels (shells).
• Each shell has fixed energy; electrons absorb or emit defined energy when moving between shells.

Ionization Energies and Trends

• Ionization energy: energy required to remove one mole of electrons from gaseous atoms.
• Ionization energy decreases down a group (more shielding, larger radius).
• Ionization energy increases across a period (greater nuclear charge, similar shielding, smaller radius).
• Successive ionization energies increase as electrons are removed from closer shells.

Periodic Trends

• Atomic radius decreases across a period (increased nuclear charge, same shell).
• Atomic radius increases down a group (extra shell added).
• Melting point trends: highest for giant structures (e.g., silicon); simple molecular substances have lower melting points.
• Metallic bonding strength increases with charge and number of delocalized electrons.

Key Terms & Definitions

• Isotope — atoms of the same element with different numbers of neutrons.
• Relative atomic mass (Ar) — weighted mean mass of an atom compared to 1/12 of carbon-12.
• Mass spectrometry — technique to determine isotopic composition and relative atomic mass.
• Quantum shell — discrete energy level where electrons reside.
• Ionization energy — energy required to remove an electron from a gaseous atom.
• Shielding — reduction in nuclear attraction due to inner electron shells.

Action Items / Next Steps

• Revise electron configuration rules and exceptions for transition metals.
• Practice calculating relative atomic mass from isotopic data.
• Review periodic trend explanations, focusing on atomic radius and ionization energy.
• Ensure clear memorization of all key definitions for exam preparation.