Electron Configurations

Line Spectra

    Spectroscopy and Light:

  • Sunlight is the same as white light, which is an electromagnetic wave
  • Spectroscopy is the study of the interaction between matter and light
  • Elements can be identified via their unique emission spectra

    Frequency and Wavelength of Electromagnetic Waves

  • Frequency (f): number of waves per second
  • Wavelength (λ): distance between two crests
  • Wavelength and frequency are inversely proportional

    Types of Spectra:

  • Continuous Spectrum: contains all frequencies/wavelengths across a range of electromagnetic radiation
  • Emission Line Spectrum: range of frequencies/wavelengths emitted during electron transition from higher to lower energy level
    • Lines converge at higher energy (high frequency)
    • True convergence only shown in hydrogen (single electron)

Electron Transitions

    Energy Levels:

  • The principal quantum number (n) determines the energy level
  • n=1 is closest to the nucleus (ground state)
  • Maximum number of electrons per energy level = 2(n)2

    Types of Transitions:

  • Absorption:
    • Low to high energy level (e.g. n=2 → n=3)
    • Creates dark lines in continuous spectrum
    • Endothermic (absorbs energy)
    • Occurs in cold gases
  • Emission:
    • High to low energy level (e.g. n=4 → n=3)
    • Creates bright lines in spectrum
    • Exothermic (releases energy)
    • Occurs in hot gases

    Spectral Series:

  • Energy transitions end at different levels:
    • n → n=1: UV radiation (highest energy)
    • n → n=2: Visible light (medium energy)
    • n → n=3: Infrared radiation (lowest energy)

Energy Sublevels

    Atomic Orbitals:

  • Region of space with a high probability of finding an electron
  • A single orbital can hold a maximum of 2 electrons
  • Main types of orbitals:
    • s: spherical shape (1 orbital) - max 2 electrons
    • p: dumbbell shape (3 orbitals) - max 6 electrons
    • d: 4-petal flower shape (5 orbitals) - max 10 electrons
    • f: complex shape (7 orbitals) - max 14 electrons

    Energy Level Distribution:

  • Each main energy level (n) is divided into specific sublevels
  • The orbital distribution follows a precise pattern:
    • nSublevelsspdfTotal orbitalsTotal e–
    11s100012
    22s, 2p130048
    33s, 3p, 3d1350918
    44s, 4p, 4d, 4f13571632

Electron Configuration Principles

    Key Principles:

  • Pauli Exclusion Principle: no more than 2 electrons can be found within a particular atomic orbital
  • Hund's Rule: all orbitals of equal energy (degenerate orbitals) must be occupied by one electron before any pairing occurs
    • This minimizes repulsion between electrons (more stable)
    • All single electrons must have the same spin
    • Paired electrons have opposite spins
  • Aufbau Principle: lower-energy orbitals are filled first (building-up principle)
    • Energy order: s < p < d < f
    • 4s fills before 3d (lower energy)
    • 4s empties before 3d (higher quantum number)

    Electron Configurations:

  • A way to show the occupied electrons orbitals for atoms and ions
  • For cations, electrons are removed from higher-energy orbitals first
  • Example configuration (aluminium 3+ ion):
    • Aluminium's electron configuration: 1s2 2s2 2p6 3s2 3p1
    • First remove electrons from 3p, then 3s
    • Final: 1s2 2s2 2p6 (3 electrons removed)
  • Abbreviated notation uses the previous noble gas as a reference
    • Carbon: [He] 2s2 2p2

    Special Cases:

  • Copper (Cu):
    • [Ar] 4s1 3d10 (NOT [Ar] 4s2 3d9)
    • More stable due to filled d-sublevel
  • Chromium (Cr):
    • [Ar] 4s1 3d5 (NOT [Ar] 4s2 3d4)
    • More stable due to half-filled d-sublevel

    Orbital Diagrams:

  • Diagrams representing electron configurations using arrows for electrons
  • Single-headed arrows (↑) are used for individual electrons
  • Opposite-facing arrows (↑↓) are used for paired electrons
  • Note: 4s fills before 3d due to energy level overlap

Ionisation Energy

    First Ionization Energy:

  • Energy required to remove one electron from a gaseous atom
  • Convergence limit at n = infinity
  • Influenced by nuclear structure:
    • Strong Nuclear Force: holds protons and neutrons together
    • Based on interactions between quarks in the nucleus

    Factors Affecting Ionization:

  • Nuclear Charge: higher charge = more energy needed
  • Distance From Nucleus: greater distance = less energy needed
  • Shielding Effect:
    • Inner shell electrons shield outer ones
    • More full shells = less ionization energy needed

    Ionization Energy Calculations:

  • The fundamental energy and frequency relationships are:
    • E = hc/λ = hf (for energy)
    • f = c/λ (for frequency)
    • Where: h is Planck's constant and c is the speed of light
  • The energy relationships show that:
    • Higher energy means higher frequency
    • Higher energy means shorter wavelength

Ionization Trends

    Periodic Trends:

  • Ionization energy increases across a period because:
    • Nuclear charge increases
    • Shielding effect increases
  • Ionization energy decreases down a group because:
    • Distance between outer electrons and nucleus increases

    Special Cases:

  • Beryllium to Boron (ionization energy decreases):
    • Boron has outer electrons in the 2p orbital
    • Beryllium only has outer electrons in the 2s orbital
    • p orbitals have higher energy than s orbitals
    • Hence electrons in boron's 2p orbital are easier to remove (less energy needed)
  • Nitrogen to Oxygen (ionization energy decreases):
    • Nitrogen has three single-occupied p orbitals (unpaired electrons)
    • Oxygen has one double-occupied p orbital (paired electrons)
    • Electron pairing in oxygen causes repulsion (easier to remove → less energy needed)

    Successive Ionization:

  • The process of removing multiple electrons in sequence
  • Each successive ionization requires more energy because:
    • The effective nuclear charge increases
    • The electrostatic attraction becomes stronger
    • Fewer electrons remain to provide shielding