Atomic Structure

• Model for atomic structure has evolved over time as knowledge and scientific understanding changes
• Plum Pudding Model:

  • Initially thought atoms consisted of a sphere of positive charge with small negative charges distributed evenly within it
• Electron Shell Model:

  • Atom consists of a small, dense central nucleus surrounded by orbiting electrons in electron shells
  • Discovered in the Rutherford scattering experiment in 1911
  • Nucleus consists of protons and neutrons giving it an overall positive charge and almost the entire mass of the atom
  • In a neutral atom, number of electrons is equal to number of protons due to relative charges
• Particle:

  • Proton: Relative charge +1, Relative mass 1
  • Neutron: Relative charge 0, Relative mass 1
  • Electron: Relative charge -1, Relative mass 1/1840
• Mass number is represented using A and can be calculated as the sum of protons and neutrons in an atom

  • Atomic number is represented using Z and is equal to the number of protons in an atom
• Relative atomic mass (Ar) is defined as the mean mass of an atom of an element, divided by one twelfth of the mean mass of an atom of the carbon-12 isotope

  • Isotopes are atoms of the same element with the same atomic number, but with a different number of neutrons, resulting in a different mass number
  • Neutral atoms of isotopes will react chemically in the same way as their proton number and electron configuration is the same, but they have different physical properties
• Ions are formed when an atom loses or gains electrons, meaning it is no longer neutral and will have an overall charge
• Mass Spectrometry is an analytical technique used to identify different isotopes and find the overall relative atomic mass of an element

  • Time of Flight (TOF) Mass Spectrometry records the time it takes for ions of each isotope to reach a detector
• Electrons are held in clouds of negative charge called orbitals

  • Different types of orbitals: s, p, d, and f
  • Each orbital can hold a different number of electrons before the next one is filled:
  • s-orbital = 2 electrons
  • p-orbital = 6 electrons
  • d-orbital = 10 electrons
  • Energy of the orbitals increases from s to d, meaning the orbitals are filled in this order
• Within an orbital, electrons pair up with opposite spin for stability

  • Three rules for writing out electron configurations:

  1. The lowest energy orbital is filled first

  2. Electrons with the same spin fill up an orbital first before pairing begins

  3. No single orbital holds more than 2 electrons
• Ionisation energy is defined as the minimum energy required to remove one mole of electrons from one mole of atoms in a gaseous state, measured in kJmol-1

  • Successive ionisation energies occur when further electrons are removed, requiring more energy as electrostatic force of attraction increases
  • First ionisation energy trends within the Periodic Table:
  • Along a Period, first ionisation energy increases
  • Down a Group, first ionisation energy decreases
• First ionisation energy of Aluminium is lower than expected due to a single pair of electrons with opposite spin, causing a natural repulsion which reduces the energy needed to remove the outer electron