Chemistry - Topic 1 - Atomic Structure

• Atom

  Mainly made up of protons and neutrons, contained in the middle, with electrons orbiting in shells
• Atom

  • Protons have a positive charge
  • Neutrons have a zero charge
  • Electrons have a negative charge
• Relative mass

  Protons and neutrons have a relative mass of 1, electrons have a relative mass of 1/2000
• Element

  Top number is the mass number (protons + neutrons), bottom number is the atomic number (protons)
• All atoms are neutral, with the number of protons equal to the number of electrons
• Ion


  Has a different number of electrons and protons, forming a stable charge
• Negative ion


  • O2- has gained 2 electrons
  • Na+ has lost 1 electron
• Isotopes


  Elements with the same number of protons but different numbers of neutrons
• Isotopes of carbon


  • Carbon-12
  • Carbon-13
  • Carbon-14
• History of atomic models

  1. Dalton - atoms are solid spheres

  2. Thomson - plum pudding model

  3. Rutherford - discovered nucleus

  4. Bohr - fixed energy levels and electron shells
• Time-of-flight mass spectrometer


  • Vaporizes sample
  • Ionizes sample
  • Accelerates ions
  • Ions drift at constant speed
  • Ions detected based on mass-to-charge ratio
• Relative atomic mass


  Average mass of an atom of an element, measured relative to carbon-12
• Relative molecular mass


  Average mass of a molecule, measured relative to carbon-12
• Relative isotopic mass


  Mass of an isotope of an element, measured relative to carbon-12
• Mass spectrum


  • Shows mass-to-charge ratio on x-axis
  • Shows abundance (percentage or relative) on y-axis
• Mass spectrum shows the isotopes present in an element
• Isotopic mass


  The mass of an isotope which makes it relatively straightforward if it had two electrons knocked off which would be quite rare
• If an isotope has two electrons knocked off

  The mass to charge ratio will be half as much
• If the abundance is a percentage, the total abundance of all isotopes must be 100%
• Most abundant isotope


  The most common isotope
• Calculating relative atomic mass

  1. Abundance of isotope A * Mass to charge ratio of A

  2. + Abundance of isotope B * Mass to charge ratio of B

  3. Divide by total abundance
• Molecular ion peak


  The last significant peak on a mass spectrum, shows the mass of the unfragmented molecule
• Electron subshells


  • S
  • P
  • D
  • F
• Electron subshells


  • S has 1 orbital, can hold 2 electrons
  • P has 3 orbitals, can hold 6 electrons
  • D has 5 orbitals, can hold 10 electrons
  • F has 7 orbitals, can hold 14 electrons
• Electron configuration


  Represented as 1s2 2s2 2p6 etc, where the number is the shell, the letter is the subshell, and the number is the electrons in that subshell
• Filling electron configurations

  1. Fill from lowest energy level upwards

  2. Fill orbitals singly first before pairing up
• Transition metal electron configurations

  Electrons can move from 4s to 3d orbitals to create more stable half-full or full 3d subshells
• Ionization


  The minimum energy required to remove one mole of electrons from one mole of atoms in the gaseous state
• Removing electrons from an ion

  1. Remove electrons from 4s first

  2. Then remove from 3D

  3. Check the numbers to ensure total matches
• Ionization energy


  Minimum amount of energy required to remove one mole of electrons from one mole of atoms in the gaseous state
• Ionization energy is always endothermic and has a positive value
• Shielding

  • The more shells or electron shells between the positive nucleus and the outer electron, the less energy is required and the weaker the attraction
• Atomic size

  • The bigger the atom, the further away the electrons are from the nucleus, the weaker the attractive force, and the less energy required to remove the outer electron
• Nuclear charge

  • The more protons in the nucleus, the bigger the attraction between the nucleus and the outer electron, and the more energy required to remove the electron
• Successive ionization energy

  The removal of more than one electron from the same atom
• There are distinctive jumps in successive ionization energies as electrons are removed from shells increasingly closer to the nucleus
• As you go down a group (e.g. group 2)

  First ionization energy decreases
• Reasons for decreasing first ionization energy down a group

  • Atomic radius increases, so electrons are further from nucleus and attractive force is weaker
  • Shielding increases, so attractive force between outer electrons and nucleus is weaker
• As you go across a period

  First ionization energy generally increases
• Reasons for increasing first ionization energy across a period

  • Nuclear charge increases, so attractive force between outer electrons and nucleus is stronger
  • Shielding is similar