Periodic trends

Cards (36)

  • Electrostatic attraction
    • Positive and negative charges attract
    • The force of attraction between two charges is described by Coulomb's law: F=k(q1q2/r^2), where q1 and q2 are the magnitudes of the charges, r is the distance between them, and k is a constant
    • The strength of any electrostatic attraction is directly proportional to the charges involved
    • The strength of any electrostatic attraction is inversely proportional to the distance between the charges squared (as the distance of the force of attraction decreases)
  • Core charge
    A measure of the attractive force between the valence electrons and the nucleus
  • Atomic radius
    The distance from the nucleus to the valence shell electrons
  • Ionisation energy
    • The process of moving an electron from an atom
    • Happens when the atom is given sufficient energy
    • The energy required to move one electron from an element in its gas phase
    • If sufficient energy is supplied, all other electrons can be ionised, and the energy required to do this is successive ionisation energy
  • Electronegativity
    The ability of an atom to attract electrons in a covalent bond
  • The atomic radius decreases
    Because the increase in nuclear charge pulls the electrons closer to the nucleus
  • Ionisation energy is the energy needed to move one electron
  • Ionisation energy increases because of increasing nuclear charge, which results in stronger attraction between the valence electrons and nucleus
  • Electrostatic force
    A force that acts between charges
  • Factors impacting electrostatic force
    • Quantity of charge (more charge = more force)
    • Distance between charges (more distance = less force)
  • Quantity of charge
    Directly proportional to electrostatic force
  • Distance between charges
    Inversely/indirectly proportional to electrostatic force
  • An atom of Lithium has 3p+, 4n°, 3e-
  • Large core charge + small distance
    Large electrostatic force
  • Smaller core charge + further away from nucleus
    Small electrostatic force
  • Full electron shells
    Shield valence electrons from the electrostatic force
  • Going down a group
    Electrons experience more shielding
  • Going left to right in a period
    Electrons experience more shielding
  • Atomic radius
    • Increases down the group (more shells)
    • Decreases going across the period (larger core charge)
  • More electrostatic force pulling electrons in
    Larger core charge
  • Ionisation energy
    The energy required to remove an electron from an atom
  • Going down the group
    Ionisation energy decreases
  • Going across the period
    Ionisation energy increases
  • Removing more electrons
    Higher ionisation energy required
  • Electronegativity
    The ability of an atom to attract an electron towards itself
  • Going down a group
    Electronegativity goes down (bigger core charge)
  • Going across a period
    Electronegativity increases (closer to the nucleus more force)
  • Fluorine is the most electronegative and has the highest ionisation energy
  • Francium is easy to remove an electron but hard to gain one
  • Electron affinity
    Change in energy of an atom when an electron is added to its valence shell
  • Going up the group
    Electron affinity increases (decreasing energy)
  • Going across a period
    Electron affinity increases (decreasing energy)
  • Chlorine has the highest electron affinity
  • Formation of ions
    Atoms with a positive or negative charge, trying to become the most stable (least electron affinity)
  • Overall net charge isn't 0 so there is imbalance in protons and electrons
  • The core charge of an atom of carbon is +6