Periodic Trends

Classification in the Periodic Table

    Metals

  • Properties: malleability, lustre, high melting points, high thermal/electrical conductivities, and tendency to lose electrons
  • Located primarily on the left-hand side of the periodic table

    Non-Metals

  • Properties: very brittle, dull, have lower melting points, low thermal/electrical conductivities, and tendency to gain electrons
  • Located primarily on the right-hand side of the periodic table

    Metalloids

  • Properties: intermediate between those of a metal and a non-metal
  • Located between metals and non-metals on the periodic table

Radius of Atoms and Ions

    Atomic Radius:

  • The radius of an element in its normal/neutral form
  • Increases down a group, decreases across a period
  • Effective Nuclear Charge: the net positive charge experienced by outer valence shell electrons, determined by the force of attraction between protons in the nucleus and the force of repulsion between inner core electrons
    • The effective nuclear charge increases across a period as the number of protons increases

    Ionic Radius:

  • The radius of an element in its ionized form
  • Positive Ions:
    • Smaller in size than their parent atom
    • Formed by a loss of electrons (typically all electrons) from outermost energy level
    • Increased attraction between nucleus and valence shell electrons
    • The greater the number of electrons removed, the smaller the ionic radius
  • Negative Ions:
    • Larger in size than their parent atom
    • Gain of electrons into the outermost energy level increases repulsion between electrons
    • Decreased attraction between nucleus and valence shell electrons
    • The greater the number of electrons gained, the larger the ionic radius

Trends in Properties of Elements

    Ionization Energy:

  • The energy required to remove the outermost electron from a species (atom, ion, molecule)
  • The first ionization energy generally decreases going down a group (less energy to remove an electron)
  • The first ionization energy generally increases going across a period (more energy to remove an electron)

    Electron Affinity:

  • The amount of energy released when a neutral atom gains an electron to form a negatively charged ion
  • The absolute value for first electron affinity generally decreases going down a group (lower affinity for gaining an electron)
  • The absolute value for first electron affinity generally increases going across a period (high affinity for gaining an electron)

    Electronegativity:

  • A measure of how much an atomic nucleus attracts the shared electrons that are involved in a covalent bond
  • Electronegativity values generally decrease down a group
  • Electronegativity values generally increase across a period

Elements in Groups 1 and 17

    Metallic Character:

  • The tendency of an element to lose electrons to form a positive ion
  • Lower ionisation energy results in a greater metallic character
  • Going down group 1, the metallic character increases

    Non-Metallic Character:

  • The tendency of an element to gain electrons to form a negative ion
  • Lower ionisation energy results in a greater metallic character
  • Going down group 17, the non-metallic character decreases

    Reactions of Group 1 Metals With Water

  • Going down group 1, elements have lower first ionisation energies due to the presence of additional energy levels
  • It is easier for elements with a lower ionisation energy to transfer their outer valence electrons to water, which increases their reactivity with water
  • Hence, reactivity with water increases going down group 1

Discontinuities in Ionisation Energy

Discontinuities in ionisation energy appear in two distinct regions on the periodic table when moving across a period:


    Between Groups 2 and 13

  • The first ionization energy of group 13 elements is generally lower than group 2 elements, going against the general trend
  • This is because group 13 elements have their outer valence electron in the p-sublevel, which is higher in energy than the s-sublevel where group 2 elements have their valence electrons
    • This makes the nucleus’ attraction weaker, and results in a lower ionization energy

    Between Groups 15 and 16

  • The first ionization energy of group 16 elements is generally lower than group 15 elements, going against the general trend
  • This is because group 15 elements have single electrons occupying each orbital within the p-sublevel, while group 16 elements have one orbital in the p-sublevel with double occupancy
    • This double occupancy results in an increased repulsion between valence electrons, which decreases the attraction between the valence electrons and the nucleus, and results in a lower ionization energy

Oxidation States

    What is an Oxidation State?

  • An oxidation state is a number assigned to an element in a compound or ion, which represents the hypothetical number of electrons lost or gained for that element
  • This is based on the assumption that an element's chemical bonds are fully ionic (electrons are completely transferred), which is not true for most substances

    Rules for Assigning Oxidation States

  • A pure element will have an oxidation state of 0
  • Monatomic (single atom) ions have an oxidation state equal to the charge of the ion
  • The oxidation state of all atoms in a neutral compound must sum up to give zero overall
  • The oxidation state of all atoms in a polyatomic ion must sum up to give the overall charge of the ion
  • The oxidation state of an atom is typically the charge of its most common ion
  • Assign positive oxidation numbers to the most metallic elements first
  • Assign negative oxidation numbers to the most non-metallic elements first

Transition Elements

    What are Transition Elements?

  • A transition element is an element with an incomplete d-sublevel that is able to form cations with an incomplete d-sublevel
  • Zinc has a filled d-sublevel both when it is an atom and an ion, hence it is not considered a transition element

    Properties of Transition Elements

  • Variable oxidation states
  • High melting points
  • Magnetic properties
  • Catalytic properties (speed up chemical reactions while remaining unchanged themselves)
  • Form coloured compounds
  • Form complex ions

    Variable Oxidation States in Transition Elements

  • Since the 4s and 3d-sublevels are close in energy, electrons from both of these sublevels are able to act as valence electrons
  • Electrons are lost first from the 4s-sublevel (this is why many transition metal ions form the +2 oxidation state)
  • In addition, electrons can also easily be lost from the 3d-sublevel without too much additional energy
  • Hence, transition elements can easily form ions with variable oxidation states as there are a number of d-electrons that can be lost in addition to the s-electrons

    Coloured Complexes of Transition Elements

  • Transition metal complex ions are formed when ligands donate lone pairs to the metal ion, creating coordination bonds
  • As ligands approach, their electrons repel the d-electrons in the metal ion, causing the d-orbitals to split into higher and lower energy levels
    • Electrons in the lower energy d-orbitals can absorb visible light to move to the higher energy orbitals
  • The energy difference (Δ) between these split d-levels corresponds to specific wavelengths of visible light that can be absorbed, and this splitting depends on factors like the metal ion, its oxidation state, and the type and arrangement of ligands
  • The wavelength absorbed determines the colour removed from white light, and the coloured complex appears as the complementary colour of the absorbed light
    • The colour wheel (in the data booklet) helps predict which colour is absorbed based on the observed colour of the complex