Lecture 4: Periodicity and the Periodic Table

The Periodic Table Basics

  • The periodic table organizes elements based on their proton number.
  • Arrangement of the Periodic Table:
    • Periods: horizontal rows (left to right)
    • Groups: vertical columns (top to bottom)
  • Example Element: Magnesium (Mg), Proton Number: 12, Atomic Mass: 24.3

Periods in the Periodic Table

  • All elements in the same period share the same number of electron shells.
  • Example:
    • Elements in Period 3 have 3 electron shells.
    • Period structure:
    • Period 1: 1 Shell
    • Period 2: 2 Shells
    • Period 3: 3 Shells
    • Period 4: 4 Shells
    • Period 5: 5 Shells
    • Period 6: 6 Shells
    • Period 7: 7 Shells

Groups in the Periodic Table

  • All elements in the same group have the same number of electrons in their outer shell.
  • Group Number: Indicates the number of valence electrons.
  • Example:
    • Group 1 has 1 outer shell electron, Group 2 has 2, and so on.

Electron Configuration

  • The periodic table can be categorized into s, p, d, and f blocks which help determine electron configurations.
  • Example of Electron Configurations:
    • Silicon (Si): 1s^2 \, 2s^2 \, 2p^6 \, 3s^2 \, 3p^2
    • Germanium (Ge): 1s^2 \, 2s^2 \, 2p^6 \, 3s^2 \, 3p^6 \, 4s^2 \, 3d^{10} \, 4p^2

Atomic Radius Trends

  • Across a Period: Atomic radius decreases due to increased nuclear charge attracting electrons closer to the nucleus.
  • Down a Group: Atomic radius increases because additional electron shells are added, increasing distance from the nucleus.

Ionisation Energies

  • Across a Period: Ionisation energy generally increases due to higher nuclear charge holding electrons more tightly.
  • Down a Group: Ionisation energy decreases as outer electrons are further from the nucleus and experience more shielding, making them easier to remove.

Summary of Period Trends

  • Ionisation Energy increases across a period while Atomic Radius decreases.
  • Ionisation Energy decreases down a group while Atomic Radius increases.

Group 1 Elements (Alkali Metals)

  • Elements: Li, Na, K, Rb, Cs, Fr
  • Characteristics:
    • Metals, lose 1 electron to form +1 ions
    • Atomic Radius: Increases down the group
    • 1st Ionisation Energy: Decreases down the group
    • Melting Point: Generally decreases down the group
    • Reactivity: Increases down the group

Reactivity of Group 1 Elements

  • Reactions with Group 7:
    • General Formula: 2X + Y_2 \rightarrow 2XY
    • Example Reaction:
    • Potassium and Chlorine: 2K + Cl_2 \rightarrow 2KCl
    • Sodium and Iodine: 2Na + I_2 \rightarrow 2NaI
  • Reactions with Oxygen:
    • General Formula: 4X + O2 \rightarrow 2X2O
    • Example:
    • Lithium: 4Li + O2 \rightarrow 2Li2O
    • Sodium: 4Na + O2 \rightarrow 2Na2O
  • Reactions with Water:
    • General Formula: 2X + 2H2O \rightarrow 2XOH + H2
    • Example:
    • Lithium: 2Li + 2H2O \rightarrow 2LiOH + H2
    • Sodium: 2Na + 2H2O \rightarrow 2NaOH + H2

Group 2 Elements (Alkaline Earth Metals)

  • Elements: Be, Mg, Ca, Sr, Ba, Ra
  • Characteristics:
    • Metals, lose 2 electrons to form +2 ions
    • Atomic Radius: Increases down the group
    • 1st Ionisation Energy: Decreases down the group
    • Melting Point: Generally decreases down the group
    • Reactivity: Increases down the group

Reactivity of Group 2 Elements

  • Reactions with Group 7:
    • General Formula: X + Y2 \rightarrow XY2
    • Example:
    • Magnesium and Bromine: Mg + Br2 \rightarrow MgBr2
    • Strontium and Chlorine: Sr + Cl2 \rightarrow SrCl2
  • Reactions with Oxygen:
    • General Formula: 2X + O_2 \rightarrow 2XO
    • Example:
    • Beryllium: 2Be + O_2 \rightarrow 2BeO
    • Magnesium: 2Mg + O_2 \rightarrow 2MgO
  • Reactions with Water:
    • General Formula: X + 2H2O \rightarrow X(OH)2 + H_2
    • Example:
    • Beryllium: Be + 2H2O \rightarrow Be(OH)2 + H_2
    • Calcium: Ca + 2H2O \rightarrow Ca(OH)2 + H_2

Group 7 Elements (Halogens)

  • Elements: F₂, Cl₂, Br₂, I₂
  • Characteristics:
    • Non-metals, gain 1 electron to form -1 ions
    • Atomic Radius: Increases down the group
    • 1st Ionisation Energy: Decreases down the group
    • Melting Point: Generally increases down the group
    • Reactivity: Decreases down the group

Reactivity of Group 7 Elements

  • Displacement Reactions:

    • More reactive halogens can displace less reactive halogens in solution
    • Example:
    • Chlorine displaces bromide: Cl2(aq) + 2Br^-(aq) \rightarrow 2Cl^-(aq) + Br2(aq)
    • Bromine displaces iodide: Br2(aq) + 2I^-(aq) \rightarrow 2Br^-(aq) + I2(aq)

  • Reaction with Water:

    • Chlorine undergoes disproportionation when reacting with water.
    • Full Equation:
      Cl2(g) + H2O(l) \rightarrow HCl + HClO
    • Ionic Equation:
      Cl2(g) + H2O(l) \rightarrow 2H^+ + Cl^- + ClO^-
    • Chlorine in HCl: -1, Chlorine in HClO: +1