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: $4X + O<em>2 \rightarrow 2X</em>2O$
  • Example:
  • Lithium: $4Li + O<em>2 \rightarrow 2Li</em>2O$
  • Sodium: $4Na + O<em>2 \rightarrow 2Na</em>2O$
• Reactions with Water:
  • General Formula: $2X + 2H<em>2O \rightarrow 2XOH + H</em>2$
  • Example:
  • Lithium: $2Li + 2H<em>2O \rightarrow 2LiOH + H</em>2$
  • Sodium: $2Na + 2H<em>2O \rightarrow 2NaOH + H</em>2$

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: $X + Y<em>2 \rightarrow XY</em>2$
  • Example:
  • Magnesium and Bromine: $Mg + Br<em>2 \rightarrow MgBr</em>2$
  • Strontium and Chlorine: $Sr + Cl<em>2 \rightarrow SrCl</em>2$
• 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: $X + 2H<em>2O \rightarrow X(OH)</em>2 + H_2$
  • Example:
  • Beryllium: $Be + 2H<em>2O \rightarrow Be(OH)</em>2 + H_2$
  • Calcium: $Ca + 2H<em>2O \rightarrow Ca(OH)</em>2 + H_2$

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: $Cl<em>2(aq) + 2Br^-(aq) \rightarrow 2Cl^-(aq) + Br</em>2(aq)$
  • Bromine displaces iodide: $Br<em>2(aq) + 2I^-(aq) \rightarrow 2Br^-(aq) + I</em>2(aq)$

• Reaction with Water:

  • Chlorine undergoes disproportionation when reacting with water.
  • Full Equation:
    $Cl<em>2(g) + H</em>2O(l) \rightarrow HCl + HClO$
  • Ionic Equation:
    $Cl<em>2(g) + H</em>2O(l) \rightarrow 2H^+ + Cl^- + ClO^-$
  • Chlorine in HCl: -1, Chlorine in HClO: +1