The Periodic Table

General

• In the Periodic Table the elements are arranged on order of increasing atomic number.
• The elements in a period show trends in physical and chemical properties that are repeated across the period.
• The similarity in properties of elements n a group is due to them havinf the same number of outer shell electrons.

Trends across Periods 2 and 3

• Electron Configuration:
  • Across the period successive elements have one more outer shell electron.
• Atomic radius:
  • Decreases going across the period:
    • Bigger nuclear charge
    • Same shielding
    • Nuclear attraction increases
• 1st ionisation energy:
  • Increases going across the period:
    • Bigger nuclear charge 
    • Same shielding
    • Nuclear attraction increases

Trends Across Periods 2 and 3 Continued

• Melting and Boiling Points:
  • Na, Mg and Al have metallic bond (high melting point).
  • Melting point increases from Na to Mg to Al because metal ion has greater charge.
  • Si has very high melting point (giant covalent structure)
  • P₄, S₈ and Cl₂ have low melting points (simple covalent molecule)
  • Ar has a very low melting points (fewest electrons, weakest van der Waal's forces).

Trends Down Groups

• Atomic Radius:
  • It increases going down the group:
    • Extra electron shell
    • Outer electron further from nucleus
    • More shielding
    • Increased nuclear charge (outweighs shielding and distance)
• 1st Ionisation Energy:
  • It decreases going down the group:
    • Extra electron shell
    • Outer electron further from the nucleus
    • More shielding
    • Increased nuclear charge (outweighs shielding and distance)

Redox Reactions of Group 2 Metals

• Group 2 elements undergo redox reactions with water and oxygen.

Water

• Fizzing will be seen (gas released).
• The group 2 eleetn will dissolve.

Example

Ca(s) + 2H₂O(l) ----> Ca(OH)₂(aq) + H₂(g)

Ca:  0 to +2 = oxidation                  

H:   +1 to 0 = reduction

Redox Reactions of Group 2 Metals - Oxygen

• Mg burns with a bright white flame.
• A white powder will be produced.

2Mg + O₂ --> 2MgO

Mg:   0 to +2 = oxidation                  

O:     0 to -2 = reduction

Trends in reactivity

• Reactivity increases down the group.
• The elements lose 2 outer electrons in these reactions.
• Down the group the outer electrons are further from nucleus and more shielded.
• Nuclear attraction is reduced, so electrons are lost more easily.
• Increased nuclear charge outweighed by greater shielding and distance.

Reactions of Group 2 Compounds - Oxides

• Adding water (in drops) to CaO produces calcium hydroxide, Ca(OH)₂(s).
• In excess water, calcium hydroxide dissolves to make limewater Ca(OH)₂(aq) whose pH is about 12.

Equation

 CaO(s) + H₂O(l) ---> Ca(OH)₂(s)

Reactions of Group 2 Compounds - Carbonates

• Group 2 carbonates decompose when heated.
• Going down the Group, the carbonates become harder to decompose (i.e. they become more stable.)

Example Equation

MgCO₃(s) ---> MgO(s) + CO₂(g)

Reactions of Group 2 Compounds - Hydroxides

• Group 2 hydroxides are alkaline and can be used to neutralise acids.
• Calcium hydroxide (Ca(OH)₂), can be used to reduce soil acidity in soil.
• Magnesium hydroxide (Mg(OH)₂), is found in milk of magnesia.

Boiling Points of the Group 7 Elements

• At room temperature:
  • Cl₂ is a pale green gas.
  • Br₂ is a brown liquid.
  • I₂ is a blue-black solid.
• In Group 7, melting and boiling point increases down the group:
  • The molecules have more electrons.
  • Stronger van der Waal's forces.

Displacement Reactions

• Reactivity decreases down the group
• This can be shown when the halogen elements higher up displace elements further down the group.

Colours of the Halogens

• In water:
  • Cl₂ - pale green
  • Br₂ - orange
  • I₂ - brown
• In a organic solvent:
  • Cl₂ - pale green
  • Br₂ - orange
  • I₂ - violet

Disproportionation reactions - Chlorine and Water

• Chlorine undergoes a redox reaction with water.
• Cl₂(g) + H₂O(l) --> HCl(aq) + HOCl(aq)
• This reaction is used in water purification.
• The risks of using chlorine:
  • Chlorine gas is toxic.
  • The formation of chlorinated hydrocarbon is possible.
• The benefits using chlorine:
  • Killing bateria outweighs the risks.

Disproportionation reactions - Chlorine and Dilute Sodium Hydroxide

• Chlorine undergoes a redox reaction with cold, dilute sodium hydroxide solution.
• This reaction is used to make bleach.
• Cl₂(g) + 2NaOH(aq) --> NaCl(aq) + NaOCl(aq) +H₂O(l)
• In both reactions the oxidation state of Cl changes from 0 to -1(Cl^-) and 0 to +1(OCl^-).
• Cl is oxidised and reduced, this is an example of disproportionation.

Reactions of Halide Ions (Cl-, Br- and I-)

• Chloride ions, bromide ions and iodide ions produce colored precipitates with silver nitrate solution (AgNO₃(aq)).
• Ag⁺(aq) + Cl^-(aq) --> AgCl(s) - white precipitate
• Ag⁺(aq) + Br^-(aq) --> AgBr(s) - cream precipitate
• Ag⁺(aq) + I^-(aq) --> AgI(s) - yellow precipitate

Compounds in Ammonia Solution

• AgCl(s) dissolves in dilute ammonia solution, NH₃(aq).
• AgBr(s) dissolves in concentrated ammonia solution but not in dilute ammonia.
• AgI(s) does not dissolve, even in concentrated ammonia solution.