Properties of Period 3 Elements & Their Oxides & Chlorides (Oxford AQA International A Level Chemistry)

Na & Mg with Water

Sodium & Magnesium

• Both sodium, Na, and magnesium, Mg, are metals and are found in Group 1 and Group 2 of the periodic table respectively

• Both have high melting points, but magnesium has a higher melting point than sodium

  • This is because of the 2+ charge of magnesium, meaning that it is has a higher charge density

• Both are silvery metals

  • Sodium is quite a soft, silvery metal which tarnishes quickly in air

  • Magnesium is harder than sodium and you will often see it as magnesium ribbon

Reactions with water

• Despite their similarities, sodium and magnesium will react with water quite differently

• Sodium with cold water:

  2Na (s) + 2H₂O (l) → 2NaOH (aq) + H₂ (g)

• This is a very vigorous, exothermic reaction

• The sodium floats on the surface of the water fizzing rapidly and melting as a result of the heat produced during the reaction

• The colourless sodium hydroxide formed will have a pH of around 13-14, so a very alkaline solution is formed

• The oxidation state of the sodium changes from 0 in its elemental state, to +1 in the sodium hydroxide

• Magnesium with cold water:

Mg (s) + 2H₂O (l)  → Mg(OH)₂ (aq) + H₂ (g)

• This is an extremely slow reaction - only a very small number of bubbles will form on the magnesium ribbon

• The magnesium hydroxide formed will have a pH of around 10 - it is less alkaline than sodium hydroxide because magnesium hydroxide is only partially soluble

• This is the key component in 'milk of magnesia'

• The oxidation state of the magnesium changes from 0 in the elemental state, to +2 in the magnesium hydroxide

• Heated magnesium with steam:

Mg (s) + H₂O (g)  → MgO (s) + H₂ (g)

• This reaction is must faster than with cold water

• The magnesium burns with a bright, white flame

• The products of this reaction are different - magnesium oxide is produced instead of magnesium hydroxide

• The oxidation state of the magnesium changes from 0 in its elemental state, to +2 in the magnesium oxide

Period 3 Oxides

• The Period 3 elements, excluding chlorine and argon, combine with oxygen to form oxides

• The oxide formed will contain the elements in their highest oxidation state

Summary of reactions with oxygen

• Sulfur can actually form two oxides - SO₂ and SO₃

  • For SO₃ to form, a catalyst must be used and the reaction must take place at a very high temperature

  • The equation for this reaction is:

2S (s) + 3O₂ (g) → 2SO₃ (g)

Ionic Oxides

• Sodium oxide, magnesium oxide and aluminium oxide are ionic oxides because the bonding exists between metals and non metals

• They have giant lattice structures and thus, high melting points

Giant Covalent Oxides

• The graph then shows a giant covalent oxide

  • Silicon dioxide

• This is covalent because both silicon and oxygen are non metals

• The millions of covalent bonds within this giant structure are extremely strong, and thus it has a high melting point

• Giant covalent structures can also be called macromolecules or giant molecules

Simple Covalent Oxides

• The graph then shows a significant drop in melting point, as we reach the simple covalent oxide molecules

  • Phosphorus(V) oxide

  • Sulfur dioxide

  • Sulfur trioxide

• These are small molecules with only weak intermolecular forces between them

• Sulfur dioxide and sulfur trioxide are both gases at room temperature, because both their melting point and boiling point are so low

  • Sulfur trioxide, SO₃, has a slightly higher melting point than sulfur dioxide, SO₂, because of the increase in intermolecular forces between the slightly larger SO₃ molecules

Summary of trends in the physical properties of Period 3 oxides

Period 3 Chlorides

• Sodium to sulfur will react directly with chlorine when heated to form chlorides

Sodium

• Sodium burns in chlorine with a bright orange flame

  • White solid sodium chloride is produced

2Na (s) + Cl₂ (g) → NaCl (s)

Magnesium

• Magnesium burns with its usual intense white flame to give white magnesium chloride

Mg (s) + Cl₂ (g) → MgCl₂ (s)

Aluminium

• Aluminium burns in a stream of chlorine to produce very pale yellow aluminium chloride

3Al (s) + 3Cl₂ (g)→ 2AlCl₃ (s)

Silicon

• If chlorine is passed over silicon powder heated in a tube, it reacts to produce silicon tetrachloride

Si (s) + 2Cl₂ (g) → SiCl₄ (l)

Phosphorus

• White phosphorus burns spontaneously in chlorine to produce phosphorus(V) chloride

P₄ (s) + 10Cl₂ (g) → 4PCl₅ (s)

Sulfur

• When chlorine is passed over heated sulphur, it reacts to form an orange liquid, S₂Cl₂

2S (s) + Cl₂ (g) → S₂Cl₂ (l)

Melting and Boiling Point

• NaCl and MgCl₂ have high melting and boiling points

  • Large amounts of energy required to break the strong attraction between the positive and negative ions

• Aluminium chloride can exist as AlCl₃ and Al₂Cl₆ depending on the conditions

• They have different structures and properties

  • AlCl₃ exists at room temperature and is giant ionic structure

  • Al₂Cl₆ exists at higher temperatures and is a dimer with covalent compounds

• PCl₅ is a white solid which sublimes at 436 K

• SiCl₄ and S₂Cl₂ have low melting and boiling points

  • Weak intermolecular forces which require little energy to overcome

  • The higher the M_r of the molecule, the stronger the intermolecular forces

Summary table to show the bonding of Period 3 chlorides

*Aluminium chloride also forms a dimer, Al₂Cl₆ formed from covalent bonds