Periodicity of Chemical Properties of the Elements in Period 3
Quick Overview
- Period 3 elements show a clear pattern in how they react with oxygen, chlorine, and water.
- Their oxides and chlorides vary in bonding, acidity/basicity, and structure.
- These trends can be explained by electronic structure, oxidation number, and bonding type.
Full Notes
Reactions with Oxygen, Chlorine, and Water
Reactions with Oxygen
| Element | Oxide Formed | Type of Oxide | Example Equation |
|---|---|---|---|
| Na | Na2O | Basic, ionic | 4Na + O2 → 2Na2O |
| Mg | MgO | Basic, ionic | 2Mg + O2 → 2MgO |
| Al | Al2O3 | Amphoteric | 4Al + 3O2 → 2Al2O3 |
| Si | SiO2 | Acidic, giant covalent | Si + O2 → SiO2 |
| P | P4O10 | Acidic, molecular | P4 + 5O2 → P4O10 |
| S | SO2 | Acidic, molecular | S + O2 → SO2 |
Reactions with Chlorine
| Element | Chloride Formed | Example Equation |
|---|---|---|
| Na | NaCl | 2Na + Cl2 → 2NaCl |
| Mg | MgCl2 | Mg + Cl2 → MgCl2 |
| Al | AlCl3 | 2Al + 3Cl2 → 2AlCl3 |
| Si | SiCl4 | Si + 2Cl2 → SiCl4 |
| P | PCl5 | PCl3 + Cl2 → PCl5 or P4 + 10Cl2 → 4PCl5 |
Reactions with Water
| Element | Behaviour | Equation | pH |
|---|---|---|---|
| Na | Vigorous | 2Na + 2H2O → 2NaOH + H2 | ~13–14 |
| Mg | Slow (cold) / Fast (steam) |
Mg + H2O → Mg(OH)2 + H2 (cold, slow) Mg + H2O → MgO + H2 (steam) |
~9–10 |
Oxidation Numbers in Oxides and Chlorides
Oxidation numbers increase across Period 3 as more valence electrons are used in bonding.
Oxides:
- Na2O: Na = +1
- MgO: Mg = +2
- Al2O3: Al = +3
- SiO2: Si = +4
- P4O10: P = +5
- SO2: S = +4
- SO3: S = +6
Chlorides:
- NaCl: Na = +1
- MgCl2: Mg = +2
- AlCl3: Al = +3
- SiCl4: Si = +4
- PCl5: P = +5
The increase in oxidation number reflects increasing use of outer electrons in bonding across the period.
Reactions of Oxides with Water and pH of Solutions
| Oxide | Reaction with Water | Product | pH |
|---|---|---|---|
| Na2O | Na2O + H2O → 2NaOH | Strongly alkaline solution | ~13–14 |
| MgO | MgO + H2O → Mg(OH)2 | Weakly alkaline solution | ~9–10 |
| Al2O3 | Insoluble | — | — |
| SiO2 | Insoluble | — | — |
| P4O10 | P4O10 + 6H2O → 4H3PO4 | Acidic solution | ~1–2 |
| SO2 | SO2 + H2O → H2SO3 | Acidic solution | ~2–3 |
| SO3 | SO3 + H2O → H2SO4 | Strongly acidic solution | ~0–1 |
Basic oxides (Na2O, MgO) release O2− ions, which form OH− in water.
Acidic oxides (P4O10, SO2, SO3) form oxacids.
Amphoteric and giant covalent oxides (Al2O3, SiO2) are insoluble in water.
Acid/Base Behaviour of Oxides and Hydroxides
Basic oxides (e.g. Na2O, MgO): React with acids to form salt + water.
Acidic oxides (e.g. SO2, P4O10): React with alkalis to form salts.
Amphoteric oxides (e.g. Al2O3): React with both acids and bases.
Hydroxides:
| Hydroxide | Acid/Base Behaviour |
|---|---|
| NaOH | Strong base |
| Mg(OH)2 | Weak base |
| Al(OH)3 | Amphoteric — reacts with acids and bases |
Reactions of Chlorides with Water
| Chloride | Behaviour in Water | pH of Solution |
|---|---|---|
| NaCl | Dissolves, no hydrolysis | Neutral (~7) |
| MgCl2 | Slight hydrolysis | Slightly acidic (~6) |
| AlCl3 | Hydrolyses to form HCl | Acidic (~3) |
| SiCl4 | Violent hydrolysis, forms HCl + SiO2 | Acidic (~2–3) |
| PCl5 | Hydrolyses to form HCl + H3PO4 | Acidic (~1–2) |
Covalent chlorides (like SiCl4 and PCl5) react vigorously with water due to hydrolysis.
Ionic chlorides (like NaCl) do not react, simply dissolving.
Explaining Trends: Bonding and Electronegativity
As you move across Period 3 from sodium (Na) to chlorine (Cl), electronegativity increases, and this affects both the type of bonding and the structure of the elements' oxides and chlorides.
Key Idea:
As electronegativity increases across the period:
- Bonding shifts from ionic to covalent (electronegativity difference decreases).
- Structures shift from giant to molecular.
- Oxides become more acidic.
- Chlorides change from soluble salts to reactive covalent compounds.
This explains why:
- Sodium oxide forms alkaline solutions.
- Silicon dioxide is insoluble but reacts with both acids and bases.
- Sulfur dioxide forms an acidic solution in water.
Left side (Na to Al):
- Metals with low electronegativity.
- Their oxides are basic and usually ionic.
- Their chlorides are also ionic, and when added to water, they simply dissolve without reacting.
Middle (Si):
- Forms giant covalent structures like SiO2.
- Its oxide does not dissolve in water, but can react with both acids and bases (amphoteric).
Right side (P to Cl):
- Non-metals with higher electronegativity.
- Their oxides are acidic, and their chlorides are covalent.
- Many chlorides react with water to form acidic solutions.
Predicting Bond Types from Properties
We can use observed chemical and physical properties to deduce bonding in a Period 3 chloride or oxide:
| Observation(s) | Likely Bonding Type |
|---|---|
| High melting point, soluble in water and conducts electricity when molten | Ionic |
| Low melting point (gas or liquid at room temperature), unable to conduct electricity when molten | Simple covalent |
| Insoluble and high melting point | Giant covalent |
Examples:
NaCl → Ionic (solid, high melting point, conductive in solution)
SiO2 → Giant covalent (very high melting point, insoluble)
SO2 → Molecular covalent (gas, low melting point, acidic solution)
Summary
- Period 3 elements react with oxygen, chlorine, and water in predictable ways.
- Oxidation numbers in oxides and chlorides increase across the period.
- Basic oxides form alkaline solutions; acidic oxides form acids; amphoteric oxides react with both acids and bases.
- Hydroxides show differences in acid/base strength (NaOH strong base, Mg(OH)2 weak, Al(OH)3 amphoteric).
- Covalent chlorides hydrolyse in water to give acidic solutions; ionic chlorides simply dissolve.
- Trends explained by increasing electronegativity: bonding shifts from ionic to covalent, structures from giant to molecular, oxides more acidic, chlorides more reactive.