Periodicity of chemical properties of the elements in Period 3 - Chemistry (9701) - Cambridge A-Level

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🧠 Mastering Periodicity: Period 3 Elements (Na to Ar)

Hello future Chemists! This chapter is incredibly important because it takes all those fundamental atomic structure concepts (like nuclear charge and shielding) and applies them to real-world trends. Period 3 elements (Sodium to Argon) are the perfect group to study because they show dramatic and predictable changes in properties as we move across the table.

Understanding these trends—known as periodicity—will help you predict the behaviour of any unknown element! Let's dive into how structure, bonding, and electrons dictate everything.

Quick Reminder: The Elements of Period 3

The elements are: Na, Mg, Al, Si, P, S, Cl, Ar. They are filling the third electron shell, meaning they all have the same amount of electron shielding from inner shells (the 1s²2s²2p⁶ electrons).

9.1 Periodicity of Physical Properties

9.1.1 Atomic and Ionic Radii

Trend in Atomic Radius

Across Period 3, the atomic radius decreases.

The Explanation: As you move from Na (Z=11) to Ar (Z=18), the number of protons (the nuclear charge) increases.
The number of occupied electron shells remains the same (three shells), so the shielding effect remains approximately constant.
The increasing nuclear charge pulls the outer electrons closer to the nucleus, resulting in a smaller atomic radius.

Trend in Ionic Radius

The trend here is trickier because we compare two different types of ions: positive ions (cations) on the left and negative ions (anions) on the right.

1. Cations (Na⁺, Mg²⁺, Al³⁺):
These ions are much smaller than their parent atoms because they have lost their entire outer (third) shell. They also decrease in size from Na⁺ to Al³⁺ because the nuclear charge increases (11+ to 13+) pulling the remaining electrons (10 electrons) tighter.

2. Anions (P^3-, S^2-, Cl^-):
These ions are much larger than their parent atoms because they have gained electrons, increasing inter-electron repulsion within the third shell, causing the cloud to swell. They decrease in size from P^3- to Cl^- because the nuclear charge increases.

Key Takeaway for Radii: The primary driver for decreasing size across the period is the increasing nuclear charge acting on electrons in the same principal quantum shell.

9.1.2 Melting Point and Structure/Bonding

The variation in melting points is complex because Period 3 elements exhibit four different fundamental types of structure and bonding. Melting points generally increase sharply to Silicon and then decrease dramatically.

Element	Structure/Bonding Type	Melting Point Trend	Explanation
Na, Mg, Al	Giant Metallic Lattice	Increases (Na < Mg < Al)	Bonds are strong. Strength increases because the number of delocalised electrons per atom increases (1 to 3), and the positive ion charge increases (1+ to 3+), leading to stronger electrostatic attraction.
Si	Giant Molecular (Covalent Network)	Highest MP	Every atom is held by four strong covalent bonds. Massive energy is required to break these bonds.
P, S, Cl	Simple Molecular	Low (S > P > Cl)	Low energy required to overcome weak van der Waals' forces between molecules. S₈ has the highest MP because it is the largest molecule (more electrons, stronger van der Waals' forces) compared to P₄ and Cl₂.
Ar	Simple Monatomic	Lowest MP	Only very weak instantaneous dipole–induced dipole forces between single atoms.

Analogy: Think of melting point as breaking the construction. Si is a massive, complex skyscraper (covalent network), while Ar is a single, detached Lego block (weak forces).

9.1.3 Electrical Conductivity

Electrical conductivity depends entirely on the presence of mobile charged particles (usually delocalised electrons or ions).

Across Period 3, conductivity increases sharply then drops to zero.

Element	Conductivity	Explanation
Na, Mg, Al	Excellent conductors	Contain mobile delocalised electrons. Conductivity increases from Na to Al because the number of delocalised electrons contributed per atom increases (1 to 3).
Si	Semiconductor	Has a giant covalent structure, but the energy gap between valence and conduction bands is small enough for some electrons to move at higher temperatures.
P, S, Cl, Ar	Non-conductors (Insulators)	All electrons are fixed in covalent bonds or filled shells (Ar). There are no mobile charge carriers.

Quick Review of Physical Properties:
1. Radius: Decreases (due to higher Z, same shielding).
2. Melting Point: High (Metallic) -> Very High (Covalent Network) -> Low (Molecular).
3. Conductivity: High (Metals) -> Low (Si) -> Zero (Non-metals).

9.2 Periodicity of Chemical Properties (Reactions)

The chemical properties depend on how the elements react, which is governed by the electronegativity and the tendency to achieve a stable electronic configuration.

9.2.1 Trend in Oxidation Number in Oxides and Chlorides

Across Period 3, the maximum oxidation state generally increases, reflecting the number of valence electrons available for bonding.

Oxides: The maximum oxidation state shown by the elements in their oxides increases from +1 (Na) to +7 (Cl).

Na₂O (+1)
MgO (+2)
Al₂O₃ (+3)
SiO₂ (+4)
P₄O₁₀ (+5)
SO₃ (+6)
Cl₂O₇ (+7)

Note: Sulfur dioxide (SO₂, S is +4) and Phosphorus pentachloride (PCl₅, P is +5) show that elements in Period 3 can expand their octet due to vacant d orbitals, unlike Period 2 elements.

Chlorides: The maximum oxidation state in chlorides also follows this trend.

NaCl (+1)
MgCl₂ (+2)
AlCl₃ (+3)
SiCl₄ (+4)
PCl₅ (+5)

Explanation (9.2.2): The increase in oxidation number corresponds to the number of outer shell electrons available for bonding. For the metals, this results in the formation of ions; for the non-metals, this involves sharing electrons, often utilizing energetically accessible vacant 3d orbitals (for Si, P, S, Cl) to expand the octet.

9.2.2 Reactions of Elements with Oxygen, Chlorine, and Water (9.2.1)

The reactions become progressively less vigorous across the period, especially towards water.

Element	Reaction with Oxygen	Reaction with Chlorine	Reaction with Water
Na	Burns strongly: \(4\text{Na} + \text{O}_2 \rightarrow 2\text{Na}_2\text{O}\)	Vigorous reaction: \(2\text{Na} + \text{Cl}_2 \rightarrow 2\text{NaCl}\)	Vigorous, rapid reaction at room temp: \(2\text{Na} + 2\text{H}_2\text{O} \rightarrow 2\text{NaOH} + \text{H}_2\)
Mg	Burns brightly: \(2\text{Mg} + \text{O}_2 \rightarrow 2\text{MgO}\)	Vigorous reaction: \(\text{Mg} + \text{Cl}_2 \rightarrow \text{MgCl}_2\)	Very slow in cold water, but reacts with steam: \(\text{Mg} + \text{H}_2\text{O} \rightarrow \text{MgO} + \text{H}_2\)
Al	Burns readily: \(4\text{Al} + 3\text{O}_2 \rightarrow 2\text{Al}_2\text{O}_3\)	Vigorous reaction: \(2\text{Al} + 3\text{Cl}_2 \rightarrow 2\text{AlCl}_3\)	No reaction (protected by inert Al₂O₃ layer)
Si	Requires heat: \(\text{Si} + \text{O}_2 \rightarrow \text{SiO}_2\)	Requires heat: \(\text{Si} + 2\text{Cl}_2 \rightarrow \text{SiCl}_4\)	No reaction
P	Burns easily: \(\text{P}_4 + 5\text{O}_2 \rightarrow \text{P}_4\text{O}_{10}\)	Vigorous reaction: \(\text{P}_4 + 10\text{Cl}_2 \rightarrow 4\text{PCl}_5\)	No reaction
S	Burns easily: \(\text{S} + \text{O}_2 \rightarrow \text{SO}_2\) (can form SO₃)	Reaction possible	No reaction
Cl	Forms complex oxides	N/A	Reacts slightly: \(\text{Cl}_2 + \text{H}_2\text{O} \rightleftharpoons \text{HCl} + \text{HOCl}\)

9.2.3 & 9.2.4 Acid/Base Behaviour of Oxides and Reactions with Water

This is the most crucial chemical trend! When dissolved in water (or reacted with acid/base), Period 3 oxides show a transition from Basic to Amphoteric to Acidic behaviour.

Memory Aid: Think BAA - Basic, Amphoteric, Acidic.

1. Basic Oxides (Metals: Na₂O, MgO)

These oxides are ionic. The O^2- ion reacts strongly with water to form OH^- ions, increasing the pH (alkaline solution).

Sodium Oxide (Na₂O): Very soluble, strongly basic.

\(\text{Na}_2\text{O}(\text{s}) + \text{H}_2\text{O}(\text{l}) \rightarrow 2\text{NaOH}(\text{aq})\) (pH ~ 14)

Magnesium Oxide (MgO): Sparingly soluble, weakly basic.

\(\text{MgO}(\text{s}) + \text{H}_2\text{O}(\text{l}) \rightarrow \text{Mg(OH)}_2(\text{aq})\) (pH ~ 9)

2. Amphoteric/Insoluble Oxides (Metalloids: Al₂O₃, SiO₂)

These oxides have stronger covalent character and are insoluble in water, but they react with both strong acids and strong bases.

Aluminium Oxide (Al₂O₃): Insoluble, Amphoteric.

Reacts with strong acid (e.g., HCl):
\(\text{Al}_2\text{O}_3(\text{s}) + 6\text{HCl}(\text{aq}) \rightarrow 2\text{AlCl}_3(\text{aq}) + 3\text{H}_2\text{O}(\text{l})\)
Reacts with strong base (e.g., hot NaOH):
\(\text{Al}_2\text{O}_3(\text{s}) + 2\text{NaOH}(\text{aq}) + 3\text{H}_2\text{O}(\text{l}) \rightarrow 2\text{Na}[\text{Al(OH)}_4](\text{aq})\) (Sodium tetrahydroxoaluminate)

Silicon Dioxide (SiO₂): Insoluble, slightly acidic. Reacts only with very strong bases (like NaOH).

3. Acidic Oxides (Non-Metals: P₄O₁₀, SO₂, SO₃)

These oxides are covalent molecules. When reacted with water, they hydrolyse completely to form strong acids, resulting in low pH solutions.

Phosphorus(V) Oxide (P₄O₁₀): Highly reactive with water, forms phosphoric acid.

\(\text{P}_4\text{O}_{10}(\text{s}) + 6\text{H}_2\text{O}(\text{l}) \rightarrow 4\text{H}_3\text{PO}_4(\text{aq})\) (pH ~ 2)

Sulfur Dioxide (SO₂): Dissolves, forms sulfurous acid (weak acid).

\(\text{SO}_2(\text{g}) + \text{H}_2\text{O}(\text{l}) \rightleftharpoons \text{H}_2\text{SO}_3(\text{aq})\) (pH ~ 3)

Sulfur Trioxide (SO₃): Highly reactive with water, forms sulfuric acid.

\(\text{SO}_3(\text{l}) + \text{H}_2\text{O}(\text{l}) \rightarrow \text{H}_2\text{SO}_4(\text{aq})\) (pH ~ 1)

9.2.5 Reactions of Chlorides with Water

The behaviour of Period 3 chlorides in water also shifts from simple dissolution to violent hydrolysis, which is directly related to the change in bonding (9.2.6 & 9.2.7).

Chloride	Structure/Bonding	Reaction with Water	pH of Solution	Equation (if hydrolysis occurs)
NaCl	Giant Ionic	Dissolves (simple hydration)	7 (Neutral)	N/A (Just dissolves)
MgCl₂	Giant Ionic (with some covalent character)	Dissolves. Mg²⁺ ion is small and polarising, causing slight hydrolysis.	~6.5 (Very slightly acidic)	\([\text{Mg}(\text{H}_2\text{O})_6]^{2+}(\text{aq}) + \text{H}_2\text{O}(\text{l}) \rightleftharpoons [\text{Mg}(\text{H}_2\text{O})_5(\text{OH})]^+\text{H}^+(\text{aq})\)
AlCl₃	Covalent (as Al₂Cl₆)	Hydrolyses violently, fumes of HCl given off. The small, highly charged Al³⁺ ion is strongly polarising.	~3 (Strongly acidic)	\(\text{AlCl}_3(\text{s}) + 3\text{H}_2\text{O}(\text{l}) \rightarrow \text{Al(OH)}_3(\text{s}) + 3\text{HCl}(\text{aq})\)
SiCl₄	Simple Molecular Covalent	Violent hydrolysis, white fumes of HCl are visible.	~2 (Very acidic)	\(\text{SiCl}_4(\text{l}) + 4\text{H}_2\text{O}(\text{l}) \rightarrow \text{Si}(\text{OH})_4(\text{s}) + 4\text{HCl}(\text{aq})\)
PCl₅	Simple Molecular Covalent	Violent hydrolysis, forming phosphoric acid and HCl gas.	~1 (Very acidic)	\(\text{PCl}_5(\text{s}) + 4\text{H}_2\text{O}(\text{l}) \rightarrow \text{H}_3\text{PO}_4(\text{aq}) + 5\text{HCl}(\text{aq})\)

Why do the covalent chlorides hydrolyse violently? (9.2.6)

Covalent chlorides like SiCl₄ and PCl₅ hydrolyse easily because the central atom (Si or P) has vacant d orbitals (3d) that can accept a lone pair of electrons from a water molecule (acting as a nucleophile). This forms an unstable intermediate which then breaks down, releasing H⁺ ions (HCl gas) and leading to a highly acidic solution.

9.2.6 & 9.2.7 Explaining Trends using Bonding and Electronegativity

We can tie all these trends together by looking at the change in bonding:

1. Sodium to Aluminium: The bonding is predominantly ionic (metal + non-metal), resulting in basic properties. However, as the metal ion charge increases (Na⁺ to Al³⁺) and ion size decreases, the power of the ion to distort the electron cloud of the anion (its polarising power) increases. This pulls the bonding towards being more covalent. AlCl₃ shows clear covalent properties (dimer Al₂Cl₆, lower melting point, violent hydrolysis).

2. Silicon to Chlorine: These compounds are formed between two non-metals, leading to covalent bonding (simple molecular or giant molecular).

Deducing Bonding from Properties (9.2.7)

If you are given physical properties, you can deduce the bonding type:

High MP, conducts electricity: Giant Metallic (Na, Mg, Al)
Extremely High MP, does not conduct: Giant Molecular Covalent (Si)
Low MP, does not conduct, hydrolyses violently in water: Simple Molecular Covalent (SiCl₄, PCl₅, SO₃)
Dissolves in water, gives neutral/basic solution: Giant Ionic (NaCl, Na₂O, MgO)

Final Summary: The Great Shift in Period 3

Period 3 demonstrates a perfect transition across the Periodic Table:

Left Side (Na, Mg, Al): Dominantly Metallic/Ionic properties.
Strong conductivity, high melting points (except Na), basic oxides.

Middle (Si, Al): Transitional / Amphoteric/Covalent Network properties.
Si has the highest MP, Al₂O₃ is amphoteric, and AlCl₃ shows covalent character.

Right Side (P, S, Cl, Ar): Dominantly Non-metallic/Covalent Molecular properties.
No conductivity, low melting points, highly acidic oxides (which hydrolyse violently).