The Periodic Table: Chemical Periodicity (AS Level Inorganic Chemistry)

Welcome to the chapter on Chemical Periodicity! This topic is all about spotting patterns and explaining why elements behave the way they do as we move across the Periodic Table. It links fundamental concepts like atomic structure and bonding (which you learned previously) to macroscopic properties and reactivity. Mastering Period 3 is key, as these elements showcase nearly all the important trends!

Quick Review: What is Periodicity?

Periodicity refers to the repeating patterns in the chemical and physical properties of elements when they are arranged in order of increasing atomic number. We focus heavily on Period 3 elements: Na, Mg, Al, Si, P, S, Cl, Ar.

1. Periodicity of Physical Properties (Section 9.1)

1.1 Trends in Atomic Radius and Ionic Radius

As you move across Period 3 (from Sodium, Na, to Argon, Ar):

1. Atomic Radius Decreases

  • Why? Across a period, electrons are added to the same principal quantum shell (the third shell).
  • Simultaneously, the nuclear charge (number of protons) increases from +11 (Na) to +17 (Cl).
  • The increasing positive charge pulls the valence electrons closer to the nucleus, causing the atom to shrink.
  • Key Takeaway: Shielding remains roughly constant, but the effective nuclear charge increases significantly.

2. Ionic Radius (Cations) Decreases

  • Group 1, 2, and 13 elements (Na, Mg, Al) form positive ions (cations).
  • Formation of these ions involves losing the entire outer shell, meaning the ions are much smaller than their parent atoms.
  • Example: Na+ has only two shells, while Na has three.
  • The size decreases from Na+ to Mg2+ to Al3+ because the nuclear charge increases while the number of electrons (10 electrons) and shells remains the same. The pull on the remaining electrons gets stronger.

3. Ionic Radius (Anions) Increases

  • Group 15, 16, and 17 elements (P, S, Cl) form negative ions (anions).
  • Anions are much larger than their parent atoms because adding electrons increases electron-electron repulsion, causing the electron cloud to expand.
  • Cl- is the largest anion because it has the smallest nuclear charge for the same number of electrons (18 electrons).

Don't worry if this seems tricky at first: Remember the "footprint" analogy. Going across a period, you’re trying to squeeze more people (protons) into the same small room (electron shell). They pull everything inwards!

1.2 Trends in Melting Point and Electrical Conductivity

The trend in melting point is not smooth—it depends entirely on the structure and bonding of the element.

Trend in Melting Point (Na to Ar):

  1. Na, Mg, Al: Very High (Metallic Bonding) - Increases from Na to Al.
  2. Si: Extremely High (Giant Molecular/Covalent) - Peak MP.
  3. P, S, Cl, Ar: Very Low (Simple Molecular) - Decreases dramatically.

Explanation based on structure (9.1.2):

  • Na, Mg, Al (Metals): High MPs due to metallic bonding (electrostatic attraction between positive ions and delocalised electrons). MP increases because the cation charge increases (Na+ to Al3+), and the number of delocalised electrons increases (1 to 3), leading to stronger metallic bonds.
  • Si (Metalloid): The highest MP. Silicon has a giant molecular structure (macromolecular, like diamond) where every atom is held by strong covalent bonds. A huge amount of energy is needed to break these bonds.
  • P₄, S₈, Cl₂, Ar (Non-metals): Very low MPs. They exist as simple molecules (e.g., P₄, S₈, Cl₂). The bonds within the molecules are strong covalent bonds, but the forces between the molecules are only weak van der Waals' forces.
  • Melting point decreases from S₈ to P₄ to Cl₂ to Ar because the size of the simple molecules decreases (S₈ is the largest, Ar is monoatomic), leading to weaker instantaneous dipole-induced dipole forces.

Trend in Electrical Conductivity (9.1.2):

Conductivity relies on the presence of mobile charge carriers (delocalised electrons or free ions).

  • Na, Mg, Al: Good conductors. Conductivity increases from Na to Al because Al contributes 3 delocalised electrons per atom compared to 1 for Na.
  • Si: Semiconductor (poor conductor).
  • P, S, Cl, Ar: Non-conductors (Insulators). These elements do not have delocalised electrons or mobile ions.

Quick Review Box (Physical Trends):

Radius: Decreases (stronger nuclear pull).
Melting Point: High (Metallic) → Highest (Giant Covalent) → Low (Simple Molecular).
Conductivity: High (Metals) → Zero (Non-metals).

2. Periodicity of Chemical Properties: Oxides (Section 9.2)

Chemical periodicity focuses on how the bonding character changes across the period, moving from elements that form highly ionic bonds (metals) to those that form increasingly covalent bonds (non-metals).

2.1 Reactions with Oxygen (9.2.1)

All Period 3 elements (except Argon) react with oxygen, often strongly, to form oxides.

  • Na: \(4\text{Na} + \text{O}_2 \rightarrow 2\text{Na}_2\text{O}\) (Sodium oxide)
  • Mg: \(2\text{Mg} + \text{O}_2 \rightarrow 2\text{MgO}\) (Magnesium oxide)
  • Al: \(4\text{Al} + 3\text{O}_2 \rightarrow 2\text{Al}_2\text{O}_3\) (Aluminium oxide)
  • Si: \(\text{Si} + \text{O}_2 \rightarrow \text{SiO}_2\) (Silicon(IV) oxide)
  • P: \(\text{P}_4 + 5\text{O}_2 \rightarrow \text{P}_4\text{O}_{10}\) (Phosphorus(V) oxide)
  • S: \(\text{S} + \text{O}_2 \rightarrow \text{SO}_2\) (Sulfur(IV) oxide, forms SO₃ if excess O₂ and catalyst present)

Trend in Bonding and Oxidation Number (9.2.2):

  • Na₂O, MgO, Al₂O₃: These are ionic compounds (metal + oxygen).
  • SiO₂, P₄O₁₀, SO₂, SO₃: These are covalent compounds (non-metal + oxygen).
  • The maximum oxidation number generally increases across the period, corresponding to the number of outer shell electrons available for bonding (Group 1: +1, Group 13: +3, Group 17: +7).
2.2 Reactions of Oxides with Water and Acid/Base Behaviour (9.2.3 & 9.2.4)

The type of oxide determines its reaction with water and its acid/base character.

Oxide Bonding & Structure Reaction with Water Acid/Base Nature
Na₂O Ionic, Giant Dissolves readily: \(\text{Na}_2\text{O} + \text{H}_2\text{O} \rightarrow 2\text{NaOH}(\text{aq})\)
(Solution pH: 14, highly alkaline)		 Strongly Basic
MgO Ionic, Giant Reacts very slowly (sparingly soluble): \(\text{MgO} + \text{H}_2\text{O} \rightarrow \text{Mg}(\text{OH})_2(\text{aq})\)
(Solution pH: 9-10, weakly alkaline)		 Basic
Al₂O₃ Ionic/Covalent, Giant Insoluble (Does not react) Amphoteric (Reacts with both acids and bases)
SiO₂ Covalent, Giant Molecular Insoluble (Does not react) Weakly Acidic (Reacts only with strong bases)
P₄O₁₀ Covalent, Simple Molecular Dissolves violently: \(\text{P}_4\text{O}_{10} + 6\text{H}_2\text{O} \rightarrow 4\text{H}_3\text{PO}_4(\text{aq})\)
(Solution pH: 1-2, highly acidic)		 Strongly Acidic
SO₂ Covalent, Simple Molecular Dissolves readily: \(\text{SO}_2 + \text{H}_2\text{O} \rightleftharpoons \text{H}_2\text{SO}_3(\text{aq})\)
(Solution pH: 3-4, acidic)		 Acidic
SO₃ Covalent, Simple Molecular Dissolves readily: \(\text{SO}_3 + \text{H}_2\text{O} \rightarrow \text{H}_2\text{SO}_4(\text{aq})\)
(Solution pH: 1-2, highly acidic)		 Strongly Acidic

Key Trend: The acidity of the oxides increases across Period 3:

Basic → Amphoteric → Acidic

Why? Explanation using Bonding and Electronegativity (9.2.6):

  • Moving left, the element (E) is a metal (low electronegativity). The E–O bond is highly ionic. When dissolved, the oxide ion (\(\text{O}^{2-}\)) acts as a strong base, reacting with water to form \(\text{OH}^{-}\). This is basic behaviour. (e.g., Na₂O)
  • Moving right, the element (E) is a non-metal (high electronegativity). The E–O bond is highly covalent. When dissolved, the element (E) pulls electrons away from water, releasing \(\text{H}^{+}\) ions. This is acidic behaviour. (e.g., SO₃)
2.3 Amphoteric Behaviour of Aluminium Oxide (\(\text{Al}_2\text{O}_3\))

An amphoteric substance can act as both an acid and a base. Aluminium oxide is insoluble in water but reacts with strong acids and strong bases.

1. Reaction with Acid (acting as a base):
\(\text{Al}_2\text{O}_3(\text{s}) + 6\text{HCl}(\text{aq}) \rightarrow 2\text{AlCl}_3(\text{aq}) + 3\text{H}_2\text{O}(\text{l})\)

2. Reaction with Base (acting as an acid, though complex ions are not required at this level, recognition of the salt formation is):
\(\text{Al}_2\text{O}_3(\text{s}) + 2\text{NaOH}(\text{aq}) + 3\text{H}_2\text{O}(\text{l}) \rightarrow 2\text{Na}[\text{Al}(\text{OH})_4](\text{aq})\) (Sodium tetrahydroxoaluminate)

Did you know? This amphoteric property allows us to separate aluminium from its impurities in mining processes, as the aluminium oxide dissolves while the impurities (like iron oxide) do not.

3. Periodicity of Chemical Properties: Chlorides (Section 9.2)

3.1 Reactions with Chlorine (9.2.1)

Elements react with chlorine to form chlorides (e.g., NaCl, \(\text{MgCl}_2\), \(\text{AlCl}_3\), \(\text{SiCl}_4\), \(\text{PCl}_5\)).

Trend in Bonding:

  • NaCl, \(\text{MgCl}_2\): Highly Ionic.
  • \(\text{AlCl}_3\): Borderline/Covalent character (due to high charge density of \(\text{Al}^{3+}\) causing polarisation of the chloride ion).
  • \(\text{SiCl}_4\), \(\text{PCl}_5\): Highly Covalent.

Why the change in bonding? This is again due to electronegativity difference and the size/charge of the cation. As the metal ion charge increases and size decreases (from Na+ to Al3+), the positive ion has a stronger ability to pull the electron cloud of the anion towards itself, causing distortion (polarisation). This leads to increased covalent character, especially for \(\text{AlCl}_3\).

3.2 Reactions of Chlorides with Water (Hydrolysis) (9.2.5)

The behaviour of Period 3 chlorides in water is a vital piece of evidence for their bonding type.

  • NaCl, \(\text{MgCl}_2\): Dissolve without reacting chemically with water. They merely dissociate into ions.
    • NaCl: Neutral solution (\(\text{pH} \approx 7\)).
    • \(\text{MgCl}_2\): Slightly acidic solution (\(\text{pH} \approx 6\)) because the small, high-charge \(\text{Mg}^{2+}\) ion weakly hydrolyses water, but this reaction is very slight compared to Al.
  • \(\text{AlCl}_3\): Undergoes hydrolysis (reaction with water), producing dense white fumes of HCl gas when heated gently, and forming an acidic solution.
    • \(\text{AlCl}_3(\text{s}) + 3\text{H}_2\text{O}(\text{l}) \rightarrow \text{Al}(\text{OH})_3(\text{s}) + 3\text{HCl}(\text{g})\)
    • The resulting solution is acidic (\(\text{pH} \approx 3\)). This is due to the high charge density of the \(\text{Al}^{3+}\) ion, which strongly polarises coordinated water molecules, leading to the release of \(\text{H}^{+}\).
  • \(\text{SiCl}_4\): Hydrolyses vigorously (a violent reaction) to produce white fumes of HCl and solid silica.
    • \(\text{SiCl}_4(\text{l}) + 2\text{H}_2\text{O}(\text{l}) \rightarrow \text{SiO}_2(\text{s}) + 4\text{HCl}(\text{g})\)
    • Solution is highly acidic (\(\text{pH} \approx 1\)).
  • \(\text{PCl}_5\): Reacts violently and completely with water (hydrolysis).
    • \(\text{PCl}_5(\text{s}) + 4\text{H}_2\text{O}(\text{l}) \rightarrow \text{H}_3\text{PO}_4(\text{aq}) + 5\text{HCl}(\text{aq})\)
    • Solution is highly acidic (\(\text{pH} \approx 1\)).

Analogy: Think of the covalent chlorides (\(\text{SiCl}_4\), \(\text{PCl}_5\)) as molecular "grenades." They react completely with water because the bonds are polar covalent and the central atom has vacant, accessible d-orbitals to accept a lone pair from the water molecule, initiating the reaction. The ionic chlorides just dissolve.

Summary of Hydrolysis Trend:

Dissolve (Neutral/Slightly Acidic) → Hydrolyse Acidically (Strong)

Key Takeaway for Chlorides: The extent of hydrolysis (reaction with water) increases across the period, resulting in increasingly acidic solutions. This change is explained by the move from ionic bonding (Na, Mg) to covalent bonding (Si, P).

4. Predicting Periodicity (Section 9.3)

One crucial skill in periodicity is applying the Period 3 trends to other parts of the table, particularly Group 2 and Group 17 (which you will study in detail later).

4.1 Predicting Properties within a Group (9.3.1)

Properties typically change gradually down a group (often related to increasing atomic size and increasing shielding).

  • Atomic Radius: Increases down a group (more shells).
  • Ionisation Energy: Decreases down a group (outer electron is further away and better shielded).
  • Reactivity of Metals (e.g., Group 1, 2): Increases down a group (easier to lose valence electrons).
  • Reactivity of Non-metals (e.g., Group 17): Decreases down a group (harder to gain an electron).

Example: If you know that Mg is a basic oxide, you can predict that Group 2 elements below it (Ca, Sr, Ba) will also form basic oxides. Since atomic size increases down the group, BaO might be expected to be even more basic than MgO.

4.2 Deducing Identity from Data (9.3.2)

In exams, you may be given a set of data (melting points, conductivity, reaction pHs) for an unknown element (X) and asked to deduce its identity or position in the Periodic Table.

Step-by-Step Deduction:

  1. Analyze Structure/Bonding (from MP/Conductivity):
    • High MP & High Conductivity → Metal (Na, Mg, Al)
    • Very High MP & Poor Conductivity → Giant Covalent (Si)
    • Low MP & Zero Conductivity → Simple Molecular Non-metal (P, S, Cl)
  2. Analyze Oxide Reactivity (Acid/Base):
    • Oxide forms strong alkali with water → Group 1 metal (e.g., Na)
    • Oxide is amphoteric → Group 13 metal (e.g., Al)
    • Oxide is strongly acidic → Right side of the period (P, S, Cl)
  3. Analyze Chloride Reactivity (Hydrolysis):
    • Chloride dissolves and is neutral → High ionic character (Na)
    • Chloride hydrolyses rapidly, forming a very acidic solution → High covalent character (Si, P, S)

By combining these observations, you can usually narrow down the element's position precisely to a specific group or even identify it (e.g., "Element X must be silicon because it is non-conductive, has a very high melting point, and forms an acidic oxide").

Key Takeaway: Periodicity is the framework for predicting chemistry. Understand the underlying reasons (nuclear charge, size, electronegativity, and structure) and the predictions follow logically.