Chemistry (9620) | Periodicity

👋 Welcome to Periodicity: The Organization of Chemistry!

Hello! This chapter on Periodicity is where we step back and look at the entire Periodic Table, not just individual elements. It's like finding the hidden pattern in a complex dataset. Understanding periodicity helps us predict how elements will behave just by knowing their position. This is central to Inorganic Chemistry!

Don't worry if some of the trends seem tricky—we'll break down the explanations step-by-step, focusing on the core reasons related to structure and bonding.

3.2.1 Classification and Period 3 Trends (International AS)

Classification: s, p, d, and f Blocks

The Periodic Table is organized based on the proton number (Atomic number, Z). Elements are categorized into blocks depending on the type of sub-shell (orbital) that contains the highest-energy or 'last' electron added in the build-up principle.

• s-Block: Groups 1 and 2 (Alkali metals and Alkaline earth metals). Outer electrons are in an s-orbital.
• p-Block: Groups 13 to 18 (e.g., Halogens, Noble gases). Outer electrons are in a p-orbital.
• d-Block: Transition metals (the middle block). Outer electrons are in a d-orbital.
• f-Block: Lanthanides and Actinides (the two rows usually placed at the bottom). Outer electrons are in an f-orbital.

Quick Tip: The block tells you about the element's general properties. For example, s-block elements are highly reactive metals, while p-block elements show a greater variety, often containing non-metals.

3.2.1.2 Physical Properties of Period 3 Elements (Sodium to Argon)

We need to examine the trends of Period 3 (Na, Mg, Al, Si, P, S, Cl, Ar) in terms of atomic radius, first ionisation energy, and melting point.

Trend 1: Atomic Radius (Decreases Across the Period)

The atomic radius gets smaller as you move from left to right across Period 3.

Explanation:

1. Nuclear Charge Increases: The proton number (Z) increases from Na (11) to Ar (18). This means the nucleus has a stronger positive charge.
2. Shielding is Constant: All Period 3 elements have their valence electrons in the third shell. The inner electrons (in shells 1 and 2) provide the same amount of shielding effect against the increasing nuclear charge.
3. Net Effect: The stronger pull from the nucleus (higher charge) on the outer electrons, combined with constant shielding, pulls the electron shells closer, making the atomic radius smaller.

Analogy: Imagine a tug-of-war. The increasing nuclear charge is like adding stronger people to the pulling side. If the people on the defensive side (the inner electrons/shielding) stay the same, the rope (the outer shell) gets pulled closer to the stronger team.

Trend 2: First Ionisation Energy (FIE) (Generally Increases Across the Period)

First Ionisation Energy (FIE) is the energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous ions with a +1 charge:
$$X(g) \rightarrow X^+(g) + e^-$$

FIE generally increases across Period 3 because the electrons are held more strongly:

1. The atomic radius decreases (electrons are closer to the nucleus).
2. The nuclear charge increases (stronger pull).

The Anomalies (The Dips):
Don't worry! This is the key evidence for sub-shells (s and p orbitals). The overall trend is upwards, but there are two noticeable dips:

• Dip 1: Mg (\(3s^2\)) to Al (\(3s^2 3p^1\)): Al has a lower FIE than Mg.

  Reason: The electron being removed from Al is in the slightly higher energy 3p orbital, which is shielded by the inner 3s electrons. It takes less energy to remove an electron from this shielded, higher-energy p orbital than from the full 3s orbital of Mg.

• Dip 2: P (\(3p^3\)) to S (\(3p^4\)): S has a lower FIE than P.

  Reason: In Sulfur, the electron being removed is the first electron to be paired up in a 3p orbital. This paired electron experiences electron-electron repulsion from its partner, making it easier to remove compared to the single electrons in Phosphorus's half-filled p orbitals (which are more stable).

Trend 3: Melting Point (Varies Significantly)

Melting point depends entirely on the structure and bonding of the element.

1. Na, Mg, Al (Metals):
   • Structure: Metallic lattice.
   • Bonding: Metallic bonding (attraction between positive ions and a sea of delocalised electrons).
   • Trend: MP increases from Na to Al. This is because the charge on the ion increases ($Na^+$ to $Al^{3+}$) and the number of delocalised electrons increases (1 to 3), leading to stronger attraction and higher MP.
2. Si (Silicon):
   • Structure: Macromolecular (Giant Covalent).
   • Bonding: Strong covalent bonds in a giant lattice.
   • Result: Si has the highest melting point in the period. Huge amounts of energy are needed to break these strong covalent bonds.
3. P, S, Cl, Ar (Non-metals):
   • Structure: Simple Molecular structures (P₄, S₈, Cl₂, Ar).
   • Bonding: Very strong covalent bonds within the molecules, but only weak induced dipole–dipole forces (van der Waals) between the molecules.
   • Result: Low melting points. Only a small amount of energy is needed to overcome the weak intermolecular forces.
   • P₄ to S₈: S₈ has a higher MP than P₄ because the S₈ molecule is larger, resulting in stronger van der Waals forces.
   • Cl₂ and Ar: Very low MPs as they are very small molecules (Cl₂) or monatomic (Ar), resulting in extremely weak van der Waals forces.

3.2.2 Group 2: The Alkaline Earth Metals (International AS)

Group 2 elements (Mg, Ca, Sr, Ba, etc.) are called the Alkaline Earth Metals. They all lose their two outer s-electrons to form $M^{2+}$ ions.

Trends in Group 2

As you move down Group 2:

1. Atomic Radius increases: Due to increasing number of electron shells.
2. First Ionisation Energy (FIE) decreases: The outer electrons are further from the nucleus and experience greater shielding from the inner shells, making them easier to remove.
3. Melting Point: Generally decreases, although trends can be irregular. The metallic bonding weakens as the size of the $M^{2+}$ ion increases and the distance between the nucleus and the delocalised electrons increases.

Reactions of Group 2 Elements with Water

Group 2 metals react with water to form the metal hydroxide and hydrogen gas:
$$M(s) + 2H_2O(l) \rightarrow M(OH)_2(aq/s) + H_2(g)$$

• Magnesium (Mg): Reacts very slowly with cold water, but reacts rapidly with steam to form MgO and H₂.
• Calcium (Ca) onwards (Sr, Ba): React vigorously with cold water.

Trend: Reactivity increases down the group, because the FIE decreases, making it easier to lose electrons and become oxidized.

Trends in Solubility of Group 2 Compounds

This is one of the most important concepts for Group 2:

• Hydroxides ($M(OH)_2$): Solubility increases down the group (Mg(OH)₂ is sparingly soluble).
• Sulfates ($MSO_4$): Solubility decreases down the group (BaSO₄ is highly insoluble).

Memory Aid: "Sulfates Sink, Hydroxides Rise" in terms of solubility down the group.

Uses of Group 2 Compounds

• Mg(OH)₂: Used as an antacid (in medicine) to neutralize excess stomach acid.
• Ca(OH)₂ (Slaked lime): Used in agriculture to neutralize acidic soil, which is essential for healthy crops.
• CaO (Calcium oxide) or CaCO₃ (Calcium carbonate): Used in flue gas desulfurization. They react with the acidic pollutant sulfur dioxide ($SO_2$) emitted by power stations to remove it before release into the atmosphere.
• Magnesium: Used in the extraction of titanium from $TiCl_4$ at high temperatures:
  $$TiCl_4 + 2Mg \rightarrow Ti + 2MgCl_2$$

Testing for Sulfate Ions

Since Barium sulfate ($BaSO_4$) is insoluble, acidified Barium Chloride solution ($BaCl_2$) is used to test for sulfate ions ($SO_4^{2-}$). If sulfates are present, a white precipitate forms:
$$Ba^{2+}(aq) + SO_4^{2-}(aq) \rightarrow BaSO_4(s)$$

Why is the $BaCl_2$ solution acidified?

The solution must be acidified (usually with dilute hydrochloric acid, $HCl$) to remove carbonate ($CO_3^{2-}$) and sulfite ($SO_3^{2-}$) ions. If these ions were present, they would also form white precipitates with $Ba^{2+}$ (e.g., $BaCO_3$), leading to a false positive result for the sulfate test. Acidification converts these interfering ions into soluble or gaseous products.

Did you know? Barium sulfate is also used in medicine as a "Barium meal" before X-rays of the digestive system. Because $BaSO_4$ is completely insoluble, it is non-toxic even though other Barium salts are highly poisonous.

3.2.3 Group 7: The Halogens (International AS)

Group 7 elements (F, Cl, Br, I, At) are the halogens. They are highly reactive non-metals that typically gain one electron to form $X^-$ ions.

Trends in Physical Properties

1. Boiling Point increases down the group:

   Halogens exist as simple diatomic molecules ($F_2$, $Cl_2$, $Br_2$, $I_2$). As you move down the group, the number of electrons increases, leading to larger molecules and stronger induced dipole–dipole forces (van der Waals forces) between them. More energy is required to overcome these forces, hence the higher boiling point.

   Observation: Fluorine and Chlorine are gases, Bromine is a liquid, Iodine is a solid.

2. Electronegativity decreases down the group:

   Electronegativity is the power of an atom to attract the pair of electrons in a covalent bond. Down the group, the atomic radius increases and the shielding increases. This means the nucleus is less able to attract the bonding electrons, decreasing electronegativity. (Fluorine is the most electronegative element.)

Trend in Oxidising Ability

Halogens are oxidising agents (they accept electrons).

• Trend: Oxidising ability decreases down the group. F₂ is the strongest oxidising agent.
• Explanation: Down the group, it becomes harder to attract an extra electron due to increasing atomic size and shielding.
• Displacement Reactions: A more reactive (stronger oxidising) halogen can displace a less reactive (weaker oxidising) halide ion from solution.
  • Chlorine displaces bromide: $$Cl_2(aq) + 2Br^-(aq) \rightarrow 2Cl^-(aq) + Br_2(aq)$$
  • Chlorine displaces iodide: $$Cl_2(aq) + 2I^-(aq) \rightarrow 2Cl^-(aq) + I_2(aq)$$

Trend in Reducing Ability

Halide ions ($X^-$) are reducing agents (they donate electrons).

• Trend: Reducing ability increases down the group. I⁻ is the strongest reducing agent.
• Explanation: The larger the ion, the weaker the attraction between the nucleus and the outer electrons, making it easier for the ion to donate an electron.
• Reaction with Concentrated Sulfuric Acid ($H_2SO_4$):
  • $NaCl$ or $NaF$: Only an acid-base reaction occurs (no redox).
    $$NaX(s) + H_2SO_4(l) \rightarrow NaHSO_4(s) + HX(g)$$
  • $NaBr$: HBr is strong enough to reduce some $H_2SO_4$ to $SO_2$. You see dense white fumes of HBr and brown fumes of $Br_2$.
  • $NaI$: HI is a much stronger reducing agent and can reduce $H_2SO_4$ all the way to $H_2S$ (rotten egg smell), producing $I_2$.

Identifying Halide Ions

Halide ions ($Cl^-$, $Br^-$, $I^-$) are identified using acidified silver nitrate solution ($AgNO_3$). A precipitate of the silver halide forms:
$$Ag^+(aq) + X^-(aq) \rightarrow AgX(s)$$

The precipitates are:

• $AgCl$: White precipitate
• $AgBr$: Cream precipitate
• $AgI$: Yellow precipitate

The Solubility in Aqueous Ammonia: To confirm which halide is present, you test the solubility of the precipitate in aqueous ammonia:

• $AgCl$: Dissolves easily in dilute ammonia.
• $AgBr$: Dissolves sparingly in concentrated ammonia.
• $AgI$: Insoluble in concentrated ammonia.

Why is the $AgNO_3$ acidified?
Similar to the sulfate test, the solution is acidified with dilute nitric acid ($HNO_3$) to prevent false positives. If carbonate ions ($CO_3^{2-}$) were present, they would form a white precipitate of $Ag_2CO_3$, which could be confused with $AgCl$. Acidification removes $CO_3^{2-}$.

3.2.3.2 Uses of Chlorine and Chlorate(I)

Chlorine is vital for water treatment to kill bacteria (disinfectant).

1. Reaction of Chlorine with Water:

   Chlorine undergoes a disproportionation reaction (where the same element is both oxidized and reduced):
   $$Cl_2(aq) + H_2O(l) \rightarrow HClO(aq) + HCl(aq)$$
   $HClO$ (chloric(I) acid, or hypochlorous acid) contains the chlorate(I) ion ($ClO^-$), which is the active disinfectant.

2. In Sunlight:

   In bright sunlight, the $HClO$ breaks down, releasing oxygen:
   $$2Cl_2(aq) + 2H_2O(l) \rightarrow 4HCl(aq) + O_2(g)$$

3. Reaction of Chlorine with Cold, Dilute Aqueous NaOH:

   This also produces a disproportionation product, used in household bleach:
   $$Cl_2(aq) + 2NaOH(aq) \rightarrow NaCl(aq) + NaClO(aq) + H_2O(l)$$
   The solution contains sodium chlorate(I) ($NaClO$), which is the key component of bleach.

Societal Balance: Chlorine is toxic, but the health benefits of clean water (preventing diseases like cholera) vastly outweigh the risks associated with its low concentration in drinking water.

3.2.4 Properties of Period 3 Elements and their Oxides and Chlorides (International A2)

This section brings together the trends in bonding, structure, and reactivity to explain what happens when Period 3 oxides and chlorides react with water. The resulting pH reveals the acidic or basic nature of these compounds.

1. Structure and Melting Point of Period 3 Oxides and Chlorides

The bonding changes from ionic on the left side of the period to covalent/molecular on the right side.

Oxides (Na₂O to SO₃)

The melting point trend follows the structure type:

• Na₂O, MgO, Al₂O₃: Have high melting points due to strong ionic bonds in a giant lattice (Al₂O₃ has some covalent character).
• SiO₂: Extremely high melting point because it has a giant covalent structure (macromolecular), requiring huge energy to break the strong covalent network.
• P₄O₁₀, SO₂, SO₃: Have low melting points because they are simple molecular structures, held together by weak intermolecular forces.

Chlorides (NaCl to PCl₅)

• NaCl, MgCl₂: High melting points (strong ionic bonding).
• Al₂Cl₆ (dimer), SiCl₄, PCl₅: Low melting points (simple molecular structures, covalent bonding).

Note: Aluminium chloride exists as a dimer, $Al_2Cl_6$, when solid or gaseous, exhibiting covalent character, but it tends to be regarded as having significant ionic character when molten. It sublimes easily.

2. Reactions with Water and pH

(a) Basic Oxides (Na₂O, MgO)

These are ionic and dissolve/react with water to form alkaline solutions (high pH, basic).

• Sodium Oxide: Dissolves easily, highly alkaline (pH 14). $$Na_2O(s) + H_2O(l) \rightarrow 2NaOH(aq)$$
• Magnesium Oxide: Sparingly soluble, but the small amount that dissolves is alkaline (pH 9-10). $$MgO(s) + H_2O(l) \rightarrow Mg(OH)_2(s/aq)$$

(b) Amphoteric Oxide (Al₂O₃) and Neutral Covalent Oxide (SiO₂)

These show transitional behaviour:

• Aluminium Oxide ($Al_2O_3$): Insoluble in water. It is amphoteric, meaning it reacts with both acids and bases (but not water directly).
• Silicon Dioxide ($SiO_2$): Insoluble in water. It has a giant covalent structure. It is chemically acidic but only reacts with very strong bases like hot, concentrated $NaOH$.

(c) Acidic Oxides (P₄O₁₀, SO₂, SO₃)

These are covalent molecules and react vigorously with water to form strong acids (low pH, acidic).

• Phosphorus(V) Oxide ($P_4O_{10}$): Forms phosphoric acid (a reasonably strong acid, pH ~1-2). $$P_4O_{10}(s) + 6H_2O(l) \rightarrow 4H_3PO_4(aq)$$
• Sulfur Dioxide ($SO_2$): Forms sulfurous acid (a weak acid). $$SO_2(g) + H_2O(l) \rightleftharpoons H_2SO_3(aq)$$
• Sulfur Trioxide ($SO_3$): Forms sulfuric acid (a very strong acid, low pH). $$SO_3(g) + H_2O(l) \rightarrow H_2SO_4(aq)$$

3. Reactions of Period 3 Chlorides with Water

This section clearly demonstrates the change from ionic to covalent bonding.

(a) Ionic Chlorides (NaCl, MgCl₂)

• NaCl: Simply dissolves (hydrates the ions). The resulting solution is neutral (pH 7).
• MgCl₂: Dissolves. The small, highly charged $Mg^{2+}$ ion can slightly polarise water molecules, leading to a very slightly acidic solution (pH 6-7).

(b) Covalent Chlorides (Al₂Cl₆, SiCl₄, PCl₅)

These chlorides hydrolyse (react strongly with water) to produce fumes of $HCl(g)$, forming highly acidic solutions (low pH).

• Aluminium Chloride ($Al_2Cl_6$): Hydrolyses vigorously, forming $Al(OH)_3$ (white precipitate) and $HCl$ gas. $$Al_2Cl_6(s) + 6H_2O(l) \rightarrow 2Al(OH)_3(s) + 6HCl(g)$$
• Silicon Tetrachloride ($SiCl_4$): Hydrolyses violently in water, producing $SiO_2$ (white solid smoke) and $HCl$ gas. $$SiCl_4(l) + 2H_2O(l) \rightarrow SiO_2(s) + 4HCl(g)$$
• Phosphorus(V) Chloride ($PCl_5$): Reacts violently, producing phosphoric acid and $HCl$ gas. $$PCl_5(s) + 4H_2O(l) \rightarrow H_3PO_4(aq) + 5HCl(g)$$

Summary of Period 3 Oxide/Chloride Acidity Trend:

Ionic/Basic (Na, Mg) $\rightarrow$ Amphoteric/Giant Covalent (Al, Si) $\rightarrow$ Covalent/Acidic (P, S, Cl)

This trend is explained by the change in bonding: Ionic compounds form basic solutions; highly covalent compounds react to form acidic solutions.

Encouragement: Periodicity synthesises much of the basic structure and bonding knowledge you learned previously. Focus on the core reasons—nuclear charge, shielding, and structure type—and the rest will click into place!