Chemistry (9620) | Properties of Period 3 elements and their oxides and chlorides

Chemistry (9620) Study Notes: Period 3 Elements and Their Compounds

Hello future chemists! This chapter is super important because it brings together everything you learned about bonding, structure, and trends in the Periodic Table. We’re going to look at the elements in Period 3 (Sodium to Argon) and see how their chemistry changes dramatically as we move across the table. It’s like watching a chemical personality shift from aggressive metal to inert gas!

Understanding these trends is key to mastering Inorganic Chemistry. Don't worry if some concepts look tricky—we'll break them down using simple language and clear steps. Let's dive in!

Key Concept: The Changing Nature of Bonding Across Period 3

As you move from left (Group 1, Sodium) to right (Group 17, Chlorine) in Period 3, the elements change from strong metals to non-metals. This change affects the type of bonding they form in their compounds:

• Left Side (Na, Mg, Al): Form compounds with highly ionic bonding.
• Middle (Si): Forms compounds with giant covalent structures.
• Right Side (P, S, Cl): Form compounds with simple molecular (covalent) structures.

1. Reactions of Period 3 Elements with Water

We only need to focus on Sodium (Na) and Magnesium (Mg) for their reactions with water.

Sodium (Na) with Water

Sodium is a highly reactive metal. It reacts instantly and vigorously with cold water.

• Reaction: Very vigorous, melting the sodium into a ball which darts around the surface.
• Products: Hydrogen gas and a highly alkaline solution of sodium hydroxide.
• Equation: \( 2\text{Na}(\text{s}) + 2\text{H}_2\text{O}(\text{l}) \rightarrow 2\text{NaOH}(\text{aq}) + \text{H}_2(\text{g}) \)
• pH: High (around pH 13 or 14).

Magnesium (Mg) with Water

Magnesium is less reactive than sodium.

• Cold Water: Reacts very slowly, eventually forming magnesium hydroxide.
• Steam (Hot Water): Reacts rapidly and brightly.
• Equation (Cold Water): \( \text{Mg}(\text{s}) + 2\text{H}_2\text{O}(\text{l}) \rightarrow \text{Mg}(\text{OH})_2(\text{aq}) + \text{H}_2(\text{g}) \)
• Equation (Steam): \( \text{Mg}(\text{s}) + \text{H}_2\text{O}(\text{g}) \rightarrow \text{MgO}(\text{s}) + \text{H}_2(\text{g}) \)
• pH: Moderate (around pH 9 to 10), as magnesium hydroxide, $\text{Mg}(\text{OH})_2$, is only sparingly soluble.

2. Properties and Reactions of Period 3 Oxides

Period 3 elements react with oxygen to form specific oxides:
\(\text{Na}_2\text{O}, \text{MgO}, \text{Al}_2\text{O}_3, \text{SiO}_2, \text{P}_4\text{O}_{10}, \text{SO}_2, \text{SO}_3\)

2.1. Explaining Melting Point Trends (Structure and Bonding)

The melting point across the period is determined entirely by the structure and bonding of the oxide.

1. High Melting Points (Na₂O, MgO): These are ionic compounds held in a strong giant ionic lattice. A lot of energy is required to overcome these powerful electrostatic forces.
2. Peak Melting Point (Al₂O₃, SiO₂): $\text{Al}_2\text{O}_3$ is mostly ionic but with some covalent character. $\text{SiO}_2$ (Silicon Dioxide, like sand) has the highest melting point because it has a giant molecular (covalent) structure. Every silicon atom is strongly bonded to four oxygen atoms in a huge network.
3. Low Melting Points (P₄O₁₀, SO₂, SO₃): These are simple molecular covalent compounds. Only weak intermolecular forces (like van der Waals) need to be broken when they melt or boil, resulting in very low melting points. They are often gases or volatile solids at room temperature.

Memory Trick: Think of a skyscraper being built. Ionic bonds are like brick walls; Giant Covalent is like solid rock (very strong); Simple Molecular is like LEGO bricks (easy to separate).

2.2. Trend in Acid-Base Character of Oxides

When oxides react with water, their resultant solution's pH shows a distinct trend: Basic $\rightarrow$ Amphoteric $\rightarrow$ Acidic.

A) Basic Oxides (Na₂O and MgO)

• Bonding: Strong ionic.
• Reaction with Water: They dissolve/react to form metal hydroxides (alkalis).
• Equations:
  \( \text{Na}_2\text{O}(\text{s}) + \text{H}_2\text{O}(\text{l}) \rightarrow 2\text{NaOH}(\text{aq}) \) (Forms highly alkaline solution, pH 13-14)
  \( \text{MgO}(\text{s}) + \text{H}_2\text{O}(\text{l}) \rightarrow \text{Mg}(\text{OH})_2(\text{s}) \) (Reacts slowly, sparingly soluble. Forms moderately alkaline solution, pH 9-10)
• Reaction with Acids: They neutralize acids (e.g., $\text{Na}_2\text{O} + 2\text{HCl} \rightarrow 2\text{NaCl} + \text{H}_2\text{O}$).

B) Amphoteric Oxide (Al₂O₃)

• Amphoteric means it can react both as an acid and as a base.
• Reaction with Water: Insoluble.
• Reaction with Acid (as a Base):
  \( \text{Al}_2\text{O}_3(\text{s}) + 6\text{H}^{+}(\text{aq}) \rightarrow 2\text{Al}^{3+}(\text{aq}) + 3\text{H}_2\text{O}(\text{l}) \)
• Reaction with Base (as an Acid):
  \( \text{Al}_2\text{O}_3(\text{s}) + 2\text{OH}^{-}(\text{aq}) + 3\text{H}_2\text{O}(\text{l}) \rightarrow 2[\text{Al}(\text{OH})_4]^{-}(\text{aq}) \)

C) Neutral Oxide (SiO₂)

• Structure: Giant molecular covalent lattice.
• Reaction with Water: Insoluble and inert. No reaction, so no pH change.

D) Acidic Oxides ($\text{P}_4\text{O}_{10}, \text{SO}_2, \text{SO}_3$)

• Bonding: Covalent.
• Reaction with Water: They react readily to form acidic solutions (low pH, typically pH 0–2).
• Phosphorus(V) oxide ($\text{P}_4\text{O}_{10}$): Reacts vigorously to form phosphoric acid ($\text{H}_3\text{PO}_4$).
  \( \text{P}_4\text{O}_{10}(\text{s}) + 6\text{H}_2\text{O}(\text{l}) \rightarrow 4\text{H}_3\text{PO}_4(\text{aq}) \)
  (The resulting anion is the phosphate ion, $\text{PO}_4^{3-}$)
• Sulfur dioxide ($\text{SO}_2$): Forms sulfurous acid ($\text{H}_2\text{SO}_3$).
  \( \text{SO}_2(\text{g}) + \text{H}_2\text{O}(\text{l}) \rightleftharpoons \text{H}_2\text{SO}_3(\text{aq}) \)
  (The resulting anion is the sulfite ion, $\text{SO}_3^{2-}$)
• Sulfur trioxide ($\text{SO}_3$): Reacts violently to form sulfuric acid ($\text{H}_2\text{SO}_4$).
  \( \text{SO}_3(\text{g}) + \text{H}_2\text{O}(\text{l}) \rightarrow \text{H}_2\text{SO}_4(\text{aq}) \)
  (The resulting anion is the sulfate ion, $\text{SO}_4^{2-}$)

3. Properties and Reactions of Period 3 Chlorides

Period 3 elements react with chlorine to form:
\(\text{NaCl}, \text{MgCl}_2, \text{Al}_2\text{Cl}_6, \text{SiCl}_4, \text{PCl}_5\)

Notice the shift again: $\text{NaCl}$ and $\text{MgCl}_2$ are ionic. $\text{Al}_2\text{Cl}_6$ exists as a covalent dimer (two $\text{AlCl}_3$ units joined) at standard conditions, showing the transition. $\text{SiCl}_4$ and $\text{PCl}_5$ are simple molecular covalent.

3.1. Melting Point Trends in Chlorides

The trend here follows the bonding:

1. High Melting Points (NaCl, MgCl₂): These are ionic compounds requiring high energy to break the strong ionic lattice.
2. Low Melting Points (Al₂Cl₆, SiCl₄, PCl₅): These are covalent molecular compounds. Their melting points are very low because only weak van der Waals forces between molecules need to be broken. $\text{Al}_2\text{Cl}_6$ sublimes (turns directly from solid to gas) at relatively low temperatures.

3.2. Reactions of Chlorides with Water (Hydrolysis)

When chlorides are added to water, their behavior (and the resulting pH) tells us about their bonding:

A) Ionic Chlorides (NaCl and MgCl₂)

• Bonding: Highly ionic.
• Reaction with Water: They dissolve easily (dissociation) but do not hydrolyse (i.e., they don't react chemically with the water molecule itself).
• pH: Neutral ($\text{NaCl}$) or near neutral ($\text{MgCl}_2$).

B) Chlorides that Hydrolyse (Al₂Cl₆, SiCl₄, PCl₅)

These covalent chlorides react vigorously with water. This process is called hydrolysis. The central atom ($\text{Al}, \text{Si}, \text{P}$) has empty orbitals and acts as an electron acceptor, making it susceptible to attack by the lone pair on the oxygen atom in water ($\text{H}_2\text{O}$).

Aluminium Chloride ($\text{Al}_2\text{Cl}_6$ or $\text{AlCl}_3$):

Although covalent, the $\text{Al}^{3+}$ ion formed has a high charge-to-size ratio. This powerful charge pulls electrons away from the $\text{O}-\text{H}$ bonds of the coordinated water molecules, making the $\text{H}$ atoms acidic (easy to release as $\text{H}^{+}$).

• Observation: Solution is strongly acidic, often producing white fumes ($\text{HCl}$) when reacting violently.
• Equation (showing acidic nature in water):
  \( [\text{Al}(\text{H}_2\text{O})_6]^{3+}(\text{aq}) \rightleftharpoons [\text{Al}(\text{H}_2\text{O})_5(\text{OH})]^{2+}(\text{aq}) + \text{H}^{+}(\text{aq}) \)
• pH: Low (pH 2-3).

Silicon Tetrachloride ($\text{SiCl}_4$):

• Observation: Rapid, vigorous reaction producing white fumes ($\text{HCl}$) and a white precipitate of silicic acid.
• Equation: \( \text{SiCl}_4(\text{l}) + 4\text{H}_2\text{O}(\text{l}) \rightarrow \text{Si}(\text{OH})_4(\text{s}) + 4\text{HCl}(\text{aq}) \)
• Resulting Acid Structure: Silicic acid ($\text{H}_4\text{SiO}_4$ or $\text{Si}(\text{OH})_4$).
• pH: Very low (highly acidic).

Phosphorus(V) Chloride ($\text{PCl}_5$):

• Observation: Very violent reaction, producing dense white fumes ($\text{HCl}$).
• Equation: \( \text{PCl}_5(\text{s}) + 4\text{H}_2\text{O}(\text{l}) \rightarrow \text{H}_3\text{PO}_4(\text{aq}) + 5\text{HCl}(\text{aq}) \)
• Resulting Acid Structure: Phosphoric acid ($\text{H}_3\text{PO}_4$).
• pH: Very low (highly acidic).

Did you know? Silicon and Phosphorus chlorides react so violently because $\text{Si}$ and $\text{P}$ atoms have available d-orbitals which allow them to temporarily expand their octet, accepting a lone pair from the water molecule easily.

3.3. Reactions with Acids and Bases (Amphoteric Properties)

Only the chlorides that form highly charged cations (like $\text{Al}^{3+}$) can react with bases in a secondary step.

Aluminium Chloride/Hydroxide Reactions:

When $\text{AlCl}_3$ hydrolyses, it forms the hydroxide $\text{Al}(\text{OH})_3$ (which is amphoteric), allowing it to react with excess alkali.

• Reaction with Base (Formation of precipitate):
  \( \text{Al}^{3+}(\text{aq}) + 3\text{OH}^{-}(\text{aq}) \rightarrow \text{Al}(\text{OH})_3(\text{s}) \)
• Reaction with Excess Base (Dissolving the precipitate):
  \( \text{Al}(\text{OH})_3(\text{s}) + \text{OH}^{-}(\text{aq}) \rightarrow [\text{Al}(\text{OH})_4]^{-}(\text{aq}) \)