• Periodicity refers to the recurring trends or patterns in the properties of elements in the periodic table.
• These trends allow us to predict the chemical and physical properties of elements based on their position in the periodic table.
Atomic radius
Decreases as we move from left to right across a period.
Increases as we move down a group.
This is due to the increasing effective nuclear charge and the addition of electron shells, which leads to a decrease in the size of the atoms.
Ionic radius
The ionic radius of an element is the radius of the ions it forms.
Decreases as we move from left to right across a period.
Increases as we move down a group.
This is due to the increasing effective nuclear charge and the addition of electron shells, which leads to a decrease in the size of the atoms.
From Na⁺ to Al³⁺
Number of energy shells are identical.
Number of electrons are identical.
From Na⁺ to Al³⁺ the number of protons increases.
Effective nuclear charge increases.
Ionic radius decreases.
From P³⁻ to Cl⁻
The anions are larger than their original atoms.
P³⁻ has three more electrons than protons.
P³⁻ ion has weakest nuclear attraction.
So, P³⁻ has largest ionic radius.
From P³⁻ to Cl⁻ ionic radius decreases.
Reactions with Oxygen
Sodium (Na) reacts with oxygen with yellow/orange flame to form sodium oxide (Na₂O).
2Na + O₂ → 2Na₂O.
Magnesium (Mg) reacts with oxygen with bright white flame to form magnesium oxide (MgO).
2Mg + O₂ → 2MgO.
Aluminum (Al) reacts with oxygen with bright white flame to form aluminum oxide (Al₂O₃).
4Al + 3O₂ → 2Al₂O₃.
Phosphorus (P) reacts with oxygen to form phosphorus trioxide (P₂O₃) and phosphorus pentoxide (P₂O₅). The reactions can be represented by the following equations:
P + O₂ → P₂O₃ (It is a white crystal)
P₂O₃ + O₂ → P₂O₅. (It is a white solid)
Sulfur (S) reacts with oxygen to form sulfur dioxide (SO₂).
2S + O₂ → 2SO₂.
Reactions with Chlorine
Sodium (Na) reacts with chlorine to form sodium chloride (NaCl).
2Na + Cl₂ → 2NaCl.
Magnesium (Mg) reacts with chlorine to form magnesium chloride (MgCl₂).
Mg + Cl₂ → MgCl₂.
Aluminum (Al) reacts with chlorine to form aluminum chloride (AlCl₃).
2Al + 3Cl₂ → 2AlCl₃
Silicon (Si) reacts with chlorine to form silicon tetrachloride (SiCl₄).
Si + 2Cl₂ → SiCl₄ (a colorless fuming liquid)
Phosphorus (P) reacts with chlorine to form phosphorus pentachloride (PCl₅).
P + 5Cl₂ → PCl₅
Reactions with Water:
Sodium (Na) reacts with water to form sodium hydroxide (NaOH) and hydrogen gas (H₂).
Na + H₂O → NaOH + H₂
Magnesium (Mg) reacts with water to form magnesium hydroxide (Mg (OH)₂) and hydrogen gas (H₂).
Mg + 2H₂O → Mg (OH)₂ + H₂
Reactions of Chloride with water:
Sodium chloride (NaCl) dissolves in water to form sodium ions (Na⁺) and chloride ions (Cl⁻).
NaCl (s) → Na⁺ (aq) + Cl⁻ (aq) [ pH at 7]
Magnesium chloride (MgCl₂) dissolves in water to form magnesium ions (Mg²⁺) and chloride ions (Cl⁻).
MgCl₂ (s) → Mg²⁺ (aq) + 2Cl⁻ (aq) [pH at 7]
Aluminum chloride (AlCl₃) reacts with water to form aluminum hydroxide (Al (OH)₃) and hydrochloric acid (HCl).
AlCl₃ + 3H₂O → Al (OH)₃ + 3HCl [pH less than7]
Silicon tetrachloride (SiCl₄) reacts with water to form silicon dioxide (SiO₂), hydrochloric acid (HCl), and hydrogen chloride gas (HCl).
SiCl₄ + 2H₂O → SiO₂ + 4HCl [pH less than 7]
Phosphorus pentachloride (PCl₅) reacts with water to form phosphoric acid (H₃PO₄) and hydrochloric acid (HCl).
PCl₅ + 6H₂O → H₃PO₄ + 5HCl [pH less than 7]
Phosphorus trichloride (PCl₃) reacts with water to form phosphoric acid (H₃PO₄) and hydrochloric acid (HCl).
PCl₃ + 3H₂O → H₃PO₄ + 3HCl [pH less than 7]
Sulfur dichloride (SCl₂) reacts with water.
SCl₂ + 2H₂O → S + SO₂ + 2HCl [pH less than 7]
Reactions of oxides with water:
Sodium oxide (Na₂O) reacts with water to form sodium hydroxide (NaOH).
Na₂O + H₂O → 2NaOH
Magnesium oxide (MgO) reacts with water to form magnesium hydroxide (Mg (OH)₂). MgO + H₂O → Mg (OH)₂
Aluminum oxide (Al₂O₃) does not react with water under normal conditions.
Phosphorus pentoxide (P₂O₅) reacts with water to form phosphoric acid (H₃PO₄).
P₂O₅ + 3H₂O → 2H₃PO₄
Sulfur dioxide (SO₂) dissolves in water to form sulfurous acid (H₂SO₃).
SO₂ + H₂O → H₂SO₃
Phosphorus trioxide (P₂O₃) reacts with water to form phosphorous acid (H₃PO₃).
P₂O₃ + 3H₂O → 2H₃PO₃
Silicon dioxide (SiO₂) does not react with water under normal conditions.
Bonds and melting points.
Chlorides
• In the chlorides (NaCl, MgCl₂, AlCl₃, SiCl₄, and PCl₅), the bonding is primarily ionic.
• This is because the elements in the third period are metals (Na, Mg, Al) or non-metals (Si, P) that form positive and negative ions when they bond.
• The presence of ionic bonds is evidenced by the high melting and boiling points of these compounds, as well as their solubility in water.
Oxides:
Acidic or basic nature:
The oxides of metals (Na, Mg, Al, Si) are generally basic in nature and react with acids to form salts and water.
The oxides of nonmetals (P, S, Cl) are generally acidic in nature and react with bases to form salts and water.
The oxide of the element in the middle of the period, Aluminum (Al), Zinc (Zn) and lead (Pb), are amphoteric and can act as both an acid and a base.
Oxides such as Carbon Monoxide (CO) or Nitrogen Monoxide (NO) are neutral oxides and do not react.
Melting points:
The melting points of the oxides of metals (Na₂O, MgO, Al₂O₃) are generally high due to the strong ionic bonding between the metal and oxygen atoms.
The melting points of the oxides of nonmetals (P₄O₁₀, SO₂, Cl₂O₇) are generally lower than those of metal oxides due to the weaker covalent bonding between the nonmetal and oxygen atoms.
Electrical conductivity in liquid state:
The oxides of metals (Na₂O, MgO, Al₂O₃) are generally not conductive in the solid state since the ions are held in fixed positions and cannot move freely.
The oxides of metals (Na₂O, MgO, Al₂O₃) are generally conductive in the molten state since the ions can move freely.
The oxides of nonmetals (P₄O₁₀, SO₂, Cl₂O₇) can conduct electricity in the dissolved state due to the presence of mobile ions.
Chemical bonding and structure:
The oxides of metals (Na₂O, MgO, Al₂O₃) are ionic in nature and have a lattice structure. The strong electrostatic attraction between the ions gives rise to the ionic bonding in these oxides.
The oxides of nonmetals (P₄O₁₀, SO₂, Cl₂O₇) are covalent in nature and have a simple molecular structure. The nonmetal atoms share electrons to form covalent bonds with oxygen atoms.
Some important reactions:
Reaction of Aluminum Chloride with Sodium hydroxide
Example: AlCl₃ + 3 NaOH → 3 NaCl + Al (OH)₃
Reaction of HCl and HClO with NaOH Example:
2NaOH + HClO + HCl → NaCl + NaClO + 2H₂O
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