Physical Properties of Ionic Compounds

• Composition: Composed of ions, not atoms or molecules.
• Bonding: Characterized by strong electrostatic forces.
• Solubility: Soluble in water; insoluble in organic solvents (e.g., benzene, propanone).
• Melting and Boiling Points: High melting and boiling points.
• State at Room Temperature: Hard, brittle solids.
• Electrical Conductivity: Conduct electricity when in a molten state or aqueous solution (dissolved in water).

Physical Properties of Covalent Compounds

• Composition: Made up of molecules, not ions.
• Bonding: Low intermolecular forces.
• Solubility: Typically insoluble in water; soluble in organic solvents (e.g., benzene, propanone).
• Melting and Boiling Points: Low melting and boiling points.
• State at Room Temperature: Mostly gases or volatile liquids.

Structural Differences Between Ionic and Covalent Compounds

1. Structure of Ionic Compounds:
   • Ions are arranged in a regular pattern, forming a crystal lattice.
   • Cations and anions alternate, creating a giant structure held together by strong ionic bonds.
   • Example: Sodium chloride (NaCl) forms a crystal lattice.
2. Effects on Physical Properties:
   • High melting and boiling points due to strong electrostatic attractions.
   • Do not conduct electricity in solid form; conduct when molten or in solution (ions can move freely). Aqueous solutions that conduct electricity are termed electrolytes.
3. Structure of Simple Covalent Compounds:
   • Atoms within molecules are held by strong covalent bonds, but molecules are attracted to each other by weak intermolecular forces (Van der Waals forces or hydrogen bonds).
4. Effects on Physical Properties:
   • Low melting and boiling points due to weak intermolecular forces; they melt or vaporize easily.
   • Do not conduct electricity (non-electrolytes) as they lack ions.

Distinguishing Ionic and Covalent Solutions Using Conductivity

Materials Needed:

• 2 cells, a bulb (3.8V, 0.3A)
• 3 connecting wires with crocodile clips
• Iron nails or graphite rods
• 40 ml of each: sodium chloride solution, sugar solution, dilute sulfuric acid, ethanol, distilled water
• 100 ml beaker

Procedure:

1. Set up the apparatus as shown in the diagram.
2. Pour 40 ml of sodium chloride solution into the beaker.
3. Dip the nails into the solution.
4. Observe the bulb; record “light” or “no light” in the results table.
5. Remove the nails and rinse the beaker and nails with distilled water.
6. Repeat steps 2 to 5 for sugar solution, dilute sulfuric acid, and ethanol.

Classification:

• Ionic Solutions: Produce light in the bulb.
• Covalent Solutions: Do not produce light in the bulb.

Common Examples of Electrolytes and Non-Electrolytes

• Electrolytes:
  • Sodium chloride solution
  • Copper (II) sulfate solution
  • Sodium hydroxide solution
  • Hydrochloric acid
  • Sulfuric acid
  • Ethanoic acid
  • Molten sulfur
• Non-Electrolytes:
  • Pure water
  • Sugar solution
  • Paraffin wax
  • Ethanol
  • Urea

Covalent Bonds and Intermolecular Forces

1. Pure Covalent Bonds

• Definition: A pure covalent bond is formed when the shared pair of electrons is equally contributed by the bonded atoms.
• Example: In ammonia (NH₃), both hydrogen and nitrogen atoms contribute one electron each to the shared pair.

2. Dative Covalent Bonds

• Definition: A dative covalent bond occurs when only one atom contributes both electrons to the shared pair.
• Example: The formation of the ammonium ion (NH₄⁺) demonstrates dative bonding, where the nitrogen atom in ammonia donates a lone pair of electrons to a hydrogen ion (H⁺).

Polar and Non-Polar Covalent Bonds

1. Polar Covalent Bonds

• Definition: A polar covalent bond forms when electrons are not shared equally due to differences in electronegativity between the atoms.
• Characteristics:
  • The more electronegative atom acquires a partial negative charge (δ⁻), while the less electronegative atom acquires a partial positive charge (δ⁺).
  • Molecules with polar covalent bonds are called polar molecules and can conduct electricity in aqueous solutions.
• Example: Hydrogen chloride (HCl) is a common example of a polar covalent compound.

2. Non-Polar Covalent Bonds

• Definition: A non-polar covalent bond occurs when electrons are shared equally between atoms, usually when they have the same electronegativity.
• Characteristics:
  • Non-polar molecules do not possess partial charges and therefore do not conduct electricity.
• Examples: Hydrogen (H₂) and chlorine (Cl₂) are examples of molecules with non-polar covalent bonds.

Types of Intermolecular Forces

There are two main types of intermolecular forces: hydrogen bonds and Van der Waals forces.

1. Hydrogen Bonding

• Definition: This type of bonding occurs between molecules containing hydrogen and a more electronegative atom (e.g., oxygen, nitrogen).
• Characteristics: The attraction between the partially negative atom and the partially positive hydrogen atom is called a hydrogen bond.

2. Van der Waals Forces

• Definition: These are weak attractive or repulsive forces that exist between molecules.
• Characteristics: Van der Waals forces are present in most molecular substances and contribute to their physical properties.

Effects of Intermolecular Forces on Physical Properties

• Boiling Point of Water: Water has a higher boiling point compared to many covalent compounds due to hydrogen bonding between its molecules.
• Surface Tension: The strong hydrogen bonds in water contribute to its high surface tension.
• Liquid State of Water: Hydrogen bonding is essential for maintaining water in its liquid state.
• Boiling Points of Halogens: Van der Waals forces contribute to the higher boiling points observed in certain elements, such as halogens.

Allotropes

Definition: Allotropes are different forms of the same element that exist in the same physical state.

• Examples:
  • Oxygen: Exists as oxygen (O₂) and ozone (O₃).
  • Sulfur: Exists as rhombic sulfur and monoclinic sulfur.
  • Carbon: Exists as graphite and diamond.

Structure and Properties of Graphite

Structure of Graphite

• Each carbon atom forms covalent bonds with three other carbon atoms, creating rings of six atoms.
• These rings form flat sheets that can slide over one another easily.
• Sheets are held together by weak forces.

Physical Properties of Graphite

• Soft and slippery.
• Dark grey in color.
• Conducts electricity due to one free electron from each carbon atom that can move and carry charge.

Uses of Graphite

• Lubricant for engines and locks.
• Mixed with clay to produce pencil lead.
• Used as electrodes in electric circuits.
• Utilized for connecting brushes in generators.

Structure and Properties of Diamond

Structure of Diamond

• Each carbon atom forms four covalent bonds with four other carbon atoms, creating a tetrahedral structure.
• This arrangement results in a very strong and rigid structure, with a melting point of 3550ºC.

Physical Properties of Diamond

• The hardest known substance on Earth.
• Colorless and sparkles when cut.
• Does not conduct electricity.

Uses of Diamond

• Made into jewelry (necklaces, rings, earrings) due to its aesthetic appeal.
• Utilized in drilling equipment and cutting tools because of its hardness.

Silicon Dioxide (Silica)

Description: Silicon dioxide (SiO₂) naturally occurs as quartz, the main mineral found in sand.

Structure of Silicon Dioxide

• Each silicon atom forms four covalent bonds with oxygen atoms, while each oxygen atom bonds with two silicon atoms.
• Results in a very hard substance with a melting point of 1710ºC.

Comparison: Diamond vs. Silicon Dioxide

• Hardness:

  • Diamond is the hardest substance on Earth.

  • Silicon dioxide is also very hard but not as hard as diamond.

• Conductivity:

  • Diamond does not conduct electricity.

  • Silicon dioxide also does not conduct electricity.

• Melting Point:

  • Diamond has a very high melting point of 3550ºC.

  • Silicon dioxide has a high melting point of 1710ºC.

Properties of Metals

• High Melting and Boiling Points: Due to strong metallic bonds between positive metal ions and a mobile sea of electrons.
• High Density: Atoms are closely packed together.
• Ductility: Metals can be drawn into thin wires.
• Malleability: Metals can be hammered into different shapes.
• Sonority: Metals produce a ringing sound.
• Shininess: Metals have a lustrous appearance.
• Good Conductors of Heat: Free electrons transfer heat throughout the metal structure.
• Good Conductors of Electricity: Mobile electrons carry charge when voltage is applied.

Uses of Metals in Relation to Their Properties

• Good Electrical Conductors:
  • Used in electrical wiring and appliances (e.g., TVs, radios, computers).
• Good Thermal Conductors:
  • Used to make cooking utensils (e.g., pots).
• Malleable:
  • Used to create shaped objects (e.g., car bodies, kitchenware).
• Ductile:
  • Used to make connecting wires (e.g., copper wire).
• Sonorous:
  • Used to make bells.
• Shiny:
  • Used in jewelry (e.g., earrings, bangles, necklaces).

Alloys

An alloy is a mixture of two or more elements, at least one of which is a metal.

Examples of Alloys

• Stainless Steel: Combination of iron and chromium.
• Brass: Mixture of copper and zinc.
• Bronze: Combination of copper and tin.

Properties and Uses of Alloys

• Brass (70% copper, 30% zinc)
  • Properties: Harder than pure copper, gold-colored.
  • Uses: Kitchenware, musical instruments, screws, radiators.
• Stainless Steel (74% iron, 18% chromium, 8% carbon)
  • Properties: Shiny, strong, rust-resistant.
  • Uses: Cutlery, surgical instruments, kitchen sinks.
• Steel (99% iron, 1% carbon)
  • Properties: Hard and strong.
  • Uses: Bridges, bodies of cars, railway tracks.
• Bronze (90% copper, 10% tin)
  • Properties: Hard, strong, shiny surface, rust-resistant.
  • Uses: Statues, monuments, medals, swords, artistic materials.