5.5 Nature of Bonding, Structure and Properties

This topic explains how the type of bonding and the structure of a substance determine its physical properties — such as melting point, boiling point, hardness, and electrical conductivity.

Main Idea

Three main factors are important when determining the properties of a substance:

• Type of particles — atoms, ions, or molecules.
• The way elementary particles are connected to each other — ionic, covalent, metallic, or weak intermolecular forces.
• The arrangement of particles — planes, layers, or three-dimensional networks.

Simple explanation: What a substance is made of, how its particles are joined, and how they are arranged — these three things decide its properties.

Core Concept

Detailed Explanation

Factor 1: Type of Particles

The types of elementary particles contained in a substance can be:

• Atoms — e.g., in diamond, graphite (giant covalent structures)
• Ions — e.g., in sodium chloride (ionic compound)
• Molecules — e.g., in water, carbon dioxide (covalent compounds)

If a substance contains ions (such as sodium chloride), it will conduct electricity when molten or dissolved.

Simple explanation: The type of particle decides whether the substance can conduct electricity.

Factor 2: The Way Elementary Particles Are Connected

• Ionic bonds
• Covalent bonds
• Metallic bonds
• Weak intermolecular forces

The stronger the bond, the higher the melting/boiling point and hardness of the substance.

Examples from the book:

• Has strong covalent bonds connecting each atom to several other atoms.
• Forms a giant covalent structure.
• The atoms are difficult to separate.
• So it is very hard and difficult to melt.

• Has strong covalent bonds between C and O atoms.
• But weak intermolecular forces between CO₂ molecules.
• The molecules are easily separated.
• So CO₂ has a low melting/boiling point.

Simple explanation: Even if the bonds inside molecules are strong, weak forces between molecules can make the substance easy to melt.

Factor 3: The Arrangement of Particles

• In planes — e.g., polymers
• In layers — e.g., clays, graphite
• In three-dimensional networks — e.g., diamond, metals

• Atoms arranged in 2-dimensional layers.
• This allows the layers to move over one another.
• This is why graphite is used in pencil writing and as a lubricant.

• Has a large three-dimensional network of carbon atoms.
• This makes it the hardest substance on Earth.

• Also have giant structures.
• Metallic bonding is strong.
• Most metals have very high melting and boiling points and are thermally stable.

Simple explanation: Arrangement matters — layers slide (graphite), 3D networks are hard (diamond), giant structures are strong (metals).

Conduction of Electricity in Ionic Compounds

Electrical conductivity is achieved by the movement of charged particles.

• Ionic compounds cannot conduct electricity in the solid state because their ions remain in fixed positions and cannot move.
• When an ionic compound is melted or dissolved in water, it is ionized.
• Its ions move freely in molten or aqueous solution.
• Therefore, electricity can pass through a molten ionic compound or its aqueous solution.

Simple explanation: Solid ionic compounds don't conduct because ions are locked in place. When molten or dissolved, ions are free to move and carry current.

Conduction of Electricity Through Acids

• Covalent compounds have no free charged particles, so they do not conduct electricity.
• However, some covalent compounds conduct electricity when dissolved in water.
• For instance, acids like HCl, H₂SO₄, HNO₃, etc.
• When these acids are dissolved in water, they ionize and form high concentrations of H⁺ ions and negatively charged ions.
• These ions can move freely in aqueous solution.
• Therefore, aqueous solutions of acids conduct electricity.

Simple explanation: Pure covalent compounds don't conduct, but acids in water form ions and do conduct.

Conduction of Electricity Through Metals

• Metals are good conductors of electricity because they have free electrons.
• These electrons are not associated with a single atom.
• These electrons begin to flow under the influence of electricity.
• Therefore, metals allow electricity to pass through.

Simple explanation: Metals have free electrons that carry electric current.

Covalent Compounds and Their Properties

Compounds that consist of covalent molecules are called covalent compounds.

• The intermolecular forces between their molecules are much weaker than the covalent bonds.
• Therefore, covalent compounds have low melting and boiling points.
• Since their molecules do not contain any free electrons or ions, they are poor conductors of electricity.

Simple explanation: Covalent compounds have weak forces between molecules → low melting/boiling points → poor conductors.

Intermolecular Forces and Their Influence on Melting and Boiling Points

Table 5.1: Melting and boiling points of some covalent compounds

Table 5.2: Melting and boiling points of some ionic compounds

Key observation: Covalent compounds usually have much lower melting points than ionic compounds.

• Ionic compounds: melting involves breaking the ionic bond — breaking the electrostatic forces between ions requires large amounts of energy. So ionic compounds have high melting and boiling points.
• Covalent compounds: melting involves breaking the intermolecular forces, which are much weaker than electrostatic forces. So less energy is required to break the intermolecular forces between covalent molecules.

Simple explanation: Ionic compounds have strong ionic bonds → high melting points. Covalent compounds have weak intermolecular forces → low melting points.

Intermediate Level

How Bonding Determines Properties

Comparison of Ionic and Covalent Compounds

Why Ionic Compounds Conduct When Molten/Dissolved

• In the solid state, ions are held in fixed positions — they cannot move.
• When melted or dissolved, the ions are free to move.
• Moving ions carry electric current → conducts electricity.

Why Metals Conduct

• Metals have free (delocalized) electrons.
• These electrons are not bound to any single atom.
• They can move freely through the metal lattice.
• When a voltage is applied, electrons flow → conducts electricity.

Why Acids Conduct in Water

• Pure acids are covalent compounds — no free ions.
• In water, acids ionize → produce H⁺ ions and negative ions.
• These ions move freely → conducts electricity.

Important Points

• - Three factors determine properties: type of particles, type of bonding, arrangement of particles.
• - Stronger bonds → higher melting/boiling points and hardness.
• - Ionic compounds: high melting/boiling points; conduct when molten/dissolved.
• - Covalent compounds: low melting/boiling points; poor conductors.
• - Metals: good conductors due to free electrons.
• - Acids in water: conduct due to ionization.
• - Graphite: layers slide → lubricant; conducts electricity.
• - Diamond: 3D network → hardest substance.
• - Ionic compounds: melting breaks ionic bonds → high energy required.
• - Covalent compounds: melting breaks weak intermolecular forces → low energy required.

Remember

• - Three factors: particles, bonding, arrangement.
• - Strong bonds → high melting point, hard.
• - Weak forces → low melting point, soft.
• - Ionic compounds: high MP/BP; conduct when molten/dissolved.
• - Covalent compounds: low MP/BP; poor conductors.
• - Metals: conduct due to free electrons.
• - Acids in water: conduct due to ionization.
• - Graphite: layers slide; conducts.
• - Diamond: 3D network; hardest.

Common Mistakes

• 1. Thinking all covalent compounds have low melting points. Diamond (a covalent compound) has a very high melting point.
• 2. Thinking ionic compounds conduct in solid state. They conduct only when molten or dissolved.
• 3. Thinking all covalent compounds are poor conductors. Acids in water conduct electricity.
• 4. Confusing intermolecular and intramolecular forces. Melting covalent compounds breaks intermolecular forces, not covalent bonds.
• 5. Thinking graphite and diamond have the same properties. They have very different structures and properties.
• 6. Forgetting that metals conduct due to free electrons. This is a key difference from ionic and covalent compounds.

Quick Revision

• Three main factors determine properties:
• Stronger bonds → higher melting/boiling point and hardness.
• Silicon dioxide (SiO₂): giant covalent structure → very hard, difficult to melt.
• Carbon dioxide (CO₂): weak intermolecular forces → low melting/boiling point.
• Graphite: 2D layers → layers slide → lubricant; conducts electricity.
• Diamond: 3D network → hardest substance.
• Metals: giant structures, strong metallic bonding → high melting/boiling points.
• Ionic compounds: conduct only when molten or dissolved.
• Covalent compounds: generally poor conductors; acids conduct in water.
• Metals: good conductors due to free electrons.
• Covalent compounds: low melting/boiling points; poor conductors.
• Ionic compounds: high melting/boiling points.

60-Second Revision

• Three factors: particles, bonding, arrangement.
• Strong bonds → high MP/BP; weak forces → low MP/BP.
• Ionic compounds: high MP/BP; conduct when molten/dissolved.
• Covalent compounds: low MP/BP; poor conductors (except acids in water).
• Metals: conduct due to free electrons.
• Graphite: layers slide, conducts.
• Diamond: hardest substance.
• Warning: melting covalent compounds breaks intermolecular forces, not covalent bonds.