Bonding and Structure

Welcome to Bonding and Structure!

Ever wondered why some substances, like salt, dissolve in water while others, like diamond, are hard enough to cut through rock? It all comes down to how atoms are "glued" together. In this chapter, we are going to explore the different types of chemical bonds and how the shape of a molecule determines its personality. Don't worry if it seems a bit abstract at first—we'll use plenty of analogies to make it stick!

Topic 2A: Bonding

1. Ionic Bonding: The "Give and Take"

Ionic bonding is the strong electrostatic attraction between oppositely charged ions. Think of it like two powerful magnets snapping together. This happens when a metal atom gives away electrons to a non-metal atom.

Factors affecting bond strength:
• Ionic Charge: The higher the charge, the stronger the attraction. An ion with a 2+ charge will pull much harder than an ion with a 1+ charge.
• Ionic Radius: Smaller ions can get closer together, making the attraction stronger. It’s like how magnets are harder to pull apart the closer they are.

Trends in Ionic Radii:
• Down a group: The radius increases because there are more electron shells.
• Isoelectronic ions (ions with the same number of electrons, like \(N^{3-}\) to \(Al^{3+}\)): As the number of protons in the nucleus increases, the "pull" on the electrons gets stronger, so the ionic radius decreases from \(N^{3-}\) to \(Al^{3+}\).

Evidence for ions: We know ions exist because they can migrate. If you put a drop of purple potassium manganate(VII) on wet filter paper and apply a voltage, the purple color moves toward the positive electrode. This proves charged particles are physically moving!

Quick Review: Ionic bonding = Metal + Non-metal. Stronger when ions are small and highly charged.

2. Covalent Bonding: The "Sharing"

A covalent bond is the strong electrostatic attraction between two nuclei and a shared pair of electrons between them. It’s like a tug-of-war where neither side wants to let go of the rope.

Dot-and-Cross Diagrams:
These show the outer shell electrons. Remember:
• Single bond = 1 shared pair.
• Double bond = 2 shared pairs.
• Triple bond = 3 shared pairs.

Dative Covalent (Co-ordinate) Bonding:
This is a special covalent bond where one atom provides both electrons for the shared pair. It’s like one friend providing the entire packed lunch for two people to share.
Examples: The Ammonium ion (\(NH_{4}^{+}\)) and Aluminium Chloride (\(Al_{2}Cl_{6}\)). In diagrams, we often show this with an arrow \(\rightarrow\) pointing away from the donor atom.

Bond Length vs. Bond Strength:
The shorter the bond, the stronger it is. Think of it like a short, tight spring versus a long, floppy one. Triple bonds are shorter and much stronger than single bonds.

3. The Shapes of Molecules (VSEPR Theory)

Molecules aren't just flat drawings; they have 3D shapes! The shape is decided by Electron Pair Repulsion Theory. Electrons are all negatively charged, so they want to stay as far away from each other as possible.

The Golden Rule: Lone pairs (unbonded pairs) push harder than bonding pairs. Lone pair-lone pair repulsion > Lone pair-bonding pair > Bonding pair-bonding pair. Each lone pair usually reduces the bond angle by about 2.5°.

Common Shapes to Memorize:
• 2 bonding pairs: Linear (180°) – Example: \(BeCl_{2}, CO_{2}\)
• 3 bonding pairs: Trigonal Planar (120°) – Example: \(BCl_{3}\)
• 4 bonding pairs: Tetrahedral (109.5°) – Example: \(CH_{4}, NH_{4}^{+}\)
• 3 bonding, 1 lone pair: Trigonal Pyramidal (107°) – Example: \(NH_{3}\)
• 2 bonding, 2 lone pairs: Bent/Non-linear (104.5°) – Example: \(H_{2}O\)
• 5 bonding pairs: Trigonal Bipyramidal (90° and 120°) – Example: \(PCl_{5}\)
• 6 bonding pairs: Octahedral (90°) – Example: \(SF_{6}\)

Mnemonic: Little Tigers Talk To Big Otter (Linear, Trigonal Planar, Tetrahedral, Trigonal Pyramidal, Bent, Octahedral).

4. Electronegativity and Polarity

Electronegativity is the ability of an atom to attract the bonding electrons in a covalent bond. Think of it as how "greedy" an atom is for electrons.

Bond Polarity: If two different atoms are bonded (like H and Cl), the more electronegative one (Cl) pulls the electrons closer, becoming slightly negative (\(\delta-\)) and leaving the other slightly positive (\(\delta+\)). This creates a polar bond.

Molecular Polarity: A molecule can have polar bonds but be non-polar overall if the shape is symmetrical (like \(CO_{2}\) or \(CCl_{4}\)). The dipoles "cancel out." If the molecule is asymmetrical (like \(H_{2}O\)), it is polar.

5. Intermolecular Forces (The Weak Connections)

These are forces between molecules, not inside them. They are much weaker than covalent or ionic bonds.

i. London Forces (Instantaneous Dipole – Induced Dipole):
These exist between all molecules. Electrons are always moving; for a split second, more might be on one side, creating a tiny temporary magnet that induces a dipole in the neighbor. Larger molecules have more electrons, so they have stronger London forces.

ii. Permanent Dipole-Dipole Interactions:
These happen between polar molecules (like HCl). The \(\delta+\) end of one molecule is attracted to the \(\delta-\) end of another.

iii. Hydrogen Bonding:
The strongest type! It only happens when Hydrogen is bonded to Fluorine, Oxygen, or Nitrogen (The "FON" elements). These are very electronegative, leaving the H very electron-poor.

Anomalous Properties of Water:
Hydrogen bonding explains why water is weird:
1. High melting/boiling point: It takes a lot of energy to break those strong hydrogen bonds.
2. Ice is less dense than water: When ice freezes, hydrogen bonds hold the molecules in an open, hexagonal lattice. This makes it float!

Boiling Point Trends:
• Alkanes: Boiling point increases with chain length (more electrons = stronger London forces). Branching lowers the boiling point because molecules can't pack as tightly.
• Hydrogen Halides: \(HF\) has a very high boiling point because of Hydrogen bonding. From \(HCl\) to \(HI\), the boiling point increases because there are more electrons (stronger London forces).

Solubility Rule of Thumb: "Like dissolves like." Polar solvents (like water) dissolve ionic and polar substances. Non-polar solvents dissolve non-polar substances.

Topic 2B: Structure

1. Giant Lattices vs. Simple Molecular

The physical properties of a substance depend on its structure.

Giant Ionic Lattices: (e.g., \(NaCl\))
• High melting points (strong bonds).
• Conduct electricity only when molten or in solution (ions are free to move).

Giant Covalent Lattices: (e.g., Diamond, Graphite, \(SiO_{2}\))
• Extremely high melting points.
• Diamond: Each Carbon bonded to 4 others. Very hard. No free electrons (insulator).
• Graphite: Each Carbon bonded to 3 others in layers. Layers slide (lubricant). Delocalised electrons between layers allow it to conduct electricity.
• Graphene: A single 2D layer of graphite. Incredibly strong and conductive.

Giant Metallic Lattices: (e.g., Magnesium, Copper)
• Metallic bonding is the attraction between metal ions and a "sea" of delocalised electrons.
• Conduct electricity as solids (electrons can flow).
• Malleable (layers of ions can slide over each other).

Simple Molecular Structures: (e.g., \(I_{2}\), Ice, \(CO_{2}\))
• Low melting/boiling points because you are only breaking weak intermolecular forces, not the strong covalent bonds inside the molecule.
• Non-conductors (no free charged particles).

Key Takeaway: If a substance has a very high melting point, it’s a Giant structure. If it conducts as a solid, it’s Metallic (or Graphite). If it only conducts when liquid/aqueous, it’s Ionic.

Don't forget!

When you are asked why a simple molecular substance like Bromine has a low boiling point, never say the covalent bonds break. The covalent bonds are fine! You are only overcoming the weak London forces between the molecules.