Welcome to Bonding in Chemistry
Welcome to one of the most fundamental chapters in AS Chemistry! Have you ever wondered why table salt forms beautiful crystal cubes, why water stays liquid at room temperature while carbon dioxide is a gas, or why diamonds are extraordinarily hard while graphite slides smoothly onto paper? The answer to all of these questions lies in chemical bonding.
Don't worry if bonding felt overwhelming at GCSE. In this chapter, we will build things up step-by-step with clear definitions, visual descriptions, analogies, and memory tricks to help you master every concept required for your CCEA AS 1 exam.
1. Ionic Bonding and Giant Ionic Lattices
What is an Ionic Bond?
An ionic bond is defined as the electrostatic attraction between oppositely charged ions acting in all directions throughout a lattice.
Ionic bonds form when electrons are transferred from metal atoms to non-metal atoms:
• Metals lose electrons to form positively charged ions (called cations).
• Non-metals gain electrons to form negatively charged ions (called anions).
Analogy: Think of ionic bonding like two magnets snapping together. The positive magnet (cation) and negative magnet (anion) strongly attract one another.
The Giant Ionic Lattice
Ionic compounds do not exist as isolated single pairs of ions. Instead, they form a giant ionic lattice. This is a regular, repeating three-dimensional arrangement where each positive ion is surrounded by negative ions, and each negative ion is surrounded by positive ions.
For example, in sodium chloride, \( \text{NaCl} \), each sodium ion, \( \text{Na}^+ \), is surrounded octahedrally by six chloride ions, \( \text{Cl}^- \), and each chloride ion is surrounded by six sodium ions. This gives a coordination number of \( 6:6 \).
Physical Properties of Ionic Compounds
1. High Melting and Boiling Points:
Giant ionic lattices contain a massive number of strong electrostatic forces of attraction between oppositely charged ions. A large amount of thermal energy is required to overcome these strong ionic bonds throughout the entire lattice.
2. Electrical Conductivity:
• Solid state: Ionic compounds do not conduct electricity because the ions are locked into fixed positions within the lattice and cannot move.
• Molten (liquid) or Aqueous (dissolved in water) state: Ionic compounds conduct electricity because the lattice structure breaks down, allowing the ions to become mobile charge carriers.
3. Brittleness:
When an ionic crystal is hit with a mechanical force, layers of ions slide over each other. This brings ions of the same charge next to each other (e.g., \( \text{Na}^+ \) next to \( \text{Na}^+ \)). The resulting strong electrostatic repulsion shatters the crystal apart.
4. Solubility in Water:
Many ionic substances dissolve in polar solvents such as water. The partially negative oxygen atoms of water molecules are attracted to the positive cations, while the partially positive hydrogen atoms are attracted to the negative anions. This process releases hydration energy which helps overcome the lattice enthalpy.
Common Mistake to Avoid: Never say that ionic compounds conduct electricity because "electrons are free to move." In ionic compounds, the mobile charge carriers are ions, not electrons!
Key Takeaway for Ionic Bonding: An ionic bond is the electrostatic attraction between oppositely charged ions in a giant lattice. Ionic compounds have high melting points, conduct only when molten or aqueous, and are brittle.
2. Covalent and Coordinate (Dative) Bonding
What is a Covalent Bond?
A covalent bond is defined as the electrostatic attraction between a shared pair of electrons and the positively charged nuclei of the bonded atoms.
Covalent bonds form between non-metal atoms when they share valence electrons to achieve stable noble-gas electron configurations:
• Single Covalent Bond: One shared pair of electrons (e.g., \( \text{Cl}_2 \), \( \text{H}_2\text{O} \), \( \text{CH}_4 \)).
• Double Covalent Bond: Two shared pairs of electrons (e.g., \( \text{O}_2 \), \( \text{CO}_2 \)).
• Triple Covalent Bond: Three shared pairs of electrons (e.g., \( \text{N}_2 \)).
Coordinate (Dative Covalent) Bonding
A coordinate bond (also known as a dative covalent bond) is a special type of covalent bond where both electrons in the shared pair are provided by the same atom.
Conditions Needed for a Coordinate Bond:
1. One atom must possess a lone pair of electrons (a pair of outer electrons not involved in bonding).
2. The other atom/ion must have an empty orbital (a vacant space) capable of accepting that pair of electrons.
Important Note: Once formed, a coordinate bond has the exact same strength, length, and chemical behavior as any standard covalent bond!
Essential CCEA Examples of Coordinate Bonds:
1. The Ammonium Ion, \( \text{NH}_4^+ \):
An ammonia molecule, \( \text{NH}_3 \), has a nitrogen atom with one lone pair. When it reacts with a hydrogen ion, \( \text{H}^+ \) (which has an empty \( 1s \) orbital and no electrons), nitrogen donates its lone pair to form the coordinate bond:
\( \text{NH}_3 + \text{H}^+ \rightarrow \text{NH}_4^+ \)
2. The Hydronium (Hydroxonium) Ion, \( \text{H}_3\text{O}^+ \):
A water molecule, \( \text{H}_2\text{O} \), has two lone pairs on the oxygen atom. Oxygen donates one lone pair to an empty \( \text{H}^+ \) ion:
\( \text{H}_2\text{O} + \text{H}^+ \rightarrow \text{H}_3\text{O}^+ \)
3. Aluminium Chloride Dimer, \( \text{Al}_2\text{Cl}_6 \):
In gaseous aluminium chloride, \( \text{AlCl}_3 \) is electron deficient (aluminium only has 6 electrons in its outer shell). Two \( \text{AlCl}_3 \) molecules join together: a chlorine atom on each molecule donates a lone pair to the aluminium atom of the neighbouring molecule, forming two coordinate bonds.
Key Takeaway for Covalent Bonding: Standard covalent bonds share electrons mutually; coordinate bonds involve one atom donating both electrons from a lone pair into a vacant orbital.
3. Metallic Bonding
What is a Metallic Bond?
A metallic bond is defined as the electrostatic attraction between positive metal ions (cations) and the sea of delocalised electrons in a giant metallic lattice.
Analogy: Imagine shiny marbles (positive metal ions) sitting tightly packed in a pool of clear honey (delocalised electrons). The honey keeps the marbles firmly stuck together, even if you push them around.
Properties of Metals Explained
1. High Electrical and Thermal Conductivity:
The delocalised electrons are free to move throughout the entire metallic structure. When a voltage is applied, these mobile electrons drift toward the positive terminal, carrying electric current. They also rapidly transfer kinetic thermal energy through the lattice.
2. High Melting and Boiling Points:
Metallic bonds are typically very strong. A significant amount of energy is required to overcome the electrostatic attraction between the metal cations and the delocalised electron cloud.
Trend note: Metallic bond strength increases when the metal ion has a higher positive charge (e.g., \( \text{Al}^{3+} > \text{Mg}^{2+} > \text{Na}^+ \)) and a smaller ionic radius, because the delocalised electrons are held closer and more tightly to the nucleus.
3. Malleability and Ductility:
• Malleable: Can be hammered or rolled into thin sheets.
• Ductile: Can be drawn out into long wires.
Why? When a force is applied to a metal, the layers of positive ions can slide past one another without breaking the bond. The sea of delocalised electrons simply adjusts and continues to hold the ions together.
Key Takeaway for Metallic Bonding: Metallic bonding consists of cations in a sea of delocalised electrons. This explains why metals conduct electricity as solids and are malleable without fracturing.
4. Electronegativity and Bond Polarity
Understanding Electronegativity
Electronegativity is the ability of an atom to attract the bonding pair of electrons in a covalent bond.
Electronegativity is measured using the Pauling scale (ranging from \( 0.7 \) to \( 4.0 \)):
• Fluorine (\( \text{F} \)) is the most electronegative element with a value of \( 4.0 \).
• Other highly electronegative elements include Oxygen (\( 3.5 \)), Nitrogen (\( 3.0 \)), and Chlorine (\( 3.0 \)).
Periodic Trends in Electronegativity:
• Across a Period (left to right): Electronegativity increases. The nuclear charge increases while shielding remains roughly constant, so the atomic radius decreases and bonding electrons are attracted more strongly.
• Down a Group (top to bottom): Electronegativity decreases. Although nuclear charge increases, the number of inner electron shells increases (more shielding) and the atomic radius increases, weakening the pull on the bonding pair.
Bond Polarity
1. Non-Polar Covalent Bonds:
When two identical atoms share electrons (such as \( \text{H}_2 \), \( \text{Cl}_2 \), or \( \text{O}_2 \)), both atoms have identical electronegativity. The electron pair is shared equally between the two nuclei.
2. Polar Covalent Bonds:
When atoms with different electronegativities form a covalent bond (such as in \( \text{H}-\text{Cl} \)), the more electronegative atom pulls the shared pair closer to itself.
• The atom gaining greater electron density gains a partial negative charge (\( \delta^- \)).
• The atom losing electron density gains a partial positive charge (\( \delta^+ \)).
This separation of charge creates a dipole.
Mnemonic Trick: Think of electronegativity as a tug-of-war for electrons. Equal strength = non-polar bond. Unequal strength = polar bond!
5. Intermolecular Forces
It is vital not to confuse intramolecular bonds (covalent bonds inside molecules holding atoms together) with intermolecular forces (weak attractions between separate molecules).
There are three main types of intermolecular forces you must master for CCEA AS Chemistry:
1. Van der Waals' Forces (London Dispersion / Induced Dipole-Dipole)
Van der Waals' forces exist between all atoms and molecules (even noble gases like \( \text{He} \) and non-polar molecules like \( \text{CH}_4 \)).
How they form (Step-by-Step):
1. Electrons inside an atom or molecule are constantly moving.
2. At any random instant, electrons may be unevenly distributed, creating a temporary, instantaneous dipole (\( \delta^+ / \delta^- \)).
3. This instantaneous dipole repels or attracts electrons in a neighbouring atom, inducing a dipole in it.
4. Weak electrostatic attractions arise between these temporary dipoles.
Factors Affecting Van der Waals' Strength:
• Number of electrons: More electrons mean a larger electron cloud, which is more easily polarisable, resulting in stronger induced dipoles and higher boiling points (e.g., boiling points increase down the halogens: \( \text{F}_2 < \text{Cl}_2 < \text{Br}_2 < \text{I}_2 \)).
• Surface area / molecular shape: Long, straight-chain molecules can pack closer together with larger contact surface area, giving stronger Van der Waals' forces than compact, branched isomers.
2. Permanent Dipole-Dipole Forces
These forces occur strictly between polar molecules (molecules with permanent dipoles, such as \( \text{HCl} \) or \( \text{CH}_3\text{Cl} \)). The \( \delta^+ \) end of one polar molecule attracts the \( \delta^- \) end of an adjacent polar molecule.
3. Hydrogen Bonding (The Strongest Intermolecular Force)
Hydrogen bonding is an exceptionally strong permanent dipole-dipole attraction.
Conditions Required for Hydrogen Bonding:
1. A hydrogen atom covalently bonded to an extremely electronegative element: specifically Nitrogen (\( \text{N} \)), Oxygen (\( \text{O} \)), or Fluorine (\( \text{F} \)).
2. A lone pair of electrons on that electronegative atom of an adjacent molecule.
Why is Hydrogen Bonding so Strong?
The large difference in electronegativity creates a very large dipole. Because the hydrogen atom has no inner screening electrons, its nucleus is exposed as an intense spot of positive charge, forming a strong attraction to the lone pair of a neighbouring \( \text{N} \), \( \text{O} \), or \( \text{F} \) atom.
Anomalous Properties of Water due to Hydrogen Bonding:
1. Unusually High Boiling and Melting Points:
Water (\( \text{H}_2\text{O} \)) has a much higher boiling point than other Group 16 hydrides (such as \( \text{H}_2\text{S} \)) because significant thermal energy is required to break the extensive hydrogen bonds between water molecules.
2. Ice is Less Dense than Liquid Water:
As water freezes, hydrogen bonds hold the water molecules in an open, three-dimensional hexagonal crystalline ring lattice. This holds the molecules further apart than in liquid water, making ice less dense so that it floats on water.
Key Takeaway for Intermolecular Forces:
Strength order: Hydrogen bonds > Permanent dipole-dipole > Van der Waals' forces. Intermolecular forces determine physical properties such as melting point and boiling point, not chemical reactivity.
6. Shapes of Simple Molecules and Ions (VSEPR Theory)
Valence Shell Electron Pair Repulsion (VSEPR) Theory
The shapes of molecules and polyatomic ions can be predicted using simple repulsion rules:
1. Electron pairs around the central atom arrange themselves in space to minimise repulsion and get as far apart as possible.
2. Lone pairs repel more strongly than bonding pairs.
Repulsion hierarchy:
Lone Pair – Lone Pair Repulsion > Lone Pair – Bonding Pair Repulsion > Bonding Pair – Bonding Pair Repulsion
3. As a rule of thumb, each lone pair reduces the ideal bond angle by approximately \( 2.5^\circ \).
Summary Table of Molecular Shapes:
Two Electron Pairs around Central Atom:
• 2 Bonding Pairs, 0 Lone Pairs:
Shape: Linear
Bond Angle: \( 180^\circ \)
Examples: \( \text{BeCl}_2 \), \( \text{CO}_2 \)
Three Electron Pairs around Central Atom:
• 3 Bonding Pairs, 0 Lone Pairs:
Shape: Trigonal Planar
Bond Angle: \( 120^\circ \)
Examples: \( \text{BF}_3 \), \( \text{AlCl}_3 \)
Four Electron Pairs around Central Atom:
• 4 Bonding Pairs, 0 Lone Pairs:
Shape: Tetrahedral
Bond Angle: \( 109.5^\circ \)
Examples: \( \text{CH}_4 \), \( \text{NH}_4^+ \), \( \text{CCl}_4 \)
• 3 Bonding Pairs, 1 Lone Pair:
Shape: Pyramidal (Trigonal Pyramidal)
Bond Angle: \( 107^\circ \) (reduced by \( 2.5^\circ \) due to 1 lone pair)
Examples: \( \text{NH}_3 \), \( \text{PCl}_3 \)
• 2 Bonding Pairs, 2 Lone Pairs:
Shape: Non-linear (Bent / V-shaped)
Bond Angle: \( 104.5^\circ \) (reduced by \( 2 \times 2.5^\circ \) due to 2 lone pairs)
Examples: \( \text{H}_2\text{O} \), \( \text{H}_2\text{S} \)
Five Electron Pairs around Central Atom:
• 5 Bonding Pairs, 0 Lone Pairs:
Shape: Trigonal Bipyramidal
Bond Angles: \( 90^\circ \) and \( 120^\circ \)
Example: \( \text{PCl}_5 \)
Six Electron Pairs around Central Atom:
• 6 Bonding Pairs, 0 Lone Pairs:
Shape: Octahedral
Bond Angle: \( 90^\circ \)
Example: \( \text{SF}_6 \)
Predicting Overall Molecular Polarity
To determine if an entire molecule is polar or non-polar:
1. Check if the bonds are polar (electronegativity difference).
2. Look at the symmetry of the molecular shape.
• If the molecule is completely symmetrical (e.g., linear \( \text{CO}_2 \), trigonal planar \( \text{BF}_3 \), tetrahedral \( \text{CCl}_4 \)), the individual dipoles cancel out, making the overall molecule non-polar.
• If the molecule is asymmetrical (e.g., non-linear \( \text{H}_2\text{O} \), pyramidal \( \text{NH}_3 \), or \( \text{CHCl}_3 \)), the dipoles do not cancel, giving a net molecular dipole (the molecule is polar).
7. Giant Covalent Structures (Macromolecules)
Some non-metal substances form massive networks of covalent bonds extending throughout a giant structure:
1. Diamond (Allotrope of Carbon)
• Each carbon atom is covalently bonded to 4 other carbon atoms in a rigid 3D tetrahedral structure.
• Properties: Extremely hard, very high melting point (huge energy needed to break strong covalent bonds), does not conduct electricity (no free ions or delocalised electrons).
2. Graphite (Allotrope of Carbon)
• Each carbon atom is covalently bonded to 3 other carbon atoms in planar hexagonal rings, arranged in layers.
• The fourth outer electron of each carbon atom is delocalised across the layer.
• Layers are held together only by weak Van der Waals' forces.
• Properties: Soft and slippery (layers slide over each other easily, used in lubricants and pencils), high melting point, conducts electricity due to delocalised electrons moving within layers.
3. Graphene
• A single two-dimensional layer of graphite, exactly one atom thick.
• Properties: Incredibly strong, lightweight, and an exceptional conductor of electricity and heat.
Quick Revision Summary Matrix
• Giant Ionic: High melting point, conducts only when molten/aqueous, soluble in water, brittle (e.g., \( \text{NaCl} \), \( \text{MgO} \)).
• Giant Metallic: High melting point, conducts as solid and liquid, malleable and ductile, insoluble (e.g., \( \text{Mg} \), \( \text{Fe} \), \( \text{Cu} \)).
• Giant Covalent: Very high melting point, insoluble, non-conductors (except graphite/graphene) (e.g., Diamond, Graphite, \( \text{SiO}_2 \)).
• Simple Molecular: Low melting and boiling points (held by weak intermolecular forces), non-conductors (e.g., \( \text{I}_2 \), \( \text{H}_2\text{O} \), \( \text{CH}_4 \), \( \text{CO}_2 \)).