Bonding

Ionic Bonding

• Ionic bonding: Electrostatic attraction between oppositely charged ions formed by electron transfer.

• Ionic crystals: Giant lattices of ions - Strong electrostatic forces of attraction between oppositely charged ions in all directions of the lattice

• Bond strength and melting point: Increase when ions are smaller and/or have higher charges.

Ionic Radii

• Positive ions: Smaller than their atoms because they lose an electron shell, increasing the nuclear attraction on the remaining electrons.

• Negative ions: Larger than their atoms because they gain electrons, increasing electron repulsion and reducing the nuclear attraction per electron.

• Isoelectronic ions: N³⁻, O²⁻, F⁻, Na⁺, Mg²⁺ and Al³⁺ have the same electronic structure as Ne. Increasing proton number increases nuclear attraction, so ionic radius decreases.

• Down a group: Ionic radius increases because ions have more electron shells.

Ionic Properties

• Melting and boiling points: High due to strong electrostatic forces between oppositely charged ions in a giant lattice. Lots of energy to overcome.

• Solubility in water: Generally good - water molecules are polar.

• Conductivity when solid: Poor because ions are fixed in the lattice and cannot move.

• Conductivity when molten: Good because ions are free to move and carry charge.

• General description: Crystalline solids with a giant ionic lattice.

Covalent bond

shared pair of electrons

Dative Covalent bonding

• Dative covalent (coordinate) bond: A shared pair of electrons comes from only one of the bonding atoms.

• Arrow direction: From the atom providing the lone pair to the electron-deficient atom.

• Shape: A dative covalent bond acts like an ordinary covalent bond when determining shape. For example, NH₄⁺ has a tetrahedral shape.

Covalent Properties / Molecular (simple)

• Melting and boiling points: Simple molecular substances have low melting and boiling points due to weak intermolecular forces. Macromolecular substances have high melting and boiling points because many strong covalent bonds must be broken.

• Solubility in water: Simple molecular substances are generally poorly soluble; macromolecular substances are insoluble.

• Conductivity when solid: Simple molecular substances do not conduct because they have no mobile ions or delocalised electrons. Diamond and sand do not conduct because their electrons are localised; graphite conducts due to delocalised electrons between layers.

• Conductivity when molten: Simple molecular and macromolecular substances generally do not conduct electricity.

• General description: Simple molecular substances are mostly gases and liquids; macromolecular substances are usually solids.

Simple molecular

Covalent bonding

With intermolecular forces (van der Waals, permanent dipoles, hydrogen bonds) between molecules

E.g Iodine Ice

Macromolecular - Graphite

Graphene

• Structure: Each carbon atom forms three covalent bonds with three other carbon atoms, creating layers of hexagonal rings.

• Delocalised electrons: The fourth electron in each carbon atom's highest energy level becomes delocalised and is free to move throughout the structure, carrying electrical charge.

• Softness and slipperiness: Weak intermolecular forces between the layers allow them to slide over each other easily, making graphite soft and slippery. Graphite is used as a lubricant.

• Low density: Layers are far apart compared with the length of covalent bonds, making graphite lightweight.

• High melting point: Strong covalent bonds require a lot of energy to break.

• Insolubility: Strong covalent bonds make graphite insoluble in water.

• Graphene is a single layer of graphite, only one atom thick. It is very strong, lightweight and conducts electricity due to its delocalised electrons.

Macromolecular Diamond

• High melting point: Strong covalent bonds between carbon atoms require a lot of energy to break.

• Hardness: Each carbon atom forms four strong covalent bonds with four other carbon atoms, creating a rigid, three-dimensional giant covalent lattice.

• Electrical conductivity: Diamond does not conduct electricity because all four outer-shell electrons are used in covalent bonds, so there are no delocalised electrons.

• Insolubility: Strong covalent bonds make diamond insoluble in water.

Metallic Bonding

Electrostatic attraction between positive metal ions and delocalised electrons.

• Factors affecting strength:

  1. Number of protons: More protons → stronger nuclear attraction.

  2. Number of delocalised electrons: More delocalised electrons per atom → stronger bonding.

  3. Size of ions: Smaller ions → stronger attraction.

• Example (Mg vs Na): Mg has stronger metallic bonding and a higher melting point because it has more protons, more delocalised electrons per atom and smaller positive ions. Therefore, more energy is needed to overcome the metallic bonding.

• More delocalised electrons mean stronger electrostatic attraction between the positive metal ions and the delocalised electrons, so metallic bonding is stronger.

Metallic properties

• Melting and boiling points: High due to strong electrostatic attraction between positive metal ions and delocalised electrons.

• Solubility in water: Insoluble.

• Conductivity when solid: Good because delocalised electrons can move through the structure and carry charge.

• Conductivity when molten: Good because delocalised electrons remain free to move and carry charge.

• General description: Shiny solids that are malleable because layers of positive ions can slide over one another while the metallic bonding is maintained

Shapes of molecules

• Bond length – Distance between the nuclei of two bonded atoms, determined using X-ray diffraction.

• Bond angle – Estimated angle between two covalent bonds using VSEPR theory.

• Wedge – Shows a bond coming out of the plane of the paper towards you.

• Dotted line – Shows a bond going away from you, behind the plane of the paper.

• Straight line – Shows a bond in the same plane as the paper.

Molecular Shapes and Bond Angles

Shape	Bonding pairs	Lone pairs	Bond angle
Linear	2	0	180°
Trigonal planar	3	0	120°
Tetrahedral	4	0	109.5°
Trigonal pyramidal	3	1	107°
Bent	2	2	104.5°
Trigonal bipyramidal	5	0	90° and 120°
Octahedral	6	0	90

Shape	Bonding pairs	Lone pairs	Bond angle
Square planar	4	2*	90°

How to use VESPR Theory

1. Identify the central atom.

2. Work out the number of electrons in the outer shell of the central atom.

3. Adjust for the overall charge of the ion: add electrons for a negative charge or subtract electrons for a positive charge.

4. Add one electron for each bond formed by the central atom.

5. Divide the total number of electrons by 2 to calculate the total number of electron pairs.

6. Identify the bonding pairs by counting the number of atoms bonded to the central atom. The remaining pairs are lone pairs.

Example: H₂O

• Oxygen is the central atom.

• Oxygen has 6 outer-shell electrons.

• Oxygen forms 2 bonds with hydrogen, giving 2 bonding pairs.

• The remaining 4 electrons form 2 lone pairs.

How to explain shape - VESPR Theory

1. State the number of bonding pairs and lone pairs of electrons around the central atom.

2. Electron pairs repel each other and arrange themselves as far apart as possible to minimise repulsion.

3. If there are no lone pairs, the bonding pairs repel equally.

4. If lone pairs are present, lone pairs repel more strongly than bonding pairs.

5. State the molecular shape and bond angle. Lone pairs reduce bond angles because they repel more strongly than bonding pairs.

Electronegativity

• Electronegativity – The relative tendency of an atom in a covalent bond to attract the bonding pair of electrons towards itself./ the power of an atom to attract electron density in a covalent bond towards itself

• ability of an atom to attract a pair of electrons towards itself in a covalent bond

• Pauling scale – Electronegativity is measured on the Pauling scale, which ranges from 0 to 4.

• Most electronegative atoms – F, O, N and Cl are among the most electronegative atoms.

• Fluorine – The most electronegative element, with a value of 4.0 on the Pauling scale. Greatest electron density.

Factors that affect Electronegativity

• Across a period: Electronegativity increases because the number of protons increases while atomic radius decreases, so the bonding electrons are attracted more strongly by the nucleus.

• Down a group: Electronegativity decreases because the distance between the nucleus and bonding electrons increases, and shielding from inner electron shells increases.

• Nuclear charge / distance between the nucleus and the outer-shell electrons / shielding of the nucleus by electrons in inner energy levels.

Intermediate Bonding

• Polar covalent bond: Forms when bonded atoms have different electronegativities (typically a difference of around 0.3–1.7).

• Unequal electron distribution: The bonding electrons are attracted more strongly towards the more electronegative atom, creating a separation of charge called a dipole.

• Partial charges: The more electronegative atom becomes δ−δ−, while the less electronegative atom becomes δ+δ+.

• The more electronegative atom has a greater electron density so electrons are more attracted to higher electron density.

• Permanent dipole: A polar covalent bond has a permanent dipole because of this unequal distribution of electrons - shift in electron density.

• Like dissolves like

Symmetric molecules - Polar and Non-polar molecules

• Symmetrical molecules: A molecule with identical bonds arranged symmetrically may be non-polar, even if its individual bonds are polar.

• Dipole cancellation: Individual bond dipoles cancel out because of the symmetrical shape of the molecule.

• Net dipole moment: If all bond dipoles cancel out, there is no net dipole moment, so the molecule is non-polar.

• Like dissolves like

Examples:

• CCl₄ – Non-polar because its tetrahedral shape is symmetrical, so the bond dipoles cancel out.

• CO₂ – Non-polar because its linear shape is symmetrical, so the bond dipoles cancel out.

• CH₃Cl – Polar because its tetrahedral shape is not symmetrical, so the bond dipoles do not cancel out.

Van der Waals’ Forces

• Van der Waals forces occur between all simple covalent molecules and individual atoms in noble gases. They do not occur between ions in ionic substances.

• They are also called transient, induced dipole–dipole interactions.

• Formation of temporary dipoles: Electrons move constantly and randomly, causing the electron density to fluctuate. This creates small, temporary dipoles.

• Formation of induced dipoles: Temporary dipoles can induce dipoles in neighbouring molecules by repelling or attracting their electrons.

• Attraction between dipoles: The δ+δ+ end of one dipole attracts the δ−δ− end of a neighbouring dipole.

Factors Affecting the Strength of Van der Waals Forces

• Number of electrons: The more electrons a molecule has, the greater the chance of forming temporary dipoles. This produces stronger Van der Waals forces, resulting in higher boiling points.

• Halogens down Group 7: Boiling points increase down the group because larger molecules have more electrons, resulting in stronger Van der Waals forces. This explains why I₂ is a solid whereas Cl₂ is a gas at room temperature.

• Alkanes: Boiling points increase across the homologous series as molecules get larger and contain more electrons, resulting in stronger Van der Waals forces.

• Molecular shape: Long-chain alkanes have a larger surface area of contact between molecules than spherical, branched alkanes. This results in stronger Van der Waals forces and higher boiling points.

Permanent dipole-dipole forces

• Occurrence: Permanent dipole–dipole forces occur between polar molecules.

• Permanent dipoles: Polar molecules have permanent dipoles due to an unequal distribution of electrons.

• Polarity: Polar molecules have a significant difference in electronegativity between bonded atoms and an asymmetrical arrangement of bond dipoles.

• Examples of polar bonds: C–Cl, C–F, C–Br, H–Cl and C=O.

• Strength: Permanent dipole–dipole forces are generally stronger than Van der Waals forces, resulting in higher boiling points for molecules of similar size and shape.

• Combined forces: Permanent dipole–dipole forces occur in addition to Van der Waals forces in polar molecules.

Hydrogen Bodning

• Occurrence: Hydrogen bonding occurs when a hydrogen atom is covalently bonded to nitrogen, oxygen or fluorine, which has a lone pair of electrons available to form hydrogen bonds.

• Examples of bonds: O–H, N–H and H–F.

• Electronegativity: There is a large electronegativity difference between hydrogen and N, O or F, resulting in a strongly polar bond.

• Diagrams: Always show the lone pairs on N, O or F, the bond dipoles and all δ+δ+ and δ−δ− charges.

• Combined forces: Hydrogen bonding occurs in addition to Van der Waals forces and, in polar molecules, permanent dipole–dipole forces.

• Strength: Hydrogen bonding is stronger than the other two types of intermolecular forces.

• Anomalously high boiling points: H₂O, NH₃ and HF have unusually high boiling points due to hydrogen bonding between their molecules.

• Group 16 hydrides: The general increase in boiling point from H₂S to H₂Te is caused by stronger Van der Waals forces as the number of electrons increases.

• Other substances: Alcohols, carboxylic acids, proteins and amides can form hydrogen bonds.

Need to be able to draw…

Solid/Liquid/Gas

Property	Solid	Liquid	Gas
Particle spacing	Very close together	Close together	Far apart
Density	Usually high	Similar to solids	Very low
Compressibility	Virtually incompressible	Virtually incompressible	Compressible
Particle movement	Vibrate about a fixed point	Move freely and randomly	Move freely and rapidly in random directions
Forces of attraction	Strong	Weaker than in solids	Very weak/negligible

Changing State: Solid to Liquid

• Energy is supplied to the particles.

• The particles gain energy and vibrate more.

• At the melting point, energy is used to overcome the forces of attraction holding the particles in fixed positions, allowing them to move past one another.