Chemistry (9620) | Bonding

3.1.3 Bonding (International AS) - Comprehensive Study Notes

Hello future chemist! This chapter is incredibly important. Why? Because the way atoms stick together (bonding) and the way molecules interact (intermolecular forces) dictate everything we observe about a substance—its melting point, its hardness, and whether it conducts electricity. Understanding bonding is the foundation of all chemistry!

3.1.3.1 Ionic Bonding: The Power of Attraction

Ionic bonding is all about opposites attracting! It occurs between metals and non-metals.

• Definition: Ionic bonding involves strong electrostatic attraction between oppositely charged ions arranged in a regular, repeating pattern called an ionic lattice.
• How it forms: Electrons are transferred (not shared) from the metal atom (forming a positive ion, or cation) to the non-metal atom (forming a negative ion, or anion).

Predicting Ion Charges (Using the Periodic Table)

We can usually predict the charge of a simple ion based on its position in the Periodic Table:

• Group 1: Lose 1 electron, form +1 ions (e.g., Na⁺).
• Group 2: Lose 2 electrons, form +2 ions (e.g., Mg²⁺).
• Group 13 (Aluminium): Loses 3 electrons, forms +3 ions (e.g., Al³⁺).
• Group 16: Gain 2 electrons, form -2 ions (e.g., O^2-).
• Group 17: Gain 1 electron, form -1 ions (e.g., Cl^-).

Constructing Formulas for Ionic Compounds

The overall charge of an ionic compound must be zero (neutral). You must balance the positive and negative charges.

Example: Magnesium (\(Mg^{2+}\)) and Chloride (\(Cl^{-}\))

1. Mg needs a +2 charge. Cl needs a -1 charge.
2. To balance, you need two Cl ions for every one Mg ion.
3. Formula: \(MgCl_2\).

You must also know the formulae of important compound ions (polyatomic ions):

• Sulfate: \((SO_4)^{2-}\)
• Hydroxide: \((OH)^{-}\)
• Nitrate: \((NO_3)^{-}\)
• Carbonate: \((CO_3)^{2-}\)
• Ammonium: \((NH_4)^{+}\) (The only common positive compound ion you need to know!)

3.1.3.2 Nature of Covalent and Dative Covalent Bonds

Covalent Bonding: The Sharing Economy

A covalent bond forms primarily between two non-metal atoms. It involves the sharing of a pair of electrons between the atoms.

• A single covalent bond is one shared pair, represented by a single line (—).
• Multiple bonds (double or triple) involve two or three shared pairs (e.g., \(C=C\) or \(C \equiv C\)).

Dative Covalent (Co-ordinate) Bonds

This is a special type of covalent bond. The sharing happens, but both electrons in the shared pair come from only one of the atoms.

• Representation: A dative bond is shown using an arrow (\(\rightarrow\)). The arrow points away from the atom donating the lone pair and towards the atom accepting it.
• Example: The reaction between ammonia (\(NH_3\)) and a hydrogen ion (\(H^{+}\)) to form an ammonium ion (\(NH_4^{+}\)). The Nitrogen atom in \(NH_3\) donates its lone pair to the empty orbital on the \(H^{+}\) ion.

Don't worry if this seems tricky at first: Once formed, a dative bond is chemically identical to a normal covalent bond; the difference is only in how it originated.

3.1.3.3 Metallic Bonding

Metallic bonding explains why metals are such good conductors and why they are shiny and malleable.

• Structure: Metals consist of a regular lattice of positive ions (cations).
• Bonding: The outer shell electrons are delocalised (free to move) throughout the structure.
• Definition: Metallic bonding is the strong electrostatic attraction between the positive metal ions and the surrounding sea of delocalised electrons.

3.1.3.5 Shapes of Simple Molecules and Ions (VSEPR Theory)

The Valence Shell Electron Pair Repulsion (VSEPR) Theory dictates the shape of a molecule. Electrons, being negatively charged, repel each other. They arrange themselves as far apart as possible to minimise this repulsion.

Electron Pair Repulsion Hierarchy

Not all electron pairs repel equally. Lone pairs (non-bonding electrons) take up more space and cause stronger repulsion than bonding pairs (BP).

Repulsion Strength:

1. Lone Pair - Lone Pair (LP-LP) is the strongest repulsion.
2. Lone Pair - Bonding Pair (LP-BP) is intermediate.
3. Bonding Pair - Bonding Pair (BP-BP) is the weakest repulsion.

This unequal repulsion is why bond angles in molecules with lone pairs (like water or ammonia) are slightly smaller than the ideal angles (like in methane).

Step-by-Step Guide to VSEPR

1. Count the total number of electron pairs (bonding pairs + lone pairs) around the central atom.
2. Determine the basic geometric arrangement that minimises repulsion (electron pair geometry).
3. Determine the final molecular shape based only on the positions of the atoms (molecular geometry).

Common Shapes (Up to six electron pairs)

Total Pairs	BP	LP	Electron Geometry	Molecular Shape	Bond Angle (approx.)
2	2	0	Linear	Linear	\(180^{\circ}\)
3	3	0	Trigonal planar	Trigonal planar	\(120^{\circ}\)
3	2	1	Trigonal planar	V-shaped (Bent)	< \(120^{\circ}\) (e.g., \(SO_2\))
4	4	0	Tetrahedral	Tetrahedral	\(109.5^{\circ}\) (e.g., \(CH_4\))
4	3	1	Tetrahedral	Trigonal pyramidal	\(107^{\circ}\) (e.g., \(NH_3\))
4	2	2	Tetrahedral	V-shaped (Bent)	\(104.5^{\circ}\) (e.g., \(H_2O\))
5	5	0	Trigonal bipyramidal	Trigonal bipyramidal	\(90^{\circ}, 120^{\circ}\)
6	6	0	Octahedral	Octahedral	\(90^{\circ}\)

3.1.3.6 Bond Polarity and Electronegativity

Electronegativity

Electrons in a covalent bond are not always shared equally. The concept of electronegativity helps us understand this.

• Definition: Electronegativity is the power of an atom to attract the shared pair of electrons in a covalent bond.
• Trends: Electronegativity generally increases across a period (atoms get smaller) and decreases down a group (atoms get larger). Fluorine is the most electronegative element.

Polar Covalent Bonds

When two atoms have different electronegativities, the electron distribution is unsymmetrical, creating a polar covalent bond.

• The more electronegative atom attracts the electrons more strongly and gains a small negative charge (\(\delta-\)).
• The less electronegative atom gains a small positive charge (\(\delta+\)).
• Showing Polarity: Use partial charges \(\delta+\) and \(\delta-\).

Molecular Dipoles (Permanent Dipoles)

A molecule may contain polar bonds, but the molecule as a whole may still be non-polar. This depends on its symmetry.

• If the molecule is asymmetrical (e.g., \(HCl\), \(H_2O\), \(NH_3\)), the dipoles do not cancel out, and the molecule has a net permanent dipole.
• If the molecule is highly symmetrical (e.g., \(CO_2\), \(CH_4\), \(CCl_4\)), the individual bond dipoles pull equally in opposite directions and cancel out. The molecule has no permanent dipole, even though it contains polar bonds.

Analogy: Think of a tug-of-war. If two equally strong teams pull exactly opposite directions (\(CO_2\) is linear, \(C \leftarrow O\) and \(C \rightarrow O\)), the rope (the molecule) doesn't move. If the pulls are uneven or angled (\(H_2O\) is bent), the rope moves, creating a net force (a permanent dipole).

3.1.3.7 Forces Between Molecules (Intermolecular Forces)

Intermolecular forces (IMFs) are weak forces of attraction between neighbouring molecules. They are crucial because they determine the melting and boiling points of molecular substances.

1. Induced Dipole–Dipole Forces (van der Waals / Dispersion / London Forces)

• These exist between all molecules (polar and non-polar).
• They arise from the constant, random movement of electrons. At any instant, the electron distribution may be uneven, creating a temporary, instantaneous dipole.
• This temporary dipole induces a dipole in a neighbouring molecule, creating a weak attraction.
• Strength: These forces increase with the number of electrons (or size/mass) of the molecule. Larger molecules have stronger van der Waals forces.

2. Permanent Dipole–Dipole Forces

• These exist only between polar molecules (molecules with a permanent dipole).
• The positive end of one polar molecule attracts the negative end of a neighbouring polar molecule.
• Strength: Stronger than van der Waals forces for molecules of similar size.

3. Hydrogen Bonding (The Strongest IMF)

Hydrogen bonding is a special, very strong type of dipole-dipole attraction that occurs only when hydrogen (H) is covalently bonded to one of the three most electronegative atoms: Nitrogen (N), Oxygen (O), or Fluorine (F) (The "NOF" rule or memory aid).

• The H atom develops a large \(\delta+\) charge, and the N, O, or F atom develops a large \(\delta-\) charge and still possesses a lone pair of electrons.
• The strong attraction is between the \(\delta+\) hydrogen atom and a lone pair on the N, O, or F atom of a neighbouring molecule.

The Importance of Hydrogen Bonding

Hydrogen bonding leads to anomalous properties, especially in water:

• Anomalous Boiling Points: Compounds like \(H_2O\), \(HF\), and \(NH_3\) have much higher boiling points than expected compared to similar molecules without H-bonds (e.g., \(H_2S\), \(HCl\)). This is because extra energy is needed to break the strong H-bonds.
• Low Density of Ice: When water freezes, H-bonds hold the molecules apart in a rigid, open lattice structure. This structure occupies more volume than liquid water, making ice less dense (which is why ice floats!).

3.1.3.4 Bonding and Physical Properties: Structure Determines Destiny

The physical properties (like melting point and conductivity) of a substance depend entirely on its structure and the type of bonding present.

The Four Types of Crystal Structure

We classify structures into four main types:

1. Molecular Structure (Simple Molecules)
2. Macromolecular Structure (Giant Covalent)
3. Ionic Structure (Giant Ionic Lattice)
4. Metallic Structure (Giant Metallic Lattice)

Relating Structure to Melting Point and Conductivity

1. Molecular Structures

• Examples: Ice (\(H_2O\)), Iodine (\(I_2\)).
• Bonding within molecules: Strong covalent bonds.
• Forces between molecules: Weak intermolecular forces (van der Waals, dipole-dipole, or H-bonds).
• Melting/Boiling Points: Low. Only weak IMFs need to be overcome to change state (energy changes are small).
• Conductivity: Do not conduct electricity in any state (no mobile charged particles).

2. Macromolecular (Giant Covalent) Structures

• Examples: Diamond, Graphite, Silicon Dioxide (\(SiO_2\)).
• Bonding: Strong covalent bonds extend throughout the entire structure.
• Melting/Boiling Points: Extremely high. Massive amounts of energy are required to break the thousands of strong covalent bonds.
• Conductivity:
  • Diamond: Non-conductor (all valence electrons are locked up in four covalent bonds).
  • Graphite: Conductor. It has layers with delocalised electrons between them, allowing charge to flow.

3. Ionic Structures

• Example: Sodium Chloride (\(NaCl\)).
• Bonding: Strong electrostatic attraction (ionic bonds) in a giant lattice.
• Melting/Boiling Points: High. Large amounts of energy needed to overcome strong attractions.
• Conductivity:
  • Solid: Non-conductor (ions are fixed in the lattice).
  • Molten (liquid) or Aqueous (dissolved): Conductors (ions become mobile and carry the charge).

4. Metallic Structures

• Example: Magnesium (\(Mg\)).
• Bonding: Strong metallic bonding (attraction between cations and delocalised electrons).
• Melting/Boiling Points: Generally high (depends on charge density and size of the ion, but stronger than molecular).
• Conductivity: Excellent conductors in solid and liquid states (due to the mobile delocalised electrons).

Energy Changes Associated with Changes of State

When a substance changes state (e.g., solid to liquid, or liquid to gas), we are measuring the energy needed to overcome the forces holding the particles together.

• Melting/Boiling (Solid \(\rightarrow\) Liquid \(\rightarrow\) Gas): This requires energy (endothermic). Energy is used to overcome the IMFs or, in giant structures, the chemical bonds.
• Condensing/Freezing (Gas \(\rightarrow\) Liquid \(\rightarrow\) Solid): This releases energy (exothermic). Energy is released as the particles form attractions (IMFs or bonds).