Welcome to Chemical Bonding!
Welcome to one of the most fundamental chapters in Chemistry! Have you ever wondered why table salt forms beautiful hard crystals, why water is a liquid at room temperature while oxygen is a gas, or why diamonds are so incredibly hard? The answer to all of these questions comes down to chemical bonding.
Atoms are sociable particles. Most of them do not like being on their own. They bond with each other to achieve a more stable, lower-energy state (often by getting a full outer shell of electrons). In this chapter, we will explore how atoms interact, the shapes molecules make in 3D space, and the forces that hold everything together.
1. Ionic Bonding
What is an Ionic Bond?
An ionic bond is the electrostatic attraction between oppositely charged ions. It forms when electrons are transferred from a metal atom to a non-metal atom.
• Metals lose electrons to form positive ions (cations).
• Non-metals gain electrons to form negative ions (anions).
Analogy: Think of ionic bonding like two oppositely charged magnetic toys snapping together. Once the transfer of electrons happens, the positive and negative charges create a powerful, non-directional pull on each other.
Dot-and-Cross Diagrams for Ionic Compounds
When drawing ionic compounds:
1. Draw square brackets around each ion.
2. Show the outer shell electrons only.
3. Use dots for electrons from one atom and crosses for electrons from another.
4. Write the overall charge as a superscript outside the top right of the bracket (e.g., \( [\text{Na}]^+ \) and \( [\text{Cl}]^- \)).
For magnesium oxide, \( \text{Mg} \) loses two electrons: \( [\text{Mg}]^{2+} \), and \( \text{O} \) gains two electrons: \( [\text{O}]^{2-} \).
Giant Ionic Lattices
Ionic compounds do not exist as isolated single pairs of ions. Instead, millions of cations and anions pack tightly together in a regular repeating 3D arrangement called a giant ionic lattice (such as in sodium chloride, \( \text{NaCl} \)). In sodium chloride, each \( \text{Na}^+ \) ion is surrounded by six \( \text{Cl}^- \) ions, and each \( \text{Cl}^- \) ion is surrounded by six \( \text{Na}^+ \) ions.
Physical Properties of Ionic Compounds
• High Melting and Boiling Points: Large amounts of thermal energy are required to overcome the strong electrostatic attractions acting in all directions throughout the giant lattice.
• Electrical Conductivity:
- Solid state: They do not conduct electricity because the ions are fixed in place in the lattice and cannot move.
- Molten or Aqueous solution: They do conduct electricity because the lattice breaks down and the ions are free to move and carry charge.
• Brittleness: When hit with a sharp force, layers of ions slide over each other. Ions of the same charge are brought next to each other, resulting in strong repulsion that shatters the crystal.
Common Mistake to Avoid: Never say "electrons move to conduct electricity in molten ionic compounds." It is always the mobile ions that carry the current!
Key Takeaway: Ionic bonding involves the transfer of electrons between metals and non-metals, forming a rigid 3D lattice with strong electrostatic forces, leading to high melting points and conductivity only when melted or dissolved.
2. Covalent and Coordinate (Dative) Bonding
What is a Covalent Bond?
A covalent bond is the electrostatic attraction between a shared pair of electrons and the positively charged nuclei of the bonded atoms. Covalent bonds typically form between non-metal atoms.
• Single bond: One shared pair of electrons (e.g., \( \text{H}_2 \), \( \text{Cl}_2 \)).
• Double bond: Two shared pairs of electrons (e.g., \( \text{O}_2 \), \( \text{CO}_2 \)).
• Triple bond: Three shared pairs of electrons (e.g., \( \text{N}_2 \)).
Coordinate (Dative Covalent) Bonding
A coordinate (or dative covalent) bond is a special type of covalent bond where both shared electrons are donated by the same atom.
For a coordinate bond to form, you need:
1. An atom with a lone pair of electrons (a pair of outer electrons not involved in bonding).
2. An atom or ion with an empty orbital (a vacant space to accept the electron pair).
Example 1: The Ammonium Ion (\( \text{NH}_4^+ \))
Ammonia (\( \text{NH}_3 \)) has a nitrogen atom with one lone pair. A hydrogen ion (\( \text{H}^+ \)) has an empty orbital (no electrons). Nitrogen donates its lone pair to the \( \text{H}^+ \) to form an ammonium ion:
\( \text{NH}_3 + \text{H}^+ \rightarrow \text{NH}_4^+ \)
Once formed, a coordinate bond is identical in strength and length to any standard covalent bond. In structural formulas, coordinate bonds are represented by an arrow (\( \rightarrow \)) pointing from the donor atom to the acceptor atom.
Example 2: Aluminium Chloride Dimer (\( \text{Al}_2\text{Cl}_6 \))
In gaseous aluminium chloride (\( \text{AlCl}_3 \)), the aluminium atom has only six outer electrons (an incomplete octet). Two \( \text{AlCl}_3 \) molecules join together: a chlorine atom from each molecule donates a lone pair to the aluminium atom of the other molecule, completing their octets.
Key Takeaway: Covalent bonds are shared pairs of electrons attracted to both nuclei. Dative covalent bonds occur when one atom supplies both electrons in the shared pair.
3. Metallic Bonding
What is a Metallic Bond?
A metallic bond is the electrostatic attraction between positive metal ions and a 'sea' of delocalised electrons.
Metal atoms lose their outer shell electrons to form a regular lattice of positive metal cations. The outer electrons become delocalised, meaning they do not belong to any single atom and are free to drift throughout the entire structure.
Properties of Metals
• Electrical and Thermal Conductivity: The delocalised electrons are free to move throughout the structure and carry electrical charge or transfer kinetic energy (heat).
• High Melting and Boiling Points: Strong electrostatic attractions between cations and delocalised electrons require significant energy to break.
• Malleability and Ductility: Metals can be hammered into sheets (malleable) and drawn into wires (ductile). Because the delocalised electrons move freely, the layers of positive ions can slide past each other without breaking the metallic bond.
Strength of Metallic Bonding: Metallic bonding is stronger when ions have a higher positive charge (e.g., \( \text{Al}^{3+} > \text{Mg}^{2+} > \text{Na}^+ \)) and smaller ionic radii, because this increases the charge density and attracts the delocalised electrons more strongly.
Key Takeaway: Metals consist of positive ions in a sea of free delocalised electrons, giving them high electrical conductivity, strength, and malleability.
4. Shapes of Molecules and Ions (VSEPR Theory)
Valence Shell Electron Pair Repulsion Theory
Molecules adopt 3D shapes to minimise repulsion between electron pairs around the central atom. The fundamental principles are:
1. Electron pairs in the outer shell of the central atom repel each other.
2. Electron pairs arrange themselves as far apart as possible to minimise repulsion.
3. Lone pairs repel more strongly than bonding pairs.
Hierarchy of repulsion:
\( \text{Lone Pair - Lone Pair} > \text{Lone Pair - Bonding Pair} > \text{Bonding Pair - Bonding Pair} \)
Rule of thumb: Each lone pair reduces the ideal bond angle by approximately \( 2.5^\circ \).
Step-by-Step Guide to Working Out Molecular Shape
1. Count the number of outer electrons on the central atom.
2. Add one electron for each atom bonded to it (and adjust for any ionic charge: add for negative charge, subtract for positive charge).
3. Divide by 2 to get the total number of electron pairs.
4. Determine the number of bonding pairs (equal to bonded atoms) and lone pairs (total pairs minus bonding pairs).
5. State the shape and bond angle.
Summary of Shapes and Angles
• 2 electron pairs (2 bonding, 0 lone):
Shape: Linear
Bond angle: \( 180^\circ \)
Example: \( \text{BeCl}_2 \), \( \text{CO}_2 \)
• 3 electron pairs (3 bonding, 0 lone):
Shape: Trigonal Planar
Bond angle: \( 120^\circ \)
Example: \( \text{BF}_3 \), \( \text{AlCl}_3 \)
• 4 electron pairs (4 bonding, 0 lone):
Shape: Tetrahedral
Bond angle: \( 109.5^\circ \)
Example: \( \text{CH}_4 \), \( \text{NH}_4^+ \)
• 4 electron pairs (3 bonding, 1 lone):
Shape: Pyramidal (Trigonal Pyramidal)
Bond angle: \( 107^\circ \) (reduced by \( 2.5^\circ \))
Example: \( \text{NH}_3 \), \( \text{H}_3\text{O}^+ \)
• 4 electron pairs (2 bonding, 2 lone):
Shape: Non-linear (Bent / V-shaped)
Bond angle: \( 104.5^\circ \) (reduced by \( 2 \times 2.5^\circ = 5^\circ \))
Example: \( \text{H}_2\text{O} \)
• 5 electron pairs (5 bonding, 0 lone):
Shape: Trigonal Bipyramidal
Bond angles: \( 90^\circ \) and \( 120^\circ \)
Example: \( \text{PCl}_5 \)
• 6 electron pairs (6 bonding, 0 lone):
Shape: Octahedral
Bond angle: \( 90^\circ \)
Example: \( \text{SF}_6 \)
Key Takeaway: Electron pairs repel to get as far apart as possible. Lone pairs repel more than bonding pairs, squeezing bond angles smaller.
5. Electronegativity and Bond Polarity
What is Electronegativity?
Electronegativity is the power of an atom to attract the bonding pair of electrons in a covalent bond.
Electronegativity is measured on the Pauling scale (from 0 to 4.0). Fluorine is the most electronegative element (\( 4.0 \)), followed by oxygen (\( 3.5 \)), nitrogen (\( 3.0 \)), and chlorine (\( 3.0 \)).
• Across a period: Electronegativity increases (nuclear charge increases, atomic radius decreases, shielding stays similar).
• Down a group: Electronegativity decreases (atomic radius increases, shielding increases, weaker pull on bonding electrons).
Polar Bonds and Dipoles
• Non-polar bond: Formed between atoms with identical or very similar electronegativity (e.g., \( \text{H}_2 \), \( \text{Cl}_2 \), \( \text{C}-\text{H} \)). The electrons are shared equally.
• Polar bond: Formed between atoms with differing electronegativities (e.g., \( \text{H}-\text{Cl} \)). The more electronegative atom pulls electron density toward itself, gaining a slight negative charge (\( \delta^- \)), leaving the other atom with a slight positive charge (\( \delta^+ \)). This separation of charge is called a permanent dipole.
Polar vs. Non-Polar Molecules
A molecule can have polar bonds but still be non-polar overall if its shape is symmetrical, causing the dipoles to cancel out!
• Carbon Dioxide (\( \text{CO}_2 \)): Linear shape. The two \( \text{C}=\text{O} \) dipoles point in directly opposite directions and cancel out. \( \text{CO}_2 \) is a non-polar molecule.
• Water (\( \text{H}_2\text{O} \)): Bent shape. The two \( \text{O}-\text{H} \) dipoles do not cancel out. \( \text{H}_2\text{O} \) has an overall dipole and is a polar molecule.
Key Takeaway: Unequal sharing of electrons causes polar bonds. Symmetrical shapes cancel dipoles (non-polar molecule), while asymmetrical shapes do not cancel dipoles (polar molecule).
6. Intermolecular Forces
Intermolecular forces are weak attractions between separate molecules. Do not confuse them with covalent bonds, which are strong attractions inside molecules.
1. Van der Waals' Forces (London Dispersion / Temporary Dipole-Induced Dipole Forces)
These forces exist between all atoms and molecules.
• Electrons are in continuous, random motion.
• At any single instant, electrons may be unevenly distributed, creating an instantaneous (temporary) dipole.
• This temporary dipole induces a dipole in a neighbouring atom or molecule.
• The two dipoles attract each other weakly.
Factor affecting strength: Larger molecules with more electrons have larger electron clouds that are more easily distorted (polarised), leading to stronger London forces and higher boiling points (e.g., \( \text{I}_2 > \text{Br}_2 > \text{Cl}_2 > \text{F}_2 \)).
2. Permanent Dipole-Dipole Attractions
These occur only between polar molecules (molecules with permanent dipoles, like \( \text{HCl} \) or \( \text{CH}_3\text{Cl} \)). The \( \delta^+ \) end of one molecule attracts the \( \delta^- \) end of another. These are generally stronger than London forces for molecules of similar size.
3. Hydrogen Bonding
Hydrogen bonding is the strongest type of intermolecular force. It is a special, strong type of permanent dipole-dipole attraction.
A hydrogen bond occurs when:
1. A hydrogen atom is directly bonded to a highly electronegative atom: Nitrogen (\( \text{N} \)), Oxygen (\( \text{O} \)), or Fluorine (\( \text{F} \)).
2. The electronegative atom contains at least one lone pair of electrons.
Mnemonic: Remember that hydrogen bonding is "N-O-F" (pronounced like "enough") fun without lone pairs!
Anomalous Properties of Water due to Hydrogen Bonding
• Unusually High Melting and Boiling Points: Water has unexpectedly high melting and boiling points compared to other Group 16 hydrides (like \( \text{H}_2\text{S} \)) because extra energy is needed to break the extensive hydrogen bond network.
• Ice is Less Dense than Liquid Water: As water freezes, the hydrogen bonds hold the \( \text{H}_2\text{O} \) molecules in a rigid, open, tetrahedral lattice. The molecules are spaced further apart than in liquid water, making ice float.
Key Takeaway: Intermolecular forces rank in strength: Hydrogen bonding > Permanent dipole-dipole > London forces. Hydrogen bonds only occur with \( \text{H} \) bonded directly to \( \text{N} \), \( \text{O} \), or \( \text{F} \).
7. Summary of Crystal Structures
Substances can be classified into four distinct physical structure types:
1. Giant Ionic (e.g., \( \text{NaCl} \), \( \text{MgO} \))
• Particles: Positive and negative ions.
• Forces: Strong ionic bonds (electrostatic attraction).
• Properties: High melting points, soluble in water, conduct electricity when molten or aqueous.
2. Giant Metallic (e.g., \( \text{Mg} \), \( \text{Fe} \), \( \text{Cu} \))
• Particles: Positive metal ions and delocalised electrons.
• Forces: Strong metallic bonds.
• Properties: High melting points, insoluble, conduct electricity in solid and liquid states, malleable.
3. Giant Covalent / Macromolecular (e.g., Diamond, Graphite, Graphene)
• Diamond: Each carbon forms 4 strong covalent bonds in a tetrahedral lattice. It is extremely hard, has a very high melting point, and does not conduct electricity.
• Graphite: Each carbon forms 3 covalent bonds in hexagonal layers, leaving one delocalised electron per carbon. Layers slide over each other (soft/lubricant) held by weak forces. It conducts electricity along the layers.
• Graphene: A single layer of graphite (one atom thick). Incredibly strong, lightweight, and an excellent electrical conductor.
4. Simple Molecular (e.g., \( \text{I}_2 \), \( \text{H}_2\text{O} \), \( \text{CO}_2 \), \( \text{CH}_4 \))
• Particles: Molecules.
• Forces: Strong covalent bonds within molecules, weak intermolecular forces between molecules.
• Properties: Low melting and boiling points (only weak intermolecular forces need to be broken), poor electrical conductors (no free ions or delocalised electrons).
Final Review Tip: When asked why simple molecules (like iodine or methane) have low melting points, always state that weak intermolecular forces are broken, NOT the strong covalent bonds!