Chemistry (9701) Study Notes: Bonding and Structure (Topics 3 & 4.2)
Hello future Chemists!
Welcome to the core of Chemistry: Bonding and Structure. This chapter explains the "why" behind every substance you encounter—why salt is hard, why water boils at a high temperature, and why diamonds are so strong. Understanding how atoms stick together is crucial because the type of bond dictates almost all of a substance's physical and chemical properties. Don't worry if some concepts seem theoretical; we will break them down using simple analogies and clear steps!
Section 1: Electronegativity and the Main Bond Types (3.1 - 3.3)
3.1 Electronegativity: The Electron Tug-of-War
Definition: Electronegativity is the power of an atom (specifically, a nucleus) to attract a pair of electrons towards itself in a covalent bond.
Analogy: Think of two atoms sharing a rope (the electron pair). Electronegativity measures how strongly each atom pulls that rope in a tug-of-war.
Factors Influencing Electronegativity
The stronger the attraction between the nucleus and the outer electrons, the higher the electronegativity. This depends on three factors:
- Nuclear Charge: A larger positive charge pulls the electrons more strongly (higher electronegativity).
- Atomic Radius: Smaller atoms have valence electrons closer to the nucleus, increasing the pull (higher electronegativity).
- Shielding: Inner electron shells block some of the nuclear charge from the valence electrons. More shells mean more shielding and a weaker pull (lower electronegativity).
Trends in the Periodic Table (Quick Review)
Electronegativity generally:
- Increases across a period (due to increasing nuclear charge and constant shielding).
- Decreases down a group (due to increasing atomic radius and increased shielding).
Predicting Bond Type
We use the difference in Pauling electronegativity values ($\Delta\text{EN}$) to predict the bond type (values will be provided in exams where needed):
- \(\Delta\text{EN} < 0.4\): Pure Covalent (Electrons shared equally). E.g., \(\text{H}_2\).
- \(0.4 < \Delta\text{EN} < 1.8\): Polar Covalent (Electrons shared unequally, creating a dipole). E.g., \(\text{HCl}\).
- \(\Delta\text{EN} > 1.8\): Ionic (Complete electron transfer, forming ions). E.g., \(\text{NaCl}\).
Key Takeaway: Electronegativity determines who wins the electron tug-of-war, which in turn defines the type of chemical bond formed.
3.2 Ionic Bonding
Definition: Ionic bonding is the strong electrostatic attraction between oppositely charged ions (positively charged cations and negatively charged anions).
Ionic bonds form when electrons are transferred from a metal atom (which loses electrons to form a cation) to a non-metal atom (which gains electrons to form an anion).
Examples:
- Sodium Chloride (\(\text{NaCl}\)): Na (metal) loses one electron to form \(\text{Na}^+\); Cl (non-metal) gains one electron to form \(\text{Cl}^-\). The strong electrostatic attraction holds these ions together.
- Magnesium Oxide (\(\text{MgO}\)): Mg transfers two electrons to O.
- Calcium Fluoride (\(\text{CaF}_2\)): Ca transfers one electron to each of two F atoms.
3.3 Metallic Bonding
Definition: Metallic bonding is the strong electrostatic attraction between positive metal ions (cations) and a 'sea' of delocalised electrons.
The valence electrons are not held by any single atom but are free to move throughout the structure. This explains why metals are such good conductors of heat and electricity.
Did you know? The malleability and ductility of metals (ability to be hammered into sheets or drawn into wires) is explained by the metallic structure. When layers of ions are slid over each other, the delocalised electrons simply shift to hold the new layers together, preventing repulsion.
Section 2: Covalent Bonding (3.4 & 3.7)
3.4 Covalent Bonding: Sharing is Caring (Sometimes)
Definition: Covalent bonding is the electrostatic attraction between the nuclei of two atoms and a shared pair of electrons.
The shared electrons occupy the space between the two nuclei, effectively holding the atoms together.
Types of Covalent Bonds (Sigma and Pi Bonds)
When atoms bond covalently, their atomic orbitals overlap.
- Sigma (\(\sigma\)) Bonds:
- Formed by the direct (head-on) overlap of orbitals (s-s, s-p, or p-p).
- Electrons are concentrated directly between the two nuclei.
- All single bonds are \(\sigma\) bonds. They allow free rotation around the bond axis.
- Pi (\(\pi\)) Bonds:
- Formed by the sideways overlap of adjacent p orbitals.
- Electrons are concentrated above and below the plane of the \(\sigma\) bond.
- A double bond consists of one \(\sigma\) bond and one \(\pi\) bond (e.g., \(\text{C}_2\text{H}_4\)).
- A triple bond consists of one \(\sigma\) bond and two \(\pi\) bonds (e.g., \(\text{N}_2\), \(\text{HCN}\)).
- \(\pi\) bonds prevent free rotation, which is key to understanding geometrical (cis/trans) isomerism.
Step-by-Step for Hybridisation:
Hybridisation is the mixing of atomic orbitals to form new hybrid orbitals suitable for bonding.
- \(\mathbf{sp^3}\) hybridisation: 1 s-orbital + 3 p-orbitals = 4 identical \(\text{sp}^3\) orbitals. Forms 4 single bonds (tetrahedral shape). Example: \(\text{CH}_4\), \(\text{C}_2\text{H}_6\) (ethane).
- \(\mathbf{sp^2}\) hybridisation: 1 s-orbital + 2 p-orbitals = 3 identical \(\text{sp}^2\) orbitals, plus 1 unhybridised p-orbital. Forms 3 \(\sigma\) bonds and 1 \(\pi\) bond (trigonal planar shape). Example: \(\text{C}_2\text{H}_4\) (ethene).
- \(\mathbf{sp}\) hybridisation: 1 s-orbital + 1 p-orbital = 2 identical \(\text{sp}\) orbitals, plus 2 unhybridised p-orbitals. Forms 2 \(\sigma\) bonds and 2 \(\pi\) bonds (linear shape). Example: \(\text{C}_2\text{H}_2\) (ethyne) or \(\text{N}_2\).
Coordinate (Dative Covalent) Bonding
This is a special type of covalent bond where the shared pair of electrons comes entirely from one atom (the donor).
- Example: The reaction between ammonia (\(\text{NH}_3\)) and hydrogen chloride (\(\text{HCl}\)) gas to form the ammonium ion (\(\text{NH}_4^+\)). The nitrogen atom donates its lone pair to the \(\text{H}^+\) ion.
- Example: The dimeric aluminium chloride molecule (\(\text{Al}_2\text{Cl}_6\)), where chlorine atoms donate lone pairs to the electron-deficient aluminium atoms.
Expanded Octet (Period 3 Elements)
Elements in Period 3 (like Sulfur (S) and Phosphorus (P)) can use their empty d orbitals to accommodate more than 8 electrons in their valence shell. They can "expand their octet".
- Examples: \(\text{PCl}_5\) (Phosphorus pentachloride) and \(\text{SF}_6\) (Sulfur hexafluoride) both have more than 8 electrons around the central atom.
Quick Review: Bond Strength and Length (3.4(a))
- Bond Energy: The energy required to break one mole of a particular covalent bond in the gaseous state. (Always positive, as energy must be put in to break bonds).
- Bond Length: The internuclear distance between two covalently bonded atoms.
Rule: Shorter bonds are stronger bonds, meaning they require more energy to break and are less reactive.
Key Takeaway: Covalent bonds come in two types (\(\sigma\) and \(\pi\)). The presence of \(\pi\) bonds affects geometry and reactivity. Period 3 elements can be electron rebels and expand their octet.
Section 3: Molecular Shapes using VSEPR Theory (3.5)
Don't worry if this seems tricky at first! The VSEPR theory (Valence Shell Electron Pair Repulsion) is a simple way to predict shapes based on how electron pairs push each other apart.
The VSEPR Principle
Electron pairs (both bonding pairs and lone pairs) around a central atom repel each other. They arrange themselves as far apart as possible to minimise this repulsion. Lone pairs repel more strongly than bonding pairs.
Step-by-Step Prediction
- Count the total number of electron pairs around the central atom.
- Determine the arrangement of these pairs (the "electron domain geometry").
- Count the number of lone pairs.
- Use the lone pairs to determine the final molecular shape (the "molecular geometry") and bond angle adjustments.
Common Examples (Syllabus Specific)
| Molecule | Electron Pairs | Lone Pairs | Shape | Bond Angle | Hybridisation |
|---|---|---|---|---|---|
| \(\text{BF}_3\) | 3 | 0 | Trigonal Planar | \(120^\circ\) | \(\text{sp}^2\) |
| \(\text{CO}_2\) | 2 | 0 | Linear | \(180^\circ\) | \(\text{sp}\) |
| \(\text{CH}_4\) | 4 | 0 | Tetrahedral | \(109.5^\circ\) | \(\text{sp}^3\) |
| \(\text{NH}_3\) | 4 | 1 | Pyramidal | \(107^\circ\) | \(\text{sp}^3\) |
| \(\text{H}_2\text{O}\) | 4 | 2 | Non-linear (Bent) | \(104.5^\circ\) | \(\text{sp}^3\) |
| \(\text{SF}_6\) | 6 | 0 | Octahedral | \(90^\circ\) | \(\text{sp}^3\text{d}^2\) |
| \(\text{PF}_5\) | 5 | 0 | Trigonal Bipyramidal | \(120^\circ\) and \(90^\circ\) | \(\text{sp}^3\text{d}\) |
Important Note: Lone pairs reduce the bond angle because they exert a stronger repulsion than bonding pairs (LP-LP > LP-BP > BP-BP).
Key Takeaway: Use VSEPR to explain bond angles and shapes. Remember that lone pairs squash the bond angles!
Section 4: Intermolecular Forces and Polarity (3.6)
Molecules themselves are held together by intramolecular forces (ionic, covalent, metallic - which are very strong). But how do molecules stick to other molecules? That's where Intermolecular Forces (IMFs) come in. These forces are much weaker than chemical bonds.
The hierarchy of strength: Covalent/Ionic/Metallic Bonds > Intermolecular Forces
2. Bond Polarity and Dipole Moments (3.6.2)
When there is a difference in electronegativity, the shared electron pair is pulled closer to the more electronegative atom, creating a bond dipole.
- The atom pulling the electrons gets a partial negative charge (\(\delta^-\)).
- The less electronegative atom gets a partial positive charge (\(\delta^+\)).
The dipole moment of a molecule is the overall measure of its polarity. If the bond dipoles cancel out due to symmetry (like in \(\text{CO}_2\) or \(\text{CH}_4\)), the molecule is non-polar. If they don't cancel (like in \(\text{H}_2\text{O}\) or \(\text{NH}_3\)), the molecule is polar.
3. Van der Waals' Forces (3.6.3)
The term Van der Waals' forces is the generic term for all intermolecular forces between neutral molecular entities, excluding forces resulting from bond formation.
Types of Van der Waals' Forces
A. Instantaneous Dipole–Induced Dipole (id-id) Forces
- Also called London Dispersion Forces.
- These arise because electrons are constantly moving. At any moment, electrons might pile up on one side of a molecule, creating a temporary, instantaneous dipole.
- This temporary dipole then induces a temporary dipole in a neighboring molecule, causing a weak attraction.
- Present in ALL molecules, whether polar or non-polar.
- Strength increases with molecular size (more electrons = greater chance of temporary dipole).
B. Permanent Dipole–Permanent Dipole (pd-pd) Forces
- These forces exist between molecules that are permanently polar (have a net dipole moment, e.g., \(\text{HCl}\)).
- The positive end of one molecule attracts the negative end of a neighbor.
1 & 3(c). Hydrogen Bonding (The Strongest IMF)
Hydrogen bonding is a special, very strong case of permanent dipole-permanent dipole attraction.
It occurs only when hydrogen (H) is covalently bonded to a highly electronegative atom: Nitrogen (N), Oxygen (O), or Fluorine (F). (Remember the mnemonic: F-O-N).
The H atom is stripped of most of its electron density and can then form a very strong electrostatic bridge to a lone pair on a nearby F, O, or N atom.
Anomalous Properties of Water (\(\text{H}_2\text{O}\))
Hydrogen bonding is responsible for water's unusual properties:
- Relatively High Melting and Boiling Points: To melt ice or boil water, you must break the strong hydrogen bonds, which requires significantly more energy than breaking normal VdW forces (as seen in comparable non-polar molecules like \(\text{H}_2\text{S}\)).
- Relatively High Surface Tension: The strong attraction between surface water molecules holds them together tightly.
- Density of Ice (Solid) compared to Liquid Water: In ice, hydrogen bonds lock the molecules into a rigid, open, lattice structure (like a cage). This structure is less dense than the liquid state, which is why ice floats.
Key Takeaway: IMFs are weak attractions *between* molecules. Hydrogen bonding (H-F, H-O, H-N) is the strongest IMF and explains why water is weird!
Section 5: Linking Structure to Physical Properties (4.2)
Chemical structure determines how a substance behaves when heated, dissolved, or subjected to an electric current. We classify solids into four main structural types.
1. Giant Ionic Structures (Lattices)
Structure: A repeating 3D lattice of alternating positive and negative ions, held by strong electrostatic forces.
Examples: \(\text{NaCl}\), \(\text{MgO}\).
Properties:
- Melting/Boiling Point: Very high, because large amounts of energy are needed to break the strong ionic bonds holding the whole lattice together.
- Electrical Conductivity: Non-conductors when solid (ions are fixed). Conductors when molten or aqueous (ions are free to move and carry charge).
- Solubility: Generally soluble in polar solvents like water, which can surround and separate the ions.
2. Simple Molecular Structures
Structure: Discrete, small molecules held together by weak intermolecular forces (VdW or H-bonds).
Examples: \(\text{I}_2\), Buckminsterfullerene (\(\text{C}_{60}\)), ice, \(\text{CO}_2\) (solid).
Properties:
- Melting/Boiling Point: Very low. Only the weak IMFs need to be overcome to melt/boil the substance; the strong covalent bonds within the molecules remain intact.
- Electrical Conductivity: Non-conductors, as there are no free charged particles (ions or electrons) available to move.
- Solubility: Depends on polarity. Non-polar molecules (\(\text{I}_2\)) dissolve in non-polar solvents. Polar molecules (like ethanol) dissolve in polar solvents (water).
3. Giant Molecular (Covalent) Structures
Structure: A vast 3D network of atoms held together entirely by strong covalent bonds. There are no individual molecules.
Examples: Diamond (\(\text{C}\)), Graphite (\(\text{C}\)), Silicon(IV) Oxide (\(\text{SiO}_2\)).
Properties:
- Melting/Boiling Point: Extremely high, as all millions of strong covalent bonds must be broken simultaneously.
- Hardness: Extremely hard (e.g., diamond, \(\text{SiO}_2\)).
- Electrical Conductivity: Generally non-conductors (e.g., diamond, \(\text{SiO}_2\)).
The special case of Graphite: Graphite is an exception! It consists of hexagonal layers of carbon atoms held by strong covalent bonds, but the layers themselves are held only by weak London forces. This makes it:
- Soft and a good lubricant (layers slide easily).
- A good electrical conductor (each carbon atom uses only three valence electrons for bonding, leaving the fourth electron delocalised between the layers, free to move).
4. Giant Metallic Structures
Structure: A lattice of positive ions in a sea of delocalised electrons (as described in 3.3).
Example: Copper (\(\text{Cu}\)), Iron (\(\text{Fe}\)).
Properties:
- Melting/Boiling Point: Generally high (strong metallic bonds).
- Electrical Conductivity: Excellent conductors (free moving delocalised electrons).
- Malleability and Ductility: High.
Key Takeaway: Strong structures (Giant Ionic, Giant Metallic, Giant Covalent) have high melting points. Simple Molecular structures have low melting points because only weak IMFs are broken.
Section 6: Visualising Bonds: Dot-and-Cross Diagrams (3.7)
Dot-and-cross diagrams show the arrangement of electrons in the outer shells of atoms or ions. Typically, electrons from one atom are shown as dots (•) and electrons from the second atom as crosses (x).
Drawing Ionic Compounds
Show the transfer of electrons. The resultant ions must be enclosed in square brackets showing the charges, and the electrons of the anion should reflect both its original electrons and the transferred electrons.
- Example: \(\text{Na}^+\) will show 8 outer electrons (originally 1 valence electron, now 8 in the second shell), with a +1 charge outside the bracket.
Drawing Covalent Compounds
Show the sharing of electrons so that most atoms achieve a stable noble gas configuration (usually 8 electrons, the octet rule).
- Show only the outer shell electrons.
- The shared electrons should be placed in the overlap area between the two atoms.
- For complex examples like \(\text{H}_2\text{O}\) or \(\text{NH}_3\), ensure that lone pairs are clearly shown alongside bonding pairs, as they are crucial for VSEPR theory.
Remember to apply these rules to tricky cases:
- Expanded Octets: For \(\text{PCl}_5\) or \(\text{SF}_6\), the central atom (P or S) must show more than 8 electrons in its outer shell (e.g., 10 for P, 12 for S).
- Odd Number of Electrons: Some species (called free radicals) may have an odd number of electrons and cannot achieve a full octet. They possess one or more unpaired electrons.
- Coordinate Bonds: Use the dot and cross notation to show that both electrons in a shared pair came from the same initial atom (e.g., the lone pair on N in \(\text{NH}_4^+\)).