Quick Simpler Start
You learn how atoms join, why molecules have different shapes, and how these bonds affect melting point, boiling point, and polarity.
When you revise this chapter, ask yourself: "Can I explain this to a friend in one easy paragraph?" If yes, you are in a good place.
Interactive Simulator
This small tool gives you a quick visual check before you move deeper into the chapter notes.
Bond Explorer
Choose the bond type and the tool explains the idea in simple student language.
Chapter 2: Chemical Bonding and Molecular Geometry
1. Objectives: What you need to know
- Explain why atoms form compounds using noble gas stability and the octet rule.
- Distinguish between ionic, covalent, polar covalent, and non-polar covalent bonding.
- Use electronegativity differences to predict bond type and polarity.
- Describe van der Waals, dipole-dipole, and hydrogen bonding.
- Explain how intermolecular forces affect boiling point and solubility.
- Use VSEPR theory to predict linear, planar, tetrahedral, pyramidal, and V-shaped molecules.
2. Big idea: Why do atoms bond?
Atoms bond because they become more stable when their outer shell arrangement becomes similar to a noble gas. Most main-group atoms aim for eight electrons in their outer shell. This is called the octet rule.
3. Exam-ready definitions
| Term | Definition | Exam clue |
|---|---|---|
| Ionic bond | Electrostatic attraction between oppositely charged ions. | Usually metal + non-metal. |
| Covalent bond | A bond formed by sharing a pair of electrons between atoms. | Usually non-metal + non-metal. |
| Electronegativity | The relative attraction of an atom for shared electrons in a chemical bond. | Higher values pull bonding electrons more strongly. |
| Polar bond | A covalent bond where electrons are shared unequally. | Creates partial charges: δ+ and δ−. |
| Intermolecular forces | Forces of attraction between molecules. | Affect boiling point, melting point, and solubility. |
| VSEPR theory | Electron pairs repel and arrange themselves as far apart as possible. | Used to predict molecular shapes. |
4. Ionic and covalent bonding
Ionic Bonding
Electrons are transferred. One atom loses electrons to become a positive ion, while another gains electrons to become a negative ion.
Example: NaCl
Covalent Bonding
Electrons are shared. Atoms share one or more pairs of electrons to achieve a stable outer shell.
Example: H₂O, CH₄, O₂, N₂
5. Polar and non-polar covalent bonds
A covalent bond is non-polar when the electrons are shared equally or almost equally. A covalent bond is polar when one atom attracts the shared pair more strongly.
6. Electronegativity
Electronegativity increases across a period from left to right and generally decreases down a group. Fluorine has the highest electronegativity.
| Electronegativity difference | Likely bond type | Example |
|---|---|---|
| Very small / zero | Non-polar covalent | Cl₂, O₂ |
| Moderate | Polar covalent | HCl, H₂O |
| Large | Ionic | NaCl, MgO |
7. Intermolecular forces
Intermolecular forces are not the same as covalent bonds. Covalent bonds hold atoms together inside molecules. Intermolecular forces attract separate molecules to each other.
Van der Waals forces
Weak attractions caused by temporary dipoles. Present between all molecules.
Dipole-dipole interactions
Attractions between permanent dipoles in polar molecules.
Hydrogen bonding
Strong intermolecular force when H is bonded to N, O, or F.
8. Molecular shapes: VSEPR theory
VSEPR means Valence Shell Electron Pair Repulsion. Electron pairs repel each other and arrange themselves as far apart as possible.
9. Formula / decision toolbox
Small difference → covalent | Moderate difference → polar covalent | Large difference → ionic
10. Worked examples
Step 1: Carbon is the central atom.
Step 2: Carbon has 4 bonding pairs and no lone pairs.
Step 3: Four electron pairs spread out as far as possible.
Answer: Tetrahedral, bond angle about 109.5°.
Answer: Water molecules form hydrogen bonds because hydrogen is bonded to oxygen. Hydrogen bonding is stronger than the intermolecular forces in H₂S, so more energy is needed to separate water molecules. Therefore water has a higher boiling point.
11. Common student mistakes
2. Confusing intermolecular forces with covalent bonds.
3. Forgetting that lone pairs change molecular shape and reduce bond angles.
4. Calling any molecule with H a hydrogen-bonded molecule. Hydrogen must be bonded to N, O, or F.
12. Examiner traps
13. Practical / application skills
• Ionic compounds usually have high melting points because strong ionic bonds extend through a crystal lattice.
• Simple covalent substances often have low boiling points because only weak intermolecular forces are overcome during boiling.
• Hydrogen bonding increases boiling points and helps explain the unusual properties of water.
• Polar substances often dissolve in polar solvents such as water.
14. Exam-style questions
Q1. Explain, using electron transfer, how sodium chloride forms. [5 marks]
Q2. Draw a dot-and-cross diagram for a molecule of water and state its shape. [5 marks]
Q3. Explain why ammonia, NH₃, is pyramidal rather than planar. [4 marks]
Q4. Hydrogen fluoride has a much higher boiling point than hydrogen chloride. Explain why. [4 marks]
15. MCQs with explanations
| Question | Answer and explanation |
|---|---|
| 1. Which bond is formed by sharing electrons? A Ionic B Covalent C Metallic D Hydrogen | B. Covalent bonding involves sharing electron pairs. Ionic bonding involves electron transfer. |
| 2. Which molecule has a tetrahedral shape? A BeCl₂ B BF₃ C CH₄ D H₂O | C. CH₄ has four bonding pairs around carbon and no lone pairs, giving a tetrahedral shape. |
| 3. Which intermolecular force is strongest in water? A Van der Waals B Hydrogen bonding C Ionic bonding D Metallic bonding | B. Water forms hydrogen bonds because hydrogen is bonded to oxygen. |
| 4. What happens to electronegativity across a period? A Increases B Decreases C Stays the same D Becomes zero | A. Electronegativity generally increases across a period from left to right. |
16. Higher Level Challenge
Question: Explain why NH₃ has a smaller bond angle than CH₄, even though both have four regions of electron density around the central atom.
Answer idea: CH₄ has four bonding pairs around carbon, giving 109.5°. NH₃ has three bonding pairs and one lone pair around nitrogen. The lone pair repels bonding pairs more strongly, pushing the N-H bonds closer together and reducing the bond angle to about 107°.
17. Last-minute revision sheet
✓ Ionic = electron transfer + attraction between ions.
✓ Covalent = sharing electron pairs.
✓ Electronegativity tells you how strongly an atom attracts bonding electrons.
✓ Polar bond = unequal sharing.
✓ Hydrogen bonding requires H bonded to N, O, or F.
✓ VSEPR: electron pairs repel and spread out.
✓ Lone pairs repel more strongly than bonding pairs.
✓ CH₄ tetrahedral, NH₃ pyramidal, H₂O V-shaped, BF₃ planar, BeCl₂ linear.
18. Self-assessment checklist
- I can explain the octet rule and noble gas stability.
- I can distinguish ionic and covalent bonding.
- I can explain polar and non-polar covalent bonds.
- I can describe electronegativity trends.
- I can identify van der Waals, dipole-dipole, and hydrogen bonding.
- I can predict simple molecular shapes using VSEPR theory.
- I can explain how lone pairs affect bond angles.
19. Mark schemes
Na has one outer electron [1]
Cl has seven outer electrons [1]
Na transfers one electron to Cl [1]
Na becomes Na+ and Cl becomes Cl- [1]
Oppositely charged ions attract strongly / ionic bond forms [1]
Q2 Water [5]
Oxygen central atom [1]
Two O-H covalent bonds shown [1]
Two lone pairs on oxygen shown [1]
Shape: V-shaped / bent [1]
Lone pairs repel bonding pairs more strongly [1]
Q3 Ammonia [4]
N has three bonding pairs and one lone pair [1]
Electron pairs repel [1]
Lone pair repels bonding pairs more strongly [1]
Shape is pyramidal / bond angle about 107° [1]
Q4 HF vs HCl [4]
HF has hydrogen bonding [1]
H is bonded to very electronegative F [1]
Hydrogen bonds are stronger than forces between HCl molecules [1]
More energy required to separate HF molecules, so boiling point is higher [1]
20. Higher Level Syllabus Patch: Missing Chapter 2 Content
This section adds the remaining Higher Level Chapter 2 points so the chapter matches the syllabus more closely: sigma and pi bonding, intramolecular versus intermolecular forces, oxidation numbers, transition-metal variable oxidation states, and naming compounds from oxidation states.
20.1 Sigma (σ) and pi (π) bonds
At Higher Level, students should know that covalent bonds can be described in terms of sigma and pi bonding.
Sigma bond, σ
A sigma bond is formed by direct overlap of orbitals along the line joining the nuclei. Every single covalent bond is a sigma bond.
Pi bond, π
A pi bond is formed by sideways overlap of orbitals above and below the line joining the nuclei. Pi bonds occur in addition to a sigma bond in double and triple bonds.
| Bond Type | Composition | Example |
|---|---|---|
| Single bond | 1 σ bond | H–H, C–H |
| Double bond | 1 σ bond + 1 π bond | O=O, C=C |
| Triple bond | 1 σ bond + 2 π bonds | N≡N, C≡C |
20.2 Intramolecular vs intermolecular forces
This distinction is a classic Higher Level trap.
| Type of Force | Acts Where? | Examples | Main Effect |
|---|---|---|---|
| Intramolecular | Inside a molecule or crystal | Covalent, ionic, metallic bonding | Determines chemical structure and bond strength |
| Intermolecular | Between separate molecules | Van der Waals, dipole-dipole, hydrogen bonding | Strongly affects boiling point, melting point and solubility |
20.3 Oxidation numbers (oxidation states)
The oxidation number of an element is the apparent charge that atom would have if the bonding electrons were assigned according to simple rules.
Reduction = decrease in oxidation number
Rules for assigning oxidation numbers
- An element in its uncombined form has oxidation number 0. Example: Na, O2, Cl2.
- A simple ion has oxidation number equal to its charge. Example: Na+ = +1, O2− = −2.
- Group 1 metals are usually +1; Group 2 metals are usually +2.
- Fluorine is always −1 in compounds.
- Oxygen is usually −2 (except peroxides and OF2).
- Hydrogen is usually +1 with non-metals and −1 in metal hydrides.
- The total oxidation numbers in a neutral compound add to 0.
- The total oxidation numbers in a polyatomic ion add to the charge on the ion.
Worked examples
H = +1, O = −2
2(+1) + S + 4(−2) = 0
2 + S − 8 = 0
S = +6
O = −2
N + 3(−2) = −1
N − 6 = −1
N = +5
20.4 Transition metals and variable oxidation states
Many transition metals show variable oxidation states. This is one reason transition-metal chemistry is rich and colourful.
| Ion / Compound | Oxidation Number of Metal | Name |
|---|---|---|
| FeCl2 | Fe = +2 | iron(II) chloride |
| FeCl3 | Fe = +3 | iron(III) chloride |
| Cu2O | Cu = +1 | copper(I) oxide |
| CuO | Cu = +2 | copper(II) oxide |
| MnO2 | Mn = +4 | manganese(IV) oxide |
20.5 Redox using oxidation numbers
Redox can be interpreted using oxidation-number changes as well as electron transfer.
Fe goes from +2 to +3, so iron is oxidised.
Cl goes from 0 to −1, so chlorine is reduced.
20.6 Symmetry and overall polarity
At Higher Level, students should be able to separate bond polarity from molecular polarity.
| Molecule | Bonds polar? | Shape / symmetry | Overall polarity |
|---|---|---|---|
| CO2 | Yes | Linear and symmetrical | Non-polar |
| BF3 | Yes | Trigonal planar and symmetrical | Non-polar |
| NH3 | Yes | Pyramidal, asymmetrical | Polar |
| H2O | Yes | V-shaped, asymmetrical | Polar |
20.7 Quick Higher Level consolidation
- I can explain the difference between sigma and pi bonds.
- I can distinguish intramolecular and intermolecular forces.
- I can assign oxidation numbers correctly.
- I can name transition-metal compounds using Roman numerals.
- I can explain redox as changes in oxidation number.
- I can use molecular symmetry to decide whether a molecule is polar overall.
21. Chapter 2 Syllabus Completion Check
This completed version now covers the main Higher Level Chapter 2 areas: ionic and covalent bonding, electronegativity, polar and non-polar bonding, sigma and pi bonding, intermolecular forces, VSEPR shapes, symmetry and molecular polarity, oxidation numbers, and transition-metal oxidation-state naming. It is designed as the patched master file for Chapter 2.