Atomic Structure and Bonding
Subject: Chemistry
Topic: 1
Cambridge Code: 0620 / 0971 / 5070
Atomic Structure
Atom - Smallest unit of an element
Subatomic Particles
| Particle | Location | Charge | Mass (amu) |
|---|---|---|---|
| Proton | Nucleus | +1 | 1 |
| Neutron | Nucleus | 0 | 1 |
| Electron | Electron shells | -1 | 0 |
Atomic Number and Mass Number
Atomic number (Z) = Number of protons = Number of electrons
Mass number (A) = Number of protons + neutrons
Isotopes
Atoms of same element with different number of neutrons
- Same atomic number, different mass number
- Similar chemical properties, different physical properties
- Different radioactivity
Example: ¹²C and ¹⁴C (both carbon, different neutrons)
Electron Configuration
Electron Configuration - Arrangement of electrons in shells
Shell Structure
- First shell (n=1): max 2 electrons
- Second shell (n=2): max 8 electrons
- Third shell (n=3): max 18 electrons
Filling rule: Fill inner shells before outer shells
Examples
- Hydrogen (1 electron): 1
- Carbon (6 electrons): 2, 4
- Oxygen (8 electrons): 2, 6
- Sodium (11 electrons): 2, 8, 1
- Chlorine (17 electrons): 2, 8, 7
Valence Electrons
Electrons in outermost shell - determine bonding
Ionic Bonding
Ionic Bond - Electrostatic attraction between oppositely charged ions
Formation
- Metal atoms LOSE electrons → Cations (positive)
- Non-metal atoms GAIN electrons → Anions (negative)
- Electrostatic attraction holds them together
Examples
NaCl (Sodium Chloride):
- Na loses 1 electron → Na⁺
- Cl gains 1 electron → Cl⁻
- Held by electrostatic force
MgO (Magnesium Oxide):
- Mg loses 2 electrons → Mg²⁺
- O gains 2 electrons → O²⁻
Properties
- High melting/boiling points
- Conduct electricity when molten/dissolved
- Brittle (breaks on impact)
- Soluble in polar solvents (e.g., water)
- Crystalline structure
Covalent Bonding
Covalent Bond - Shared pair of electrons
Types
Single Bond: 1 shared pair (e.g., H-H) Double Bond: 2 shared pairs (e.g., O=O) Triple Bond: 3 shared pairs (e.g., N≡N)
Polar vs. Non-Polar
Non-polar covalent:
- Electrons shared equally
- Similar electronegativity
- E.g., Cl-Cl, C-H
Polar covalent:
- Electrons shared unequally
- Different electronegativity
- Slight charge separation (δ+ and δ-)
- E.g., H-Cl, H-O
Molecular vs. Giant Covalent
Molecular covalent:
- Atoms bonded by covalent bonds
- Weak intermolecular forces
- Low melting point
- E.g., CO₂, H₂O, ethanol
Giant covalent:
- All atoms covalently bonded throughout
- Very high melting point
- Insoluble
- Hard or soft but brittle
- E.g., Diamond, Silicon dioxide
Metallic Bonding
Metallic Bond - Electrostatic attraction between cations and delocalized electrons
Structure
- Metal atoms lose valence electrons
- Form cation lattice
- Electrons move freely ("electron sea")
Properties
- High melting/boiling point
- Conduct electricity (solid and molten)
- Malleable and ductile
- Shiny luster
- Generally insoluble (except ionic salts)
Electronegativity
Electronegativity - Ability to attract electrons in a bond
Trends
- Increases ACROSS a period
- Decreases DOWN a group
- Fluorine is most electronegative
- Electronegativity difference determines bond type
Bond Classification
- Difference < 0.4: Non-polar covalent
- Difference 0.4-1.7: Polar covalent
- Difference > 1.7: Ionic (approximately)
Key Points
- Protons and neutrons in nucleus, electrons in shells
- Isotopes: same Z, different A
- Ionic: electron transfer, electrostatic attraction
- Covalent: electron sharing
- Metallic: delocalized electrons
- Bonding type depends on electronegativity
Practice Questions
- What is the atomic number and mass number of ³⁵Cl?
- Compare ionic and covalent bonding
- Explain why NaCl conducts electricity when molten
- Draw electron configuration for sodium
- Classify C-C, H-Cl, and Na-Cl bonds
Revision Tips
- Learn electron configuration rules
- Understand electronegativity trends
- Know properties of each bonding type
- Practice calculating subatomic particles
- Draw bonding diagrams