Atomic Structure
Concise revision notes aligned to the O Level syllabus.
Subatomic Particles
Every atom is made up of three types of subatomic particle. Protons and neutrons sit in the dense central nucleus; electrons occupy shells (energy levels) around the nucleus.
| Particle | Relative charge | Relative mass | Location |
|---|---|---|---|
| Proton | +1 | 1 | Nucleus |
| Neutron | 0 | 1 | Nucleus |
| Electron | −1 | Negligible (≈ 1/1836) | Shells around nucleus |
Almost all the mass of an atom is concentrated in the nucleus. The atom is mostly empty space — electrons are tiny compared to the nucleus and have negligible mass.
Atomic Number and Mass Number
Two key numbers define every atom:
- Atomic number (Z) — number of protons in the nucleus. This identifies the element: all carbon atoms have Z = 6.
- Mass number (A) — total number of protons + neutrons (nucleons) in the nucleus.
Number of neutrons = Mass number − Atomic number
In a neutral atom, the number of electrons always equals the number of protons.
| Example | Atomic number | Mass number | Protons | Neutrons | Electrons |
|---|---|---|---|---|---|
| Carbon-12 | 6 | 12 | 6 | 6 | 6 |
| Carbon-14 | 6 | 14 | 6 | 8 | 6 |
| Sodium-23 | 11 | 23 | 11 | 12 | 11 |
| Chlorine-35 | 17 | 35 | 17 | 18 | 17 |
The atomic number never changes for a given element — it defines what element it is. The mass number can vary (isotopes). Always subtract atomic number from mass number to get the neutron count.
Isotopes
Isotopes are atoms of the same element with the same atomic number but different mass numbers — i.e. different numbers of neutrons.
- They have identical chemical properties because they have the same number of outer-shell electrons (which determine chemical behaviour).
- They differ in physical properties such as density and rate of diffusion.
Relative atomic mass (Aᵣ) is the weighted average mass of all naturally occurring isotopes of an element, relative to carbon-12.
Example — chlorine exists as Cl-35 (75%) and Cl-37 (25%):
Aᵣ = (35 × 75 + 37 × 25) / 100 = 35.5
Isotopes have the same chemical properties because chemical reactions depend on the electron arrangement, not the number of neutrons. The neutron count only affects mass-related physical properties.
Electron Configuration
Electrons fill shells from the innermost outward. For O Level you need to know:
| Shell | Maximum electrons |
|---|---|
| 1st (innermost) | 2 |
| 2nd | 8 |
| 3rd | 8 (for elements up to Z = 20) |
The configuration is written as a sequence of numbers separated by commas, e.g. 2,8,5 for phosphorus (Z = 15).
| Element | Z | Configuration |
|---|---|---|
| Hydrogen | 1 | 2,1 → actually 1 |
| Carbon | 6 | 2,4 |
| Oxygen | 8 | 2,6 |
| Sodium | 11 | 2,8,1 |
| Chlorine | 17 | 2,8,7 |
| Argon | 18 | 2,8,8 |
| Potassium | 19 | 2,8,8,1 |
| Calcium | 20 | 2,8,8,2 |
The number of outer-shell electrons equals the group number (for Groups 1–7). Elements in Group 0 (noble gases) have full outer shells.
The period number tells you how many shells an atom has; the group number tells you how many electrons are in the outermost shell. These two facts let you reconstruct any configuration for Z ≤ 20.
Ions
An ion is a charged atom (or group of atoms) formed when electrons are gained or lost.
- Cation (positive ion): formed when an atom loses electrons.
Na → Na⁺ + e⁻— sodium loses its one outer electron; configuration goes from 2,8,1 to 2,8. - Anion (negative ion): formed when an atom gains electrons.
Cl + e⁻ → Cl⁻— chlorine gains one electron; configuration goes from 2,8,7 to 2,8,8.
Ions are stable because they achieve a full outer shell — the same electron configuration as a noble gas.
| Ion | Charge | Electrons lost/gained | Noble gas configuration |
|---|---|---|---|
| Na⁺ | +1 | Lost 1 | 2,8 (Ne) |
| Mg²⁺ | +2 | Lost 2 | 2,8 (Ne) |
| Al³⁺ | +3 | Lost 3 | 2,8 (Ne) |
| Cl⁻ | −1 | Gained 1 | 2,8,8 (Ar) |
| O²⁻ | −2 | Gained 2 | 2,8 (Ne) |
| S²⁻ | −2 | Gained 2 | 2,8,8 (Ar) |
Metals form positive ions (they lose electrons); non-metals form negative ions (they gain electrons). The size of the charge equals the number of electrons lost or gained.
Ionic Bonding
Ionic bonding occurs between a metal and a non-metal. Electrons are transferred from the metal to the non-metal, forming oppositely charged ions. The strong electrostatic attraction between these ions holds the compound together.
Example — sodium chloride (NaCl):
- Na (2,8,1) loses 1 electron → Na⁺ (2,8)
- Cl (2,8,7) gains 1 electron → Cl⁻ (2,8,8)
- Na⁺ and Cl⁻ attract each other and arrange into a giant ionic lattice.
Properties of ionic compounds:
| Property | Explanation |
|---|---|
| High melting/boiling points | Strong electrostatic forces in the lattice require a lot of energy to break |
| Conduct when molten or dissolved | Ions are free to move and carry charge |
| Do not conduct when solid | Ions are fixed in position in the lattice |
| Often soluble in water | Water molecules can pull apart the ions |
Ionic compounds conduct electricity when molten or in solution because the ions are free to move. In the solid state the ions are fixed — no conduction is possible.
Covalent Bonding
Covalent bonding occurs between two non-metals. Atoms share pairs of electrons, each pair forming one covalent bond. Both atoms achieve a full outer shell this way.
| Bond type | Shared pairs | Example |
|---|---|---|
| Single bond | 1 pair | H–H (H₂), H–Cl (HCl) |
| Double bond | 2 pairs | O=O (O₂), O=C=O (CO₂) |
| Triple bond | 3 pairs | N≡N (N₂) |
Typical bonding numbers:
- H forms 1 bond
- O forms 2 bonds
- N forms 3 bonds
- C forms 4 bonds
Simple molecular substances (H₂O, CO₂, CH₄): the covalent bonds within each molecule are strong, but the intermolecular forces between molecules are weak → low melting and boiling points, do not conduct electricity.
Giant covalent structures (diamond, graphite, SiO₂): covalent bonds extend throughout the entire structure → very high melting points.
Do not confuse the strength of covalent bonds with the melting point of a simple covalent substance. Simple molecules melt easily because the weak forces between molecules (not the bonds within them) are what you need to overcome.
Metallic Bonding
Metallic bonding exists in metals and alloys. Metal atoms release their outer electrons into a sea of delocalised electrons that moves freely throughout the structure. The resulting positive metal ions are held in a regular lattice by electrostatic attraction to this electron sea.
Properties of metals explained by metallic bonding:
| Property | Explanation |
|---|---|
| Good electrical conductivity | Delocalised electrons move freely and carry charge |
| Good thermal conductivity | Delocalised electrons transfer energy rapidly |
| High melting/boiling points | Strong attraction between positive ions and electron sea |
| Malleability and ductility | Layers of ions can slide over each other without breaking the bonding |
| Lustrous (shiny) | Delocalised electrons reflect light |
Metals conduct because electrons (not ions) are free to move. This contrasts with ionic solutions, where ions carry the charge. Distinguishing which charge carrier is responsible is a common exam question.
Comparing the Three Bond Types
| Feature | Ionic | Covalent | Metallic |
|---|---|---|---|
| Electron behaviour | Transferred | Shared | Delocalised (sea) |
| Bonded species | Metal + non-metal | Non-metals | Metal atoms |
| Melting point | High | Low (simple) / Very high (giant) | High |
| Electrical conductivity | Only when molten/dissolved | Generally no | Yes (solid and liquid) |
| Example | NaCl, MgO, CaCl₂ | H₂O, CO₂, CH₄, diamond | Fe, Cu, Al |
The key differentiator: ionic compounds conduct only when free ions can move (molten/dissolved); metals conduct in all states because electrons are always free; simple covalent molecules never conduct because there are no free charged particles.
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