Key concepts in chemistry
Atomic structure
Every atom consists of a tiny, dense nucleus containing protons and (usually) neutrons, surrounded by shells of electrons. Key particles:
| Particle | Relative mass | Relative charge | Location |
|---|---|---|---|
| Proton | 1 | +1 | Nucleus |
| Neutron | 1 | 0 | Nucleus |
| Electron | 1/1836 (≈0) | −1 | Shells around nucleus |
Atomic number (Z): the number of protons. This uniquely defines each element. Mass number A: protons + neutrons. Number of neutrons = A − Z.
Electronic configuration
Electrons fill shells in order: shell 1 holds up to 2, shell 2 holds up to 8, shell 3 holds up to 8 (at GCSE). Example: sodium (Z=11) → 2, 8, 1.
The number of outer-shell electrons determines an element's chemical properties and its group in the periodic table.
Isotopes
Isotopes are atoms of the same element with the same number of protons but different numbers of neutrons (different mass numbers). Example: carbon-12 (⁶₁₂C) and carbon-14 (⁶₁₄C) are both carbon but ¹²C has 6 neutrons and ¹⁴C has 8 neutrons.
Relative atomic mass (Aᵣ): the weighted average mass of all naturally occurring isotopes relative to ¹/₁₂ the mass of a carbon-12 atom. Example: chlorine has Aᵣ ≈ 35.5 because it is ~75% ³⁵Cl and ~25% ³⁷Cl.
The periodic table
The periodic table arranges elements in order of increasing atomic number. Key features:
- Periods (rows): elements in the same period have the same number of electron shells. Period 2 has 8 elements (Li → Ne).
- Groups (columns): elements in the same group have the same number of outer electrons → similar chemistry. Group 1 = 1 outer electron; Group 7 = 7 outer electrons; Group 0 = full outer shell.
- Metals vs non-metals: the dividing staircase runs from B→At. Metals are on the left; non-metals on the right. Metalloids (semi-metals) straddle the line.
- Transition metals: the central d-block (Groups 3–12). They form coloured compounds, have variable oxidation states, and act as catalysts.
Edexcel exam tip: Mendeleev
Mendeleev arranged elements by atomic mass (1869), leaving gaps for undiscovered elements. Modern table uses atomic number — this resolved anomalies (e.g. Ar/K) and is confirmed by Moseley's X-ray work.
Types of substance and bonding
| Bond type | Between | Result | Typical properties |
|---|---|---|---|
| Ionic | Metal + non-metal | Giant ionic lattice | High m.p., conducts when molten/dissolved, brittle |
| Covalent (simple) | Non-metals | Discrete molecules | Low m.p., poor conductors |
| Covalent (giant) | Non-metals (e.g. diamond) | Giant lattice | Very high m.p., poor conductors (except graphite) |
| Metallic | Metal atoms | Giant metallic lattice | Good conductor, malleable, ductile |
Moles and relative formula mass (Mᵣ)
Relative formula mass (Mᵣ): sum of Aᵣ values for all atoms in the formula. H₂O: (2 × 1) + 16 = 18. NaCl: 23 + 35.5 = 58.5. CaCO₃: 40 + 12 + (3 × 16) = 100.
Mole: 1 mole = 6.02 × 10²³ particles (Avogadro constant). n = mass ÷ Mᵣ (moles = mass in grams divided by relative formula mass).
This formula is the basis of all quantitative chemistry (yield calculations, titrations, concentrations).
⚠Common mistakes— Common mistakes (Edexcel examiner traps)
- Confusing atomic number and mass number: atomic number = protons; mass number = protons + neutrons.
- Isotope definition: isotopes have the same protons (same element), different neutrons — NOT same mass.
- Group numbering: Group 1 = alkali metals (NOT hydrogen, which has unique chemistry).
- Mole calculation direction: n = m/Mᵣ, so mass = n × Mᵣ. Don't invert.
- Relative atomic mass for isotopes: Aᵣ is the weighted mean (accounting for abundance), not a simple average.
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