🧪 AS 1: Basic Concepts in Physical & Inorganic Chemistry
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CCEA GCE Chemistry (1110) · Unit AS 1: Basic Concepts in Physical and Inorganic Chemistry
Mini-Lesson
AS 1: Basic Concepts in Physical & Inorganic Chemistry
Unit AS 1 lays the foundations: atomic structure, the mole, formulae and equations, bonding, shapes of molecules, intermolecular forces and oxidation numbers.
Everything you meet later — energetics, kinetics, equilibrium and organic mechanisms — is built on these ideas, so the calculations here need to become automatic.
Work through each screen, answer the questions as you go (several are full A-level calculations) and collect ⭐ stars. Press Start when you're ready.
Atomic structure
Isotopes, mass spectrometry and Ar
An atom is defined by its proton number. Isotopes have the same number of protons but different numbers of neutrons, so they have identical chemical behaviour (same electron arrangement) but different masses.
Ar = Σ(isotopic mass × abundance) ÷ Σ(abundances)the weighted mean mass, on the scale where ¹²C = exactly 12
A mass spectrometer ionises the sample, accelerates the ions, separates them by mass-to-charge ratio (m/z) and detects them. Each peak is one isotope; the peak heights give the abundances.
Calculate
Your turn — calculation 1
1Magnesium has three isotopes: 78.6% of mass 24, 10.1% of mass 25 and 11.3% of mass 26. Calculate the relative atomic mass of magnesium, to 1 decimal place.
Electrons fill orbitals in order of energy: 1s 2s 2p 3s 3p 4s 3d 4p. Note that 4s fills before 3d, but the 4s electrons are removed first when the atom is ionised.
The first ionisation energy is the energy needed to remove one electron from each atom in one mole of gaseous atoms:
X(g) → X⁺(g) + e⁻note the state symbols — gaseous, and one mole. Both are needed for the definition mark.
It depends on nuclear charge, atomic radius and shielding by inner shells.
The two anomalies across a period: Al is lower than Mg (Al's outer electron is in a higher-energy, more shielded 3p orbital), and S is lower than P (in S, two electrons are paired in one 3p orbital and repel each other).
Quick check
Think it through
?The first ionisation energies of consecutive Period 3 elements are: Na 496, Mg 738, Al 578, Si 789 kJ mol⁻¹. Why is aluminium lower than magnesium?
Amount of substance
The mole and the ideal gas equation
n = m ÷ Mmoles = mass (g) ÷ molar mass (g mol⁻¹)
n = c × Vmoles = concentration (mol dm⁻³) × volume (dm³)
pV = nRTp in Pa · V in m³ · T in K · R = 8.314 J K⁻¹ mol⁻¹
Empirical formula = the simplest whole-number ratio of atoms. Method: divide the mass (or %) of each element by its Ar, then divide all the answers by the smallest.
Molecular formula = a whole-number multiple of the empirical formula, found by comparing the empirical formula mass with the Mr.
Unit discipline for pV = nRT: kPa → ×1000 (Pa); cm³ → ÷10⁶ (m³); dm³ → ÷1000 (m³); °C → +273 (K). Most lost marks in this topic are unit conversions, not chemistry.
Calculate
Your turn — calculation 2
2A sample of an iron oxide contains 2.80 g of iron and 1.20 g of oxygen. (Ar: Fe = 55.8, O = 16.0). How many oxygen atoms are in the empirical formula?
O atoms
Hint: n(Fe) = 2.80 ÷ 55.8; n(O) = 1.20 ÷ 16.0. Divide both by the smaller value.
30.400 g of a gas occupies 250 cm³ at 100 kPa and 300 K. Calculate its molar mass in g mol⁻¹, to 3 significant figures. (R = 8.314 J K⁻¹ mol⁻¹)
g mol⁻¹
Hint: Convert: p = 100 000 Pa, V = 250 × 10⁻⁶ m³. n = pV ÷ RT, then M = m ÷ n.
Method
n = (100 000 × 2.50 × 10⁻⁴) ÷ (8.314 × 300) = 25.0 ÷ 2494.2 = 0.01002 mol. M = 0.400 ÷ 0.01002 = 39.9 g mol⁻¹ — the gas is argon.
Calculate
Your turn — calculation 4
45.30 g of anhydrous sodium carbonate, Na₂CO₃ (M = 106.0), is dissolved and made up to 250 cm³ of solution. Calculate the concentration in mol dm⁻³.
mol dm⁻³
Hint: n = 5.30 ÷ 106.0. Then c = n ÷ 0.250 dm³.
Method
n = 5.30 ÷ 106.0 = 0.0500 mol. c = 0.0500 ÷ 0.250 = 0.200 mol dm⁻³.
Calculate
Your turn — calculation 5
5Calculate the percentage by mass of nitrogen in ammonium nitrate, NH₄NO₃. (Ar: H = 1.0, N = 14.0, O = 16.0)
%
Hint: M(NH₄NO₃) = 14.0 + 4.0 + 14.0 + 48.0 = 80.0. There are 2 N atoms → 28.0. Then (28.0 ÷ 80.0) × 100.
Method
M = 80.0 g mol⁻¹; mass of N = 2 × 14.0 = 28.0. % N = (28.0 ÷ 80.0) × 100 = 35.0% — the reason ammonium nitrate makes such an efficient fertiliser.
Bonding
Ionic, covalent and metallic bonding
Ionic — electrons are transferred from a metal to a non-metal; the giant lattice is held by the electrostatic attraction between oppositely charged ions. High melting point; conducts only when molten or aqueous, when the ions are free to move.
Covalent — a shared pair of electrons. Simple molecular substances (CO₂, Cl₂) have weak intermolecular forces between the molecules, so they melt at low temperatures — the strong covalent bonds inside the molecules are not broken on melting. Giant covalent substances (SiO₂, diamond) melt only when covalent bonds break, so they are extremely high-melting.
Metallic — a lattice of positive ions in a sea of delocalised electrons. Good conductors; malleable, because layers can slide without breaking the bonding.
Dative covalent (coordinate) — both electrons of the shared pair come from the same atom, e.g. NH₄⁺ and H₃O⁺.
The classic misconception: melting CO₂ does not break C=O bonds. It only overcomes the weak London forces between molecules. That single distinction explains why CO₂ sublimes at −78 °C while SiO₂ melts above 1600 °C.
Sort it
Which type of bonding?
Tap a substance, then tap the type of bonding it has.
⚡ Ionic
🔩 Metallic
🔗 Covalent
Shapes and polarity
Electron pair repulsion and intermolecular forces
Electron pairs around a central atom repel to get as far apart as possible. Lone pairs repel more strongly than bonding pairs, closing the bond angle by about 2.5° each.
Electronegativity is the power of an atom to attract the shared pair in a covalent bond. It increases across a period and decreases down a group; fluorine is the most electronegative element.
Three intermolecular forces, in increasing strength:
London (van der Waals) forces — present in all molecules; stronger the more electrons a molecule has.
Permanent dipole–dipole — between polar molecules.
Hydrogen bonding — only when H is bonded directly to N, O or F.
A molecule with polar bonds can still be non-polar overall. CO₂ is linear and CCl₄ is tetrahedral, so their bond dipoles are symmetrical and cancel. H₂O is bent, so its dipoles do not cancel and the molecule is polar.
Match it
Match the molecule to its shape
Tap a molecule on the left, then its shape and bond angle on the right.
Molecule
Shape and angle
Quick check
Think it through
?Which molecule has hydrogen bonding as its strongest intermolecular force?
Quick check
Think it through
?Why does water have a much higher boiling point than hydrogen sulfide, H₂S, even though H₂S has more electrons?
Redox
Oxidation numbers
Oxidation is loss of electrons (the oxidation number rises); reduction is gain (it falls). The oxidising agent is itself reduced; the reducing agent is itself oxidised.
Uncombined element = 0; a simple ion = its charge.
O is usually −2 (−1 in peroxides); H is usually +1 (−1 in metal hydrides); F is always −1.
The oxidation numbers sum to 0 in a neutral compound and to the charge in an ion.
Worked example — S in SO₄²⁻
x + 4(−2) = −2 → x − 8 = −2 → x = +6.
Quick check
Think it through
?What is the oxidation number of nitrogen in the nitrate ion, NO₃⁻?
Quick check
Think it through
?Magnesium chloride melts at 714 °C, but solid carbon dioxide sublimes at −78 °C. Which explanation is correct?
Recap
The big ideas to know
Isotopes: Ar = weighted mean from the mass spectrum
The mole: n = m/M · n = cV · pV = nRT (Pa, m³, K)
Formulae: empirical = simplest ratio (÷ by Ar, then ÷ by the smallest)
Ionisation energy: defined for one mole of gaseous atoms; dips at Al (3p) and S (paired 3p)
Bonding: ionic (transfer) · covalent (share) · metallic (delocalised sea) · dative (both electrons from one atom)