Lesson 1: Atomic Structure and the Periodic Table

Every substance you have ever touched is built from atoms, and there are only about 90 naturally occurring kinds of them. Chemistry begins with a model of the atom that explains why those kinds — the elements — behave so differently, and with the table that organises them: the periodic table. This lesson builds that model and shows how the table predicts chemical behaviour before you ever mix two substances.

Inside the atom

An atom consists of a tiny, dense nucleus containing protons (charge +1, mass approximately 1 u) and neutrons (charge 0, mass approximately 1 u), surrounded by electrons (charge −1, mass roughly 1/1836 u, usually treated as negligible). The atomic number Z is the number of protons: it defines the element. Carbon always has Z = 6; a nucleus with 7 protons is nitrogen, whatever else it contains. The mass number A counts protons plus neutrons. In a neutral atom, electrons = protons; if electrons are gained or lost, the atom becomes an ion with a net charge.

Atoms of the same element with different neutron counts are isotopes. They have identical chemistry (chemistry is governed by electrons) but different masses and sometimes different nuclear stability. Chlorine, for example, is a mixture of ³⁵Cl and ³⁷Cl — which is why its relative atomic mass on the periodic table, 35.45, is not a whole number and corresponds to no actual chlorine atom.

Electron arrangement and the periodic table

Electrons occupy shells (energy levels), and within shells, orbitals. The first shell holds 2 electrons, the second 8, the third 18, filling in a pattern summarised by electron configurations such as magnesium: 1s² 2s² 2p⁶ 3s². The electrons in the outermost shell — the valence electrons — do almost all the chemistry, because they are the ones lost, gained or shared when atoms react.

The periodic table arranges elements by increasing Z. Groups (columns) share the same number of valence electrons, so they react similarly: Group 1 metals all react vigorously with water. Periods (rows) correspond to filling a new shell. Across a period, nuclear charge rises while electrons enter the same shell, so atoms get smaller and attract electrons more strongly — the pattern behind the periodic trends in atomic radius, ionisation energy and electronegativity that later lessons use constantly.

Worked example 1: average atomic mass

A sample of chlorine contains 75.77% ³⁵Cl (atomic mass 34.97 u) and 24.23% ³⁷Cl (atomic mass 36.97 u). Calculate the relative atomic mass of chlorine. Step 1: convert percentages to fractions: 0.7577 and 0.2423. Step 2: weight each isotope mass by its abundance and add: (0.7577 × 34.97) + (0.2423 × 36.97) = 26.50 + 8.96 = 35.45 u (4 significant figures, matching the data). Notice the answer lies much closer to 35 than to 37, because the lighter isotope dominates. You can reverse the logic: an element with two isotopes of mass 62.93 u (69.17%) and 64.93 u (30.83%) has an average of (0.6917 × 62.93) + (0.3083 × 64.93) = 43.53 + 20.01 = 63.55 u — that is copper, identifiable from its mass spectrum alone.

Worked example 2: particles in an ion

Determine the numbers of protons, neutrons and electrons in ⁵⁶Fe and in the Fe³⁺ ion (Z = 26). For ⁵⁶Fe: protons = Z = 26. Neutrons = A − Z = 56 − 26 = 30. A neutral atom has electrons = protons = 26. For Fe³⁺: the nucleus is unchanged (26 protons, 30 neutrons), but three electrons have been removed to give the 3+ charge, so electrons = 26 − 3 = 23. The method matters more than the numbers: ions change only the electron count, never the proton count — changing protons changes the element itself.

In the lab: flame tests and their limits

Flame tests identify metal ions by the colour their electrons produce when heated electrons fall back to lower energy levels, emitting light of characteristic wavelengths. In a school determination, a nichrome wire is dipped in concentrated acid, then in the sample, and held at the edge of a blue Bunsen flame: lithium gives crimson red, sodium an intense orange-yellow, potassium lilac, calcium orange-red, copper blue-green.

The results carry real experimental uncertainty. A single observation is one trial with a subjective colour judgement — "orange-red" versus "red" is an opinion, not a measurement. Sodium is a near-universal contaminant (from glass, dust and fingers), and its intense yellow can mask potassium's lilac entirely; analysts view potassium flames through blue cobalt glass to filter the sodium light out. A careful worker runs each test three times, cleans the wire in acid until it gives no colour of its own, and records a blank. If the wire is not cleaned between samples, residue from the previous sample produces a mixed colour — a systematic error that replication alone will not reveal, because every repeat repeats the same contamination.

Chemistry in the real world

Mass spectrometry — the instrument behind Worked example 1 — is how forensic scientists confirm the identity of a drug sample and how geologists date rocks from isotope ratios such as ²⁰⁶Pb/²³⁸U. The same instrument sits in hospitals, where breath tests using ¹³C-labelled urea detect the bacterium that causes stomach ulcers.

Common mistakes

• Confusing mass number with relative atomic mass. A is an integer for one isotope; the periodic table value is an abundance-weighted average over all isotopes.
• Letting ions change the proton count. Fe²⁺ and Fe³⁺ differ only in electrons.
• Giving electrons in shells as (2, 8, 18) capacities for configuration problems at this level. Follow the configuration pattern (1s, 2s, 2p, 3s, 3p, 4s…) and count from it.
• Reporting average atomic mass to the wrong precision. Keep at least 4 significant figures; rounding 35.45 to 35 will cost marks and, in later lessons, moles.

Exam Experience

1. Which statement about isotopes of the same element is correct? A) They have different numbers of protons. B) They have identical chemical properties but different masses. C) They have different numbers of electrons in neutral atoms. D) They occupy different positions in the periodic table. Answer: B. Isotopes differ only in neutron number, so mass differs; electron arrangement — and therefore chemistry — is unchanged. A and C describe different elements or ions, and D is wrong because position is set by Z.

2. What is the electron configuration of a calcium atom (Z = 20)? A) 1s² 2s² 2p⁶ 3s² 3p⁶ 3d² B) 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² C) 2, 8, 10 D) 1s² 2s² 2p⁶ 3s² 3p⁸ Answer: B. The 4s orbital fills before 3d; Ca ends …3p⁶ 4s². A puts electrons in 3d too early; C misallocates the third shell; D exceeds the p-subshell capacity of 6.

3. Moving down Group 1 from lithium to caesium, reactivity with water increases mainly because: A) the atoms gain protons. B) the outer electron is further from the nucleus and less strongly attracted, so it is lost more easily. C) the metals become less dense. D) the number of electron shells decreases. Answer: B. Reactivity in Group 1 is about losing the single valence electron; increased distance and shielding weaken its attraction, lowering the energy needed. D states the opposite of the trend.

Constructed response (4 marks). Boron has two isotopes: ¹⁰B (19.9%, mass 10.013 u) and ¹¹B (80.1%, mass 11.009 u). (a) Calculate the relative atomic mass of boron. (b) State the number of neutrons in a ¹¹B atom (Z = 5). Scoring guide: (a) 1 mark for correct weighting set-up (0.199 × 10.013) + (0.801 × 11.009); 1 mark for the value 10.81 u; 1 mark for 4 significant figures / sensible precision. Accept 10.81 ± 0.01. (b) 1 mark for 11 − 5 = 6 neutrons. Common error: weighting by mass number only (10 and 11) gives 10.80 — award the set-up mark only if exact isotopic masses were clearly intended.

Reading a mass spectrum: worked example 3

A mass spectrometer vaporises a sample, ionises the atoms, and sorts them by mass-to-charge ratio, printing a stick spectrum of relative abundance against m/z. Magnesium's spectrum shows three peaks: m/z 24 (78.99%, mass 23.985 u), 25 (10.00%, 24.986 u) and 26 (11.01%, 25.983 u). The relative atomic mass is the same weighted average as Worked example 1: (0.7899 × 23.985) + (0.1000 × 24.986) + (0.1101 × 25.983) = 18.94 + 2.50 + 2.86 = 24.31 u — the periodic table value. Two features make mass spectra more powerful than a single average: the *pattern* of peaks is a fingerprint (chlorine's 3 : 1 pair at 35/37 is unmistakable), and in molecular work (Lesson 17) the heaviest peak often reveals the molecular mass directly.

Ionisation energy: the trend that runs the table

The first ionisation energy — the energy to remove one mole of electrons from one mole of gaseous atoms — ties the table's geometry to its chemistry. Across a period it generally rises (greater nuclear charge, similar shielding, smaller atoms). Down a group it falls (the electron is further away and better shielded). The exceptions teach the structure: aluminium dips below magnesium because its outer electron sits in a higher-energy 3p orbital, and sulfur dips below phosphorus because paired electrons in one p-orbital repel. Group 1's falling ionisation energies (Li 520, Na 496, K 419, Rb 403, Cs 376 kJ/mol) are exactly why reactivity with water explodes down the group — Lesson 1's flame colours come with fireworks attached.

Successive ionisation energies: reading the shells

Remove electrons one at a time from sodium and the energies (kJ/mol) run: 496, 4562, 6910, 9543… The first electron leaves cheaply — it is the lone 3s electron, far from the nucleus and shielded by a full inner shell. The second costs nine times more because it must be torn from the stable n = 2 shell, much closer to an effectively stronger nuclear pull. That single jump is direct experimental evidence for shell structure: the table of successive ionisation energies is the atom's X-ray. Magnesium shows its jump after two electrons (738, 1451, then 7733), aluminium after three — the jump position counts the valence electrons, confirming group numbers from physics rather than from reaction patterns. Examiners love this data because it cannot be memorised as a slogan; it must be read.