Atomic Theory and Models

Why Atomic Structure Matters in CDS

The atom is the smallest unit of an element that takes part in a chemical reaction. The word “atom” comes from the Greek atomos, meaning “indivisible” — a name that turned out to be only partly true, because we now know the atom can be split into smaller particles. Understanding its structure is the foundation for the periodic table, chemical bonding, valency and almost every other chemistry topic in the CDS syllabus.

In the CDS General Science paper, atomic-structure questions are usually direct, single-fact recall — who discovered the electron, what is the charge on a proton, which model introduced the nucleus, or which shell holds how many electrons. They demand no algebra, so a few hours of focused memorisation here gives a very high marks-per-minute return compared with calculation-heavy physics topics.

The best way to study this chapter is as a story of correction: each scientist accepted what worked in the earlier model and repaired its single biggest flaw. Once you see it as a chain, the dates, names and discoveries stop being random facts and start making logical sense, which makes them far easier to recall under exam pressure.

Dalton's Atomic Theory (1808)

John Dalton gave the first scientific theory of the atom. Its key postulates were:

• All matter is made of tiny indivisible particles called atoms.
• Atoms can neither be created nor destroyed (basis of conservation of mass).
• Atoms of the same element are identical in mass and properties; atoms of different elements differ.
• Atoms combine in simple whole-number ratios to form compounds.

Dalton’s theory beautifully explained the laws of chemical combination — the law of conservation of mass (matter is neither created nor destroyed in a reaction) and the law of constant proportions (a compound always contains the same elements in the same fixed ratio by mass). For example, pure water always has hydrogen and oxygen in a 1:8 mass ratio, no matter the source. Dalton’s atoms explained why these laws hold.

Its weakness was the very idea that gave the atom its name: Dalton believed the atom was indivisible and structureless. The discovery of charged particles within the atom — starting with the electron — soon proved this wrong, while leaving the rest of his theory broadly intact and still useful today.

The Three Sub-atomic Particles

By the early 1900s, scientists found the atom was not indivisible — it contains three fundamental particles. This single table is the most frequently tested fact in the chapter.

• Electron (e⁻): charge −1, negligible mass (1/1837 u), discovered by J. J. Thomson (1897) in cathode-ray experiments.
• Proton (p⁺): charge +1, mass ≈ 1 u, discovered by E. Goldstein (anode/canal rays); the name “proton” is credited to Rutherford.
• Neutron (n⁰): no charge, mass ≈ 1 u, discovered by James Chadwick (1932).

Thomson's Plum-Pudding Model (1898)

Having discovered the electron, J. J. Thomson proposed that the atom is a sphere of positive charge with negatively charged electrons embedded in it — like seeds in a watermelon or plums in a pudding.

• The positive and negative charges are equal, so the atom as a whole is neutral.
• This was the first model to show the atom had internal structure rather than being a featureless ball.

Thomson reached the electron by studying cathode rays — the glowing stream produced in a discharge tube at low pressure. These rays bent towards the positive plate, proving they were made of negatively charged particles, and they behaved identically whatever gas or electrode metal was used, proving the electron is a universal building block of all matter.

The model’s flaw: it gave no place to a concentrated nucleus and spread the positive charge thinly throughout the atom. It therefore could not explain why some particles in Rutherford’s later scattering experiment bounced almost straight back — a thin haze of positive charge could never produce such a violent deflection.

Rutherford's Nuclear Model (1911)

Ernest Rutherford fired fast-moving, positively charged α-particles at a thin gold foil. The observations changed chemistry forever:

• Most particles passed straight through → the atom is mostly empty space.
• A few were deflected by large angles → there is a small, dense, positively charged centre.
• Very few (about 1 in 12,000) bounced straight back → this centre is extremely tiny but holds nearly all the mass.

From this Rutherford concluded the atom has a tiny dense nucleus at the centre, carrying the positive charge and nearly all the mass, with electrons revolving around it like planets around the Sun. To grasp the scale: if the atom were the size of a sports stadium, the nucleus would be no bigger than a pea at the centre — which is why the atom is described as mostly empty space.

This model is sometimes called the planetary model of the atom. It correctly placed the proton-bearing nucleus at the heart of the atom, an idea that has never been overturned. What it could not do was explain how the orbiting electrons remained stable.

Bohr's Model of the Atom (1913)

Niels Bohr refined Rutherford’s model by adding quantum ideas:

• Electrons revolve only in certain fixed circular paths called orbits or energy shells (K, L, M, N…).
• While moving in a permitted shell, an electron does not radiate energy — these are “stationary states”.
• Energy is absorbed or emitted only when an electron jumps from one shell to another (ΔE = E₂ − E₁).

By declaring that electrons in a permitted shell simply do not lose energy, Bohr removed the “spiralling collapse” problem of the Rutherford model. When an electron absorbs energy it jumps to a higher shell (excitation); when it falls back to a lower shell it emits the energy as light of a definite colour. This is exactly why heated hydrogen gives a few sharp coloured lines rather than a continuous rainbow — Bohr’s model explained the line spectrum of hydrogen for the first time.

The shells are labelled with both numbers and letters: n = 1 is the K shell, n = 2 is L, n = 3 is M, n = 4 is N, and so on outward. The closer a shell is to the nucleus, the lower its energy and the more tightly its electrons are held.

Atomic Number, Mass Number and Isotopes

Two numbers fix the identity of any atom:

• Atomic number (Z) = number of protons = number of electrons (in a neutral atom). It decides which element it is.
• Mass number (A) = number of protons + number of neutrons.

So number of neutrons = A − Z. An atom is written as _Z^AX, e.g. ₆¹²C means carbon with 6 protons and 6 neutrons.

• Isotopes: same Z, different A (same protons, different neutrons) — e.g. ¹H, ²H, ³H.
• Isobars: same A, different Z — e.g. ₂₀⁴⁰Ca and ₁₈⁴⁰Ar.
• Isotones: same number of neutrons, different Z and A — e.g. ₆¹⁴C and ₈¹⁶O both have 8 neutrons.

Because isotopes have the same number of protons and electrons, they have identical chemical properties but slightly different masses (and therefore different physical properties such as density). This is why chlorine’s atomic mass is the fractional 35.5 u — it is the weighted average of its two main isotopes, chlorine-35 and chlorine-37, present in roughly a 3:1 ratio in nature.

Writing Electronic Configurations

Electrons fill shells from the innermost outward, obeying the 2n² rule and the octet limit on the outermost shell. The number of electrons in the outermost shell is the valence electron count, which decides chemical reactivity and valency.

• Hydrogen (Z=1): K=1
• Carbon (Z=6): K=2, L=4
• Sodium (Z=11): K=2, L=8, M=1
• Chlorine (Z=17): K=2, L=8, M=7

Sodium has 1 valence electron (tends to lose it → +1), chlorine has 7 (tends to gain one → −1). Together they form NaCl, common table salt. The rule of thumb is simple: atoms with 1, 2 or 3 valence electrons usually lose them to form positive ions, while those with 5, 6 or 7 usually gain electrons to form negative ions, both heading towards a stable octet.

There is also an order in which shells fill. The K shell (max 2) fills first, then L (max 8), then M. Note a subtle but tested point: although the M shell can ultimately hold 18 electrons, while it is the outermost shell it stops at 8, and the next two electrons begin the N shell before M fills further. This is why potassium (Z = 19) is written 2, 8, 8, 1 and not 2, 8, 9.

Worked Example: Identifying an Atom

Common Mistakes to Avoid

• Confusing mass number with atomic number — Z = protons, A = protons + neutrons.
• Crediting Rutherford with discovering the neutron — that was Chadwick.
• Thinking the outermost shell can hold up to 2n² electrons — it is capped at 8.
• Saying the electron has the same mass as a proton — the electron is about 1837 times lighter.
• Mixing up the plum-pudding (Thomson) and nuclear (Rutherford) models.

Previous-Year Style Question

Quick Revision