Structure of the Atom

Why this topic matters for NDA

Chemistry in the NDA General Ability Test rewards students who know the basics cold. Structure of the Atom is one of those high-yield basics — the questions are direct, factual and rarely tricky. Because the General Ability Test packs Physics, Chemistry, Biology, History, Geography and current affairs into one paper, every guaranteed mark counts, and this chapter delivers exactly that kind of reliable mark.

You will typically see one or two questions on atomic number, mass number, isotopes, or which scientist discovered which particle. These are pure recall marks if you have the fundamentals clear. The same concepts also support harder topics later — the periodic table, chemical bonding and radioactivity all build directly on what you learn here, so time spent now pays off across the whole syllabus.

What exactly is an atom?

An atom is the smallest particle of an element that can take part in a chemical reaction. The word comes from the Greek atomos, meaning indivisible — though we now know atoms are made of even smaller parts. Atoms are incredibly small: their radius is measured in nanometres (1 nm = 10⁻⁹ m), far too tiny to see even with an ordinary microscope.

John Dalton (1808) first proposed the modern atomic theory. Its main points are that matter is made of tiny indivisible atoms, atoms of the same element are identical in mass and properties, atoms of different elements differ, and atoms combine in fixed whole-number ratios to form compounds. This theory neatly explained the laws of chemical combination such as the law of conservation of mass and the law of constant proportions.

It is important to distinguish an atom from a molecule. A molecule is formed when two or more atoms join together — for example two oxygen atoms form an O₂ molecule, and atoms of hydrogen and oxygen combine to form a water (H₂O) molecule. The atom remains the basic building block.

The three sub-atomic particles

Every atom is built from three fundamental particles. Knowing their charge, mass and discoverer is the single most important set of facts in this chapter.

• Electron (e⁻): negative charge, almost negligible mass, discovered by J.J. Thomson (1897).
• Proton (p⁺): positive charge, mass ≈ 1 atomic mass unit (u), discovered by E. Goldstein / Rutherford.
• Neutron (n⁰): no charge (neutral), mass ≈ 1 u, discovered by James Chadwick (1932).

The proton was identified through the study of positive rays (anode rays) produced by Goldstein, while Rutherford later confirmed it as a fundamental particle of the nucleus. The neutron was the last to be found because, having no charge, it is hard to detect; Chadwick's discovery finally explained why an atom's mass is greater than the mass of its protons alone.

Nucleus and electron shells

The nucleus is the dense central core of the atom. It contains protons and neutrons (together called nucleons) and carries almost the entire mass of the atom along with a net positive charge.

Electrons revolve around the nucleus in fixed circular paths called shells or energy levels, labelled K, L, M, N (or n = 1, 2, 3, 4). An atom as a whole is electrically neutral because the number of protons equals the number of electrons.

Atomic number and mass number

These two numbers define an atom and are the most frequently tested calculations in this chapter.

An element is written as _Z^AX. For example, carbon is written ₆¹²C, meaning Z = 6 (6 protons, 6 electrons) and A = 12 (so neutrons = 12 − 6 = 6).

The atomic number Z is the true identity of an element — change Z and you get a completely different element. This is exactly how the modern periodic table is arranged: elements are placed in order of increasing atomic number, not atomic mass.

Note that mass number A is always a whole number because it simply counts nucleons. The atomic mass printed in the periodic table, however, is often a decimal (for example chlorine is 35.5) because it is the average mass of all the naturally occurring isotopes of that element, weighted by how common each one is.

Isotopes, isobars and isotones

NDA loves these three look-alike terms. Learn them as a contrast set so you never mix them up.

• Isotopes: same atomic number (Z), different mass number (A). They are the same element with different numbers of neutrons. Example: ₁¹H, ₁²H (deuterium), ₁³H (tritium).
• Isobars: different Z, same mass number A. Different elements. Example: ₂₀⁴⁰Ca and ₁₈⁴⁰Ar.
• Isotones: same number of neutrons, different Z and A.

Because isotopes of the same element have the same number of electrons, they show identical chemical properties; they differ only in physical properties that depend on mass, such as density. This is why a sample of an element behaves chemically as one substance even though it may be a mix of isotopes.

Uses of isotopes are a favourite NDA fact:

• Uranium-235 — fuel in nuclear reactors and atom bombs.
• Cobalt-60 — used in radiotherapy to treat cancer.
• Carbon-14 — used in carbon dating to find the age of fossils and ancient objects.
• Iodine-131 — used to detect and treat goitre (thyroid disorders).

Thomson's model of the atom

After discovering the electron, J.J. Thomson proposed the first model of the atom in 1904.

He pictured the atom as a sphere of positive charge with electrons embedded in it — like seeds in a watermelon, or plums in a pudding. This is famously called the plum pudding model (also the watermelon model).

Rutherford's nuclear model

Ernest Rutherford fired fast-moving α-particles (positively charged) at a thin gold foil. Most passed straight through, a few were deflected, and very few bounced straight back.

From this gold foil experiment he concluded:

• Most of the atom is empty space (most particles passed through).
• There is a tiny, dense, positively charged centre called the nucleus (a few particles bounced back).
• Electrons revolve around the nucleus.

Bohr's model and energy levels

Niels Bohr (1913) improved Rutherford's model. His key idea: electrons revolve only in certain fixed orbits of definite energy called stationary states or energy levels.

• While in a fixed orbit, an electron does not radiate energy — this solves the stability problem.
• Energy is absorbed when an electron jumps to a higher orbit, and energy is emitted (often as light) when it falls back to a lower orbit.
• Energy increases as we move outward: K < L < M < N.

The fixed orbits are sometimes called quantised energy levels — an electron can only have certain allowed energies, never values in between. This single idea explained why elements give off light of specific colours when heated, something earlier models could not do.

Electronic configuration and the 2n² rule

The arrangement of electrons in the various shells is called the electronic configuration. Electrons fill shells starting from the one nearest the nucleus.

The electrons in the outermost shell are called valence electrons, and they decide the chemical behaviour of the element. There is also a filling rule worth remembering: even though an inner shell like M can eventually hold 18 electrons, it does not fill completely before the next shell starts — for the lighter elements the outer shell never takes more than 8 before electrons begin entering the following shell.

Valency from electronic configuration

Valency is the combining capacity of an atom — how many electrons it gains, loses or shares to complete its outer shell (usually to reach 8, the octet).

• If the outermost shell has 1, 2 or 3 electrons, valency = number of valence electrons (the atom loses them). Example: Na (2,8,1) → valency 1.
• If the outermost shell has 5, 6 or 7 electrons, valency = 8 − (valence electrons), because the atom gains electrons. Example: Cl (2,8,7) → valency 8 − 7 = 1.
• Atoms with a full outer shell (2 or 8) have valency 0 — the noble gases, which are unreactive.

Previous-year question and quick recap