Structure of the Atom

By the late 1800s, experiments with electricity through gases and salt solutions showed that matter contains charged particles, so the atom cannot be the indivisible, uncuttable thing its name (Greek 'a-tomio', not to be cut) originally meant.

J.J. Thomson discovered the electron from cathode ray tube experiments. Cathode rays travel from the negative electrode to the positive electrode, and in an electric field they bend towards the positive plate, showing they carry a negative charge. Thomson found this negative particle is present in atoms of every element, so it is a universal, fundamental constituent of matter. An electron carries a charge of −1 unit (−1.6 × 10⁻¹⁹ C) and a mass about 1/2000th that of a hydrogen atom.

E. Goldstein discovered canal rays (also called anode rays) — streams of positively charged particles produced in a discharge tube with a perforated cathode, moving opposite to cathode rays. This led to the discovery of the proton, a particle of charge +1 unit (+1.6 × 10⁻¹⁹ C) and mass about 1 atomic mass unit (u), roughly 2000 times the mass of an electron.

Since ordinary matter is electrically neutral, an atom must carry equal numbers of positive and negative charges: the positive charge of its protons exactly balances the negative charge of its electrons.

Thomson proposed that an atom is a uniformly positively charged sphere, with electrons embedded in it like seeds in a watermelon or plums in a pudding — this is why it is called the plum-pudding model (also the watermelon model).

The important features are: (i) the positive charge is spread throughout the sphere, not concentrated anywhere; (ii) the negatively charged electrons are studded inside this sphere; (iii) the magnitude of the positive charge equals the total negative charge of the electrons, so the atom as a whole is electrically neutral.

This model could explain the electrical neutrality of the atom, but it failed to explain the results of later experiments — particularly Rutherford's alpha-particle scattering experiment, which showed that positive charge is not spread out but concentrated in a tiny region.

Ernest Rutherford directed fast-moving alpha (α) particles (doubly charged helium ions, He²⁺) at a thin gold foil (about 100 nm thick) with a circular fluorescent zinc sulphide screen around it to detect where the particles landed. This is the alpha-particle scattering experiment.

Observations: (i) Most of the fast-moving alpha particles passed straight through the gold foil without any deflection. (ii) Some particles were deflected by small angles. (iii) A very small fraction (about 1 in 12000) bounced back almost 180°.

Conclusions: (i) Since most alpha particles passed through undeflected, most of the space inside the atom is empty. (ii) Since a few positively charged, fast alpha particles were deflected, they must have been repelled by a small region of concentrated positive charge — this region is the nucleus. (iii) Since a very few particles bounced straight back, nearly all the mass of the atom is concentrated in this very small nucleus.

From these results Rutherford proposed a new model: a tiny, dense, positively charged nucleus at the centre carrying almost the entire mass of the atom, with electrons revolving around it in circular orbits, much as planets revolve around the sun. The size of the nucleus is very small compared to the size of the atom (roughly 1/100000th).

Drawback: according to the electromagnetic theory of that time, a charged particle moving in a circular path (an accelerating charge) must continuously radiate energy. If an electron kept losing energy this way, it would spiral inward and eventually fall into the nucleus, making the atom unstable. But atoms are known to be stable, so Rutherford's model could not explain this.

To fix the instability problem of Rutherford's model, Niels Bohr proposed a new model with two main postulates. (i) Only certain special orbits, called discrete orbits or shells, are allowed inside the atom; these are the only orbits in which an electron is permitted to move. (ii) While revolving in these fixed, discrete orbits, an electron does not radiate energy, even though it is accelerating.

These shells, or energy levels, are represented by the letters K, L, M, N, ... going outward from the nucleus, corresponding to n = 1, 2, 3, 4, .... An electron stays at a fixed energy in a given shell; it neither loses nor gains energy as long as it stays in that shell, which is why the atom is stable in Bohr's picture. Electrons can jump between shells only by absorbing or releasing a fixed amount of energy.

Bohr's model gave a working, stable picture of the atom that matched the chemistry of the elements very well, and it is the model used through the rest of school-level chemistry (shells, valence electrons, valency), even though the modern quantum-mechanical model has since replaced the idea of a sharp circular orbit with a probability cloud.

A third sub-atomic particle, the neutron, was discovered by J. Chadwick in 1932. Neutrons are present in the nucleus of all atoms except the ordinary hydrogen atom (protium), which has only one proton and no neutron.

A neutron has no charge (it is electrically neutral) and its mass is nearly equal to that of a proton, about 1 u. Together, protons and neutrons make up the mass of the atom and are collectively called nucleons. The number of neutrons in an atom is found by subtracting the atomic number (number of protons) from the mass number: number of neutrons = mass number − atomic number.

Bohr and Bury worked out the rules for filling electrons into shells, known as the Bohr-Bury scheme. The shells are numbered n = 1, 2, 3, 4, ... from the nucleus outward and named K, L, M, N, .... The maximum number of electrons that can be accommodated in a shell is given by the formula 2n².

So the K shell (n = 1) can hold a maximum of 2 electrons, the L shell (n = 2) a maximum of 8, the M shell (n = 3) a maximum of 18, and the N shell (n = 4) a maximum of 32.

Two more practical rules apply at the school level: (i) the outermost shell of an atom can hold a maximum of 8 electrons (even if 2n² allows more); (ii) electrons are not accommodated in a given shell unless the inner shells are filled first, i.e. shells are filled in a stepwise manner starting from K.

For example, sodium (atomic number 11) has the electron distribution 2, 8, 1 — 2 in K, 8 in L, and the remaining 1 in M. Magnesium (atomic number 12) is 2, 8, 2. Argon (atomic number 18) is 2, 8, 8, filling K and L completely and putting all 8 remaining electrons in M (even though M could hold up to 18, the outermost-shell-of-8 rule caps it here). This pattern of numbers, written for each element, is called its electronic configuration.

The electrons present in the outermost shell of an atom are called valence electrons. It is these electrons that decide how an atom combines with other atoms, which is why the outermost shell is also called the valence shell.

Atoms with a completely filled outermost shell (8 electrons, or 2 for a K-shell-only atom like helium) show almost no chemical activity — these are the noble gases, and their combining capacity, or valency, is taken as zero.

For other atoms, valency is found from the number of valence electrons: (i) if the outermost shell has 1, 2, 3 or 4 electrons, the valency equals that number directly (the atom tends to lose these electrons); (ii) if the outermost shell has 5, 6 or 7 electrons, the valency is found by subtracting that number from 8 (the atom tends to gain electrons to complete its octet).

For example, sodium (2, 8, 1) has 1 valence electron, so its valency is 1. Magnesium (2, 8, 2) has valency 2. Chlorine (2, 8, 7) has 7 valence electrons, so its valency is 8 − 7 = 1. Oxygen (2, 6) has 6 valence electrons, so its valency is 8 − 6 = 2.

Every element has a fixed atomic number (Z), equal to the number of protons present in the nucleus of one atom of that element. Because an atom is neutral, the atomic number also equals the number of electrons in a neutral atom. Atomic number is the most fundamental identity of an element — it is what distinguishes one element from another.

The mass number (A) of an atom is the total number of protons and neutrons present in its nucleus (its nucleons); electrons contribute negligible mass and are not counted. Protons and neutrons, being of almost equal mass and both present in the nucleus, together decide the mass of an atom.

An element is represented in full by writing its mass number as a left superscript and atomic number as a left subscript before its chemical symbol, for instance sodium is written as ²³₁₁Na, meaning 11 protons, 11 electrons and 23 − 11 = 12 neutrons.

Isotopes are atoms of the same element that have the same atomic number but different mass numbers, because they differ in the number of neutrons while having the same number of protons. Hydrogen has three isotopes: protium (¹₁H, no neutron), deuterium (²₁H, 1 neutron) and tritium (³₁H, 2 neutrons). Carbon has isotopes carbon-12, carbon-13 and carbon-14; chlorine occurs in nature mainly as chlorine-35 and chlorine-37 in the ratio 3:1.

Because isotopes of an element have the same number of electrons and hence the same valence-electron arrangement, their chemical properties are identical; only their physical properties (like density) differ, since these depend on mass. This is why isotopes are placed at the same position in the periodic table.

Isotopes have important uses: an isotope of uranium (U-235) is used as fuel in nuclear reactors; an isotope of cobalt (Co-60) is used in the treatment of cancer; an isotope of iodine (I-131) is used in the treatment of goitre.

The average atomic mass of an element with isotopes is a weighted average of the isotopic masses, taken according to their relative natural abundance — for chlorine, (35 × 75 + 37 × 25)/100 = 35.5 u, which is why chlorine's atomic mass in the periodic table is not a whole number.

Isobars, by contrast, are atoms of different elements that have the same mass number but different atomic numbers, because their combined proton-plus-neutron counts happen to match even though the number of protons differs. A classic pair is calcium (atomic number 20) and argon (atomic number 18), both with mass number 40. Since isobars have different numbers of protons and electrons, they are different elements with different chemical properties.