An atom is the smallest particle of an element that can take part in a chemical reaction. For a long time atoms were considered indivisible, but experiments in the late 19th and early 20th century revealed subatomic particles. Electron: discovered by J.J. Thomson (1897) using cathode ray tube experiments. He found negatively charged particles in all elements — these were electrons. Proton: discovered by Ernest Rutherford (1919) — positively charged particles in the nucleus. Neutron: discovered by James Chadwick (1932) — neutral particles in the nucleus. Charge and mass: electron (-1, negligible mass), proton (+1, 1 u), neutron (0, 1 u). Atomic number (Z) = number of protons. Mass number (A) = protons + neutrons.

J.J. Thomson proposed the first model of the atom in 1904, known as the "Plum Pudding Model" or "Watermelon Model." According to this model: The atom is a positively charged sphere (like a watermelon) with negatively charged electrons embedded in it (like seeds). The positive and negative charges balance each other making the atom electrically neutral. This model explained why atoms are electrically neutral. Limitation: It could not explain the results of Rutherford's gold foil experiment — it could not account for the existence of a concentrated positive nucleus.

Ernest Rutherford (1911) bombarded a thin gold foil with high-speed alpha (α) particles (positively charged helium nuclei). He placed a fluorescent screen around the foil to detect where particles went. Observations: (1) Most α-particles passed straight through the foil. (2) Some were deflected at small angles. (3) A very few (1 in 20,000) bounced straight back. Conclusions: (1) Most of the atom is empty space (most particles passed through). (2) Positive charge and most mass are concentrated in a tiny central region called the nucleus. (3) Electrons revolve around the nucleus at a distance. Limitation: According to classical physics, an electron in circular orbit should continuously lose energy and spiral into the nucleus — the atom should be unstable. Rutherford's model could not explain atomic spectra either.

Niels Bohr (1913) improved Rutherford's model to overcome its limitations, using Planck's quantum theory. Postulates: (1) Electrons revolve around the nucleus in fixed circular paths called orbits or shells or energy levels (K, L, M, N...). (2) As long as an electron stays in a given orbit it does not lose energy — these are called stationary states. (3) Electrons can jump to higher energy levels by absorbing energy and fall back to lower levels by emitting energy (as light). This explained atomic spectra. Significance: Successfully explained the hydrogen spectrum. Shells are designated K (n=1), L (n=2), M (n=3), N (n=4). Limitation: Could not explain spectra of multi-electron atoms.

Electrons are arranged in energy shells around the nucleus. Rules for filling electrons: (1) Maximum electrons in a shell = 2n² (where n = shell number). K shell: 2×1² = 2; L shell: 2×2² = 8; M shell: 2×3² = 18. (2) The outermost shell can hold maximum 8 electrons. (3) Shells are filled in order (K first, then L, then M...). Examples: Hydrogen (1): K=1. Helium (2): K=2. Lithium (3): K=2, L=1. Carbon (6): K=2, L=4. Sodium (11): K=2, L=8, M=1. Chlorine (17): K=2, L=8, M=7. Calcium (20): K=2, L=8, M=8, N=2. Valence electrons = electrons in the outermost shell. Valency = how many electrons an atom needs to give/take to complete its outermost shell.

Isotopes: atoms of the same element that have the same atomic number (same protons) but different mass numbers (different neutrons). They have the same chemical properties but different physical properties. Examples: Hydrogen isotopes — Protium (¹H, 0 neutrons), Deuterium (²H, 1 neutron), Tritium (³H, 2 neutrons). Carbon isotopes — ¹²C (6n) and ¹⁴C (8n, used in radiocarbon dating). Chlorine — ³⁵Cl and ³⁷Cl (average mass 35.5). Uses of isotopes: ¹³¹I in treating thyroid cancer, ²³⁵U in nuclear fuel, ¹⁴C in carbon dating. Isobars: atoms of different elements with the same mass number but different atomic numbers. Examples: ⁴⁰Ca (Z=20) and ⁴⁰Ar (Z=18). Isobars have same mass number, different protons/neutrons.

Atomic number (Z): number of protons in the nucleus. It uniquely identifies an element. In a neutral atom, number of electrons = number of protons = Z. Valency: the combining capacity of an element. Determined by: electrons in outermost shell if ≤ 4 (valency = number of outermost electrons); if > 4, valency = 8 − outermost electrons. Examples: Carbon (4 outermost) → valency 4. Oxygen (6 outermost) → valency 2. Chlorine (7 outermost) → valency 1. Noble gases (He, Ne, Ar): 8 outermost electrons (He: 2) → valency 0 (do not react). The concept of valency explains how atoms bond: sodium (valency 1) combines with chlorine (valency 1) to form NaCl.

The nucleus is the tiny, dense, positively charged centre of the atom. Despite protons repelling each other (same positive charge), the nucleus is held together by the strong nuclear force — the strongest of the four fundamental forces, which acts over very short ranges. The nucleus contains protons and neutrons (collectively called nucleons). Nuclear properties: Radius of nucleus ≈ 10⁻¹⁵ m (nucleus is 100,000 times smaller than atom). Almost all mass of atom is in nucleus. Density of nucleus is enormously high (~10¹⁷ kg/m³). Radioactivity: some nuclei are unstable and spontaneously emit radiation (alpha, beta, gamma) to become stable — discovered by Henri Becquerel (1896). Marie and Pierre Curie studied radioactivity of uranium and polonium.

Comparison of atomic models: Thomson's model (1904) — positive sphere with embedded electrons; explained neutrality; failed to explain Rutherford's results. Rutherford's model (1911) — nuclear model with electrons orbiting nucleus; explained gold foil experiment; failed to explain stability and spectra. Bohr's model (1913) — electrons in fixed energy shells; explained hydrogen spectrum; failed for multi-electron atoms. Modern quantum mechanical model — electrons in probability clouds (orbitals), not fixed paths. Key facts: Protons and neutrons are in the nucleus. Electrons are outside nucleus. Proton mass = neutron mass ≈ 1 u = 1.67 × 10⁻²⁷ kg. Electron mass ≈ 1/1840 of proton. Electron charge = −1.6 × 10⁻¹⁹ C. Proton charge = +1.6 × 10⁻¹⁹ C.

Atomic number (Z) is the number of protons in the nucleus of an atom; it identifies the element. Mass number (A) is the total number of protons and neutrons (nucleons) in the nucleus. Number of neutrons = A − Z. In a neutral atom, number of electrons = number of protons = Z. An atom is written as ᴬ_Z X. Example: sodium ²³₁₁Na has 11 protons, 11 electrons and 23 − 11 = 12 neutrons. Electrons have negligible mass, so the mass of an atom is concentrated in its nucleus.

Electron: discovered by J.J. Thomson (1897) from cathode rays; charge −1, mass about 1/1837 of a proton. Proton: positive rays (canal rays) were observed by E. Goldstein (1886); Rutherford later identified the proton; charge +1, mass 1 u. Neutron: discovered by James Chadwick (1932); no charge, mass about 1 u. Protons and neutrons lie in the nucleus; electrons move around it. Nearly all the mass of the atom is in the nucleus, which is about 100,000 times smaller than the atom.

Rutherford's model: electrons revolve around the nucleus in circular paths. According to classical physics, a revolving charged particle is accelerated and should radiate energy, lose it and spiral into the nucleus, so the atom would collapse. The model could not explain the stability of the atom. Bohr's model overcame this by proposing that electrons revolve only in certain fixed circular orbits (energy levels or shells) called stationary orbits, in which they do not radiate energy. Bohr's model explains the stability of the atom and the hydrogen spectrum.

The electrons in the outermost shell are the valence electrons. Atoms are most stable when the outermost shell is full: an octet (8 electrons) for most elements, and a duplet (2 electrons) for helium and the K shell. Atoms combine by gaining, losing or sharing electrons to reach this stable arrangement. Valency is the combining capacity of an atom. If the valence electrons are 4 or fewer, valency equals the number of valence electrons; if more than 4, valency = 8 − valence electrons. Examples: Na (2,8,1) valency 1; Mg (2,8,2) valency 2; Al (2,8,3) valency 3; O (2,6) valency 2; Cl (2,8,7) valency 1; Ne and Ar valency 0.

Isotopes are atoms of the same element with the same atomic number but different mass numbers (different number of neutrons). Examples: hydrogen has protium (¹H), deuterium (²H) and tritium (³H); carbon has ¹²C and ¹⁴C; chlorine has ³⁵Cl and ³⁷Cl. Isotopes have the same chemical properties because they have the same electronic configuration, but slightly different physical properties. Uses: uranium-235 as fuel in nuclear reactors; iodine-131 for treating goitre; cobalt-60 in cancer treatment; carbon-14 in carbon dating. The atomic mass of chlorine, 35.5 u, is the average of its isotopes in the ratio about 3:1.

Shell capacity is 2n²: K = 2, L = 8, M = 18, N = 32. Configurations (K,L,M,N): H 1; He 2; Li 2,1; Be 2,2; B 2,3; C 2,4; N 2,5; O 2,6; F 2,7; Ne 2,8; Na 2,8,1; Mg 2,8,2; Al 2,8,3; Si 2,8,4; P 2,8,5; S 2,8,6; Cl 2,8,7; Ar 2,8,8; K 2,8,8,1; Ca 2,8,8,2. For the first 20 elements the outermost shell never holds more than 8 electrons and the third shell is filled up to 8 before the fourth shell starts.