Chapter 8: Journey Inside the Atom

• Acharya Kanada (India) proposed that matter, divided repeatedly, reaches indivisible parmanus. Leucippus and Democritus (Greece) called them atomos (indivisible).
• These were ideas, not experimental results.
• In 1808, John Dalton gave the first scientific atomic theory: all matter is made of indivisible atoms.
• Radioactivity later showed that atoms are made of smaller particles.

• In 1897, J. J. Thomson studied cathode rays in a gas discharge tube. They are streams of negatively charged particles, electrons, much lighter than atoms.
• Cathode rays were the same whatever the cathode material or gas, so electrons are present in all atoms. The charge of an electron is −1.602 × 10⁻¹⁹ C, taken as −1.
• Plum pudding (watermelon) model: a sphere of positive charge with electrons embedded in it.

• In 1911, Geiger and Marsden (under Rutherford) fired α-particles at thin gold foil. Most passed straight through, some were deflected sharply, a few bounced back.
• Conclusions: most of the atom is empty space; the positive charge and most of the mass are in a tiny, dense nucleus; electrons revolve around it (planetary model).
• Atom diameter ≈ 10⁻¹⁰ m, nucleus diameter ≈ 10⁻¹⁵ m (about 10⁵ times smaller).
• Limitation: an orbiting electron accelerates, should lose energy and spiral into the nucleus, but atoms are stable.
• Rutherford also showed that the nucleus carries positive charge in the form of protons.

• Electrons move in fixed circular paths called stationary states, orbits, shells or energy levels (K, L, M, N or n = 1, 2, 3, 4).
• An electron does not lose energy while in a shell. It changes shell only by absorbing or releasing a fixed amount of energy equal to the difference of the levels.
• The shell nearest the nucleus (K) has the least energy; energy increases outward.
• Each shell holds only a certain number of electrons. Bohr's model was later improved by the quantum mechanical model.

• James Chadwick discovered the neutron in 1932: mass nearly equal to a proton, no charge. It explains why a helium atom is four times as heavy as hydrogen.
• Relative charges: electron −1, proton +1, neutron 0. Protons and neutrons are in the nucleus; the electron's mass is negligible.
• Heavier atoms need more neutrons than protons to hold the nucleus together (nuclear force).

• Dalton used pictorial symbols; Berzelius (1813) used letters from Latin names. IUPAC approves names and symbols today.
• Rules: first letter capital, second small (Al, Co). Some symbols come from Latin, Greek or German names: Fe (ferrum), Hg (hydrargyros), W (wolfram), Na, K, Cu, Ag, Au, Pb.
• Atomic number (Z) = number of protons = number of electrons in a neutral atom. It identifies the element.
• Mass number (A) = protons + neutrons (nucleons). Notation: ^A_ZX, e.g. ¹²₆C.

• Maximum electrons in a shell = 2n²: K = 2, L = 8, M = 18.
• The outermost shell holds at most 8 electrons. Shells fill in order K, L, M, N.
• Examples: 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.

• The outermost shell is the valence shell; its electrons are valence electrons. A complete octet (2 for helium) makes an atom stable and unreactive.
• Valency is the number of electrons gained, lost or shared to complete the octet.
• Fewer than four valence electrons: the atom tends to lose them (Na valency 1). More than four: tends to gain (O valency 2). Four: shares (C valency 4).

• Isotopes: same atomic number, different mass numbers (¹H, ²H, ³H; ¹²C, ¹³C, ¹⁴C). They have similar chemical properties (same electrons) but different physical properties.
• Uses: ²³⁵U fuel in reactors, ⁶⁰Co cancer treatment, ¹³¹I thyroid, ¹⁴C dating fossils.
• Average atomic mass is weighted by abundance. Chlorine: 35 × 0.75 + 37 × 0.25 = 35.5 u.
• Isobars: same mass number, different atomic numbers (Ca, K, Ar all with A = 40).

Dalton (indivisible atom) → Thomson (positive sphere with electrons) → Rutherford (nuclear model) → Bohr (energy levels) → modern quantum mechanical model (electron clouds) → still being discovered.