Basic Atomic Structure

Our understanding of atoms evolved over centuries through key experiments and models:

• Dalton's atomic theory (1807) — Matter consists of indivisible atoms; atoms of the same element are identical in mass and properties; compounds form when atoms of different elements combine in simple whole-number ratios. These postulates explained the laws of definite and multiple proportions.
• Thomson's “plum-pudding” model (1897) — The discovery of the electron via cathode-ray experiments showed atoms contain negatively charged subatomic particles embedded in a diffuse positive charge.
• Millikan's oil-drop experiment (1909) — Measured the charge of an individual electron (1.602 × 10^-19 C), allowing precise determination of its mass (9.109 × 10^-28 g).
• Rutherford's gold-foil experiment (1911) — Most alpha particles passed straight through a thin gold foil, but a few deflected sharply. This proved the atom's mass is concentrated in a tiny, dense, positively charged nucleus, with electrons occupying the surrounding space.
• Bohr model and beyond — Later work introduced energy levels and quantum mechanics, but the nuclear model remains the foundation of our picture of the atom.

Three subatomic particles account for the structure of every atom:

Protons and neutrons reside in the extremely small, dense nucleus, which has a radius on the order of 10^-15 m. Electrons occupy the vast space around the nucleus (atomic radius ∼10^-10 m). Because an electron's mass is roughly 1/1836 that of a proton, nearly all of an atom's mass comes from its nucleus. In a neutral atom, the number of protons equals the number of electrons, so the charges balance to zero.

Two integers define the composition of any specific nuclide:

• Atomic number (Z) — the number of protons in the nucleus. This determines the element's identity. Every carbon atom has Z = 6; every iron atom has Z = 26.
• Mass number (A) — the total number of protons plus neutrons: A = Z + N, where N is the neutron count.

From these two values you can find any particle count:

• Number of protons = Z
• Number of neutrons = A − Z
• Number of electrons = Z (for a neutral atom) or Z − charge (for an ion)

For example, a neutral atom of gold has Z = 79 and A = 197, so it contains 79 protons, 118 neutrons (197 − 79), and 79 electrons.

When an atom gains or loses electrons it forms an ion. Metals typically lose electrons to form cations (positive charge), while nonmetals gain electrons to form anions (negative charge). The number of protons and neutrons does not change when an ion forms.

To find electron count in an ion, start from the neutral atom count (Z) and adjust by the charge:

• Na (Z = 11) → Na⁺: 11 − 1 = 10 electrons
• Cl (Z = 17) → Cl^-: 17 + 1 = 18 electrons
• Fe (Z = 26) → Fe³⁺: 26 − 3 = 23 electrons
• O (Z = 8) → O^2-: 8 + 2 = 10 electrons

The general formula: electrons = Z − (charge), where the charge carries its sign. A +3 charge means 3 fewer electrons; a −2 charge means 2 more electrons.

Isotopes are atoms of the same element (same number of protons) that differ in the number of neutrons, and therefore differ in mass number. For example, carbon has three naturally occurring isotopes:

• ¹²C — 6 protons, 6 neutrons (mass number 12)
• ¹³C — 6 protons, 7 neutrons (mass number 13)
• ¹⁴C — 6 protons, 8 neutrons (mass number 14)

Isotopic notation places the mass number as a superscript and the atomic number as a subscript to the left of the element symbol. In plain text this is often written as carbon-12 or C-12. All isotopes of an element have the same chemical properties (same electron configuration) but different nuclear properties and atomic masses.

The atomic mass listed on the periodic table is a weighted average of the masses of all naturally occurring isotopes, reflecting their relative abundances. The calculation follows this pattern:

average atomic mass = ∑(fractional abundance × isotopic mass)

For example, chlorine has two major isotopes: ³⁵Cl (mass 34.969 amu, 75.76% abundant) and ³⁷Cl (mass 36.966 amu, 24.24% abundant):

average = (0.7576 × 34.969) + (0.2424 × 36.966) = 26.50 + 8.96 = 35.46 amu

This explains why the periodic-table value for chlorine (35.45) is not a whole number — it reflects the mixture of isotopes found in nature, not the mass of any single atom. The more abundant isotope “pulls” the average closer to its mass.

Atomic-structure problems are straightforward once you know the patterns, but a few errors appear frequently:

• Confusing atomic number with mass number. The atomic number (Z) counts only protons; the mass number (A) counts protons plus neutrons. If a problem says “sodium-23,” the 23 is A, not Z.
• Forgetting to adjust electrons for ions. A neutral atom has electrons = Z. For cations, subtract the charge; for anions, add the magnitude of the charge. Writing Fe³⁺ with 26 electrons (neutral count) instead of 23 is a common slip.
• Using mass number in average-mass calculations. The weighted-average formula uses precise isotopic masses (e.g., 34.969 amu for ³⁵Cl), not the whole-number mass numbers.
• Mixing up isotope notation. In the symbol ^A_ZX, the superscript is the mass number and the subscript is the atomic number — not the other way round.

Quick self-check: after solving any particle-count problem, verify that protons + neutrons = mass number and that the net charge equals protons − electrons.