Atomic and Nuclear Structure · 9 min read
Development of the Atomic Model
The atomic model is a textbook case of how scientific knowledge changes with new evidence. Each model explained the data available at the time, and each was revised when a new experiment produced results it could not explain. The test expects you to match each scientist with the experiment, the observation, and the conclusion, and to recognize that models are tools that are refined rather than simply "proven wrong."
John Dalton proposed that all matter is made of indivisible atoms, that atoms of one element are identical in mass and properties, and that compounds form when atoms combine in fixed whole-number ratios. His theory explained the laws of conservation of mass, definite proportions, and multiple proportions. J. J. Thomson used cathode-ray tubes and showed that the rays were deflected by electric and magnetic fields toward the positive plate, so they were streams of negatively charged particles far lighter than any atom. These electrons came from every metal tested, so atoms had internal parts. Thomson proposed the plum-pudding model: electrons embedded in a diffuse sphere of positive charge.
Robert Millikan's oil-drop experiment balanced the electric force on tiny charged oil drops against gravity. Every drop carried a whole-number multiple of one smallest charge, 1.60 × 10⁻¹⁹ C, the elementary charge. Combined with Thomson's charge-to-mass ratio, this gave the electron's mass. Ernest Rutherford's gold-foil experiment fired alpha particles at thin gold foil. Most passed straight through, but a few were deflected at large angles and some bounced almost straight back. Rutherford concluded that the atom is mostly empty space with a tiny, dense, positively charged nucleus containing nearly all of the mass.
KEY EXPERIMENTS BEHIND THE ATOMIC MODEL Scientist Evidence Observation
Matter is made of indivisible atoms
Elements combine in fixed whole-number
Dalton Mass ratios in compounds ratios
Rays bend toward the positive plate
Atoms contain negative electrons; plum- pudding model
Charges are multiples of 1.60 × 10⁻¹⁹ C
Charge is quantized; electron mass found
Most alpha particles pass through; a few
Small, dense, positive nucleus; mostly empty space
Only specific wavelengths are emitted
Electrons occupy fixed energy levels
Neutral radiation knocks protons out of
The nucleus contains neutral neutrons
Rutherford's model could not explain why electrons do not spiral into the nucleus or why heated hydrogen emits only certain colors. Niels Bohr proposed that the electron in hydrogen occupies only certain circular orbits, each with a fixed energy. When an electron drops from a higher level to a lower one, the atom emits a photon whose energy equals the difference between the levels: ΔE = hf. Because only certain energy differences exist, only certain frequencies appear, giving a line (emission) spectrum. Absorption works in reverse: an atom absorbs only photons that exactly match a level gap, producing dark lines in a continuous spectrum.
BOHR MODEL AND PHOTON EMISSION n = 3 photon visible
Electron falls from n=3 to n=2
Schematic, not to scale. An electron that drops to a lower level emits one photon with energy equal to the gap (ΔE = hf). For hydrogen, drops to n = 2 give visible lines; drops to n = 1 give larger gaps and ultraviolet photons.
A student says Rutherford's experiment proved that Thomson was a poor scientist. The teacher reframes the history as a model-revision cycle: Thomson's model explained cathode rays, and only the unexpected large-angle scattering required a new model. Asking students which result Thomson's model predicts for the gold foil (all particles pass with tiny deflections) makes the logic of the revision clear.
Chadwick and the Quantum Model
James Chadwick identified the neutron by showing that a neutral radiation from beryllium bombarded with alpha particles could knock protons out of paraffin; its particles had about the mass of a proton but no charge. This explained why atomic masses are larger than the proton count alone predicts. The current quantum-mechanical model, built on de Broglie's matter waves, Schrödinger's wave equation, and Heisenberg's uncertainty principle, replaces fixed orbits with orbitals: regions where an electron is likely to be found. Bohr's energy levels survive as principal energy levels, but exact electron paths cannot be known.
Link each observation to its inference: rays bending toward the positive plate means negative particles; large-angle rebounds mean a small dense nucleus; most particles passing through means mostly empty space; discrete spectral lines mean quantized energy levels.
Every atom has a nucleus of protons and neutrons (together called nucleons) surrounded by electrons. The proton and neutron have nearly equal masses, each close to 1 atomic mass unit (u or amu), while the electron is about 1/1836 of a proton's mass. Because electrons contribute so little mass, nearly all of an atom's mass is in the nucleus, even though the nucleus is roughly 1/100,000 of the atom's diameter.
THE THREE SUBATOMIC PARTICLES Particle Charge
Atomic Number and Mass Number
The atomic number (Z) is the number of protons, and it alone identifies the element: every carbon atom has 6 protons. The mass number (A) is the total number of protons plus neutrons, so the number of neutrons equals A − Z. Nuclide notation writes the mass number as a left superscript and the atomic number as a left subscript, as in ¹²₆C or ²³⁵₉₂U, or in words as carbon-12 and uranium-235. In a neutral atom the electron count equals Z. An ion has gained or lost electrons, never protons: Na⁺ has 11 protons and 10 electrons, and Cl⁻ has 17 protons and 18 electrons.
Isotopes are atoms of the same element with different numbers of neutrons and therefore different mass numbers. Carbon-12, carbon-13, and carbon-14 all have 6 protons, but they have 6, 7, and 8 neutrons. Isotopes of an element have essentially identical chemical behavior because chemistry depends on electrons, but they differ in mass, density, and nuclear stability. Some isotopes, such as carbon-14 and uranium-238, are radioactive. Atoms with the same mass number but different atomic numbers, such as ¹⁴C and ¹⁴N, are different elements, not isotopes.
Atomic number (Z)
Number of protons; defines the element and its place on the periodic table.
Mass number (A)
Protons plus neutrons in one specific nucleus; always a whole number.
Weighted average mass of an element's naturally occurring isotopes, in u; usually not a whole number.
Same element, same proton count, different neutron count.
The atomic mass printed on the periodic table is a weighted average of the masses of an element's natural isotopes, each multiplied by its fractional abundance. The average is pulled toward the most abundant isotope. Mass spectrometry measures both the masses and the abundances: the instrument separates ions by mass-to-charge ratio, and the heights of the peaks give relative abundances. Average atomic mass in u also equals molar mass in grams per mole, which links
Convert percentages to fractions: 0.75 and 0.25.
Check: the answer lies between 35 and 37 and closer to 35, the more abundant isotope. The table value is 35.45
The same method runs in reverse. If an element with isotopes of mass 10 u and 11 u has an average atomic mass of 10.8 u, let x be the fraction of the heavier isotope: 10(1 − x) + 11x = 10.8, so x = 0.8. The element (boron) is about 80% boron-11 and 20% boron-10. On the test, a quick estimate of where the average lies between the isotope masses often eliminates two or three answer choices without full arithmetic.
Students often confuse mass number with atomic mass. Mass number belongs to one isotope and is a whole count of nucleons; atomic mass is an average for the element. No single chlorine atom has a mass of 35.45 u.
A useful classroom diagnostic is to ask students to describe a ³⁷Cl⁻ ion completely. A correct answer gives 17 protons (from Z), 20 neutrons (37 − 17), and 18 electrons (one extra for the 1− charge). Errors reveal which idea is shaky: subtracting the charge instead of adding, or confusing mass number with neutron number.
In the quantum-mechanical model, each electron is described by its principal energy level (n = 1, 2, 3, …), its sublevel (s, p, d, f), and its orbital. An orbital holds at most two electrons. The s sublevel has 1 orbital (2 electrons), p has 3 orbitals (6 electrons), d has 5 orbitals (10 electrons), and f has 7 orbitals (14 electrons). Energy level n holds a maximum of 2n² electrons: 2, 8, 18, and 32 for n = 1 to 4. An s orbital is spherical, and p orbitals are dumbbell-shaped along the x, y, and z
Electrons fill the lowest-energy orbitals first Aufbau principle
Pauli exclusion No two electrons in an atom share all four quantum principle numbers
Maximum two electrons per orbital, with opposite spins
Within a sublevel, electrons occupy empty orbitals singly
Nitrogen's three 2p electrons are unpaired, with parallel spins
Following the Aufbau order, oxygen (8 electrons) is 1s² 2s² 2p⁴, and iron (26 electrons) is 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁶. The 4s sublevel fills before 3d because it is slightly lower in energy for these atoms, but when a transition metal forms an ion, the 4s electrons are removed first: Fe²⁺ is [Ar] 3d⁶. Chromium ([Ar] 4s¹ 3d⁵) and copper ([Ar] 4s¹ 3d¹⁰) are the common exceptions, because half-filled and filled d sublevels are especially stable.
Noble-Gas Notation and Valence Electrons
Noble-gas (shorthand) notation replaces the core electrons with the symbol of the previous noble gas in brackets. Sodium is [Ne] 3s¹, and sulfur is [Ne] 3s² 3p⁴. Valence electrons are those in the outermost principal energy level, and they control bonding and chemical behavior. For main-group elements, the number of valence electrons equals the last digit of the group number: group 1 has 1, group 2 has 2, group 13 has 3, and group 17 has 7. Noble gases have 8 (helium has 2), a full outer s and p set that explains their low reactivity.
Students often write potassium as 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹ because they expect level 3 to fill completely before level 4. A teacher can have students locate potassium on the table first: it sits in period 4 of the s block, so its last electron must be 4s. Using the table as a map of the filling order replaces memorized diagonal charts.
Blocks of the Periodic Table
The periodic table's shape follows from electron configuration. The period (row) number equals the highest principal energy level that is occupied. The table divides into blocks named for the sublevel being filled: the s block is groups 1 and 2 plus helium, the p block is groups 13 to 18, the d block is the transition metals in groups 3 to 12, and the f block is the lanthanides and actinides. Elements in the same group share a valence configuration, which is why they share chemical properties: lithium, sodium, and potassium all end in ns¹ and all react vigorously with water.
By Hund's rule the three 3p electrons occupy separate orbitals, so there are 3 unpaired electrons.
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