Trace how the atomic model changed with evidence and write electron configurations

Learning goal: Explain what evidence drove each change in the atomic model from Dalton to the quantum mechanical model, use emission and absorption spectra as evidence that electrons occupy discrete energy levels, and write orbital diagrams and extended and noble-gas electron configurations for any element in the first three periods.

Models are built from evidence, then replaced. An atomic model is a description of the atom that explains the experiments done so far and predicts new ones. When a new experiment produces results the model cannot explain, the model is revised or replaced. The history of the atom is therefore a history of experiments, and each model below is tied to the evidence that produced it. Physical Science already gave you protons, neutrons and electrons, atomic number, mass number, ions and isotopes. This course asks a sharper question: how do we know, and where exactly are the electrons?

Dalton: indivisible spheres. In the early 1800s John Dalton proposed that each element consists of tiny indivisible spheres, all identical for that element, which combine in whole-number ratios to form compounds. His evidence was chemical: masses of elements combine in fixed proportions. The model says nothing about the inside of the atom, and that is not a failure. Thinking of atoms as hard spheres still explains many physical properties, such as why a substance is a solid, liquid or gas at room temperature, and chemists still use it whenever the inside of the atom does not matter.

Thomson: the electron and the plum pudding. Late in the 1800s J. J. Thomson studied electrical discharges in cathode-ray tubes, glass tubes with most of the air pumped out and a high voltage across them. The glowing ray inside was bent by electric and magnetic fields in the direction expected for negative particles, and the bending was the same whatever gas or metal was used. So atoms of every element contain the same negative particle, later named the electron, and an atom is divisible. Because atoms are neutral overall, Thomson pictured electrons dotted through a spread-out sphere of positive charge, like fruit in a pudding.

Rutherford: a tiny, dense, positive nucleus. In the years up to 1911, when he published the nuclear model, Ernest Rutherford's team fired alpha particles, small fast positive particles, at a very thin gold foil. If the plum pudding model were right, the spread-out positive charge could only nudge the alpha particles slightly. Almost all of them did pass straight through, but a small fraction were deflected through large angles and a few bounced nearly straight back. The only explanation was that the positive charge and almost all of the mass are concentrated in a very small nucleus, and most of the atom is empty space. Rutherford's model put the nucleus at the center with the electrons somewhere outside.

Bohr: electrons on allowed orbits. A nucleus with orbiting electrons had a problem: classical physics said an orbiting charge should lose energy and spiral in. Niels Bohr used a different kind of evidence, atomic spectra. When a gas is excited by electricity and its light is spread out by a prism or grating, you do not see a continuous rainbow. You see a few sharp colored lines, an emission spectrum, and each element has its own set. Bohr proposed that electrons can occupy only certain allowed orbits, each with a fixed energy, like planets at fixed distances from the sun. An atom is normally in its ground state, with its electrons in the lowest available orbits. If an electron absorbs a photon whose energy exactly matches the gap to a higher orbit, the atom becomes excited; a photon with any other energy is not absorbed for a jump between levels (a photon with enough energy can instead remove the electron entirely, which is ionization). The excited state lasts a tiny fraction of a second, then the electron drops back and emits a photon whose energy equals the difference between the two levels. Because photon energy sets the frequency, and frequency sets the color, each allowed jump gives one line. The absorption spectrum, dark lines where a continuous light passes through the cool gas, shows the same gaps. That each element has its own lines means each has its own set of energy levels, which is why a flame test or a spectral tube identifies an element.

Schrödinger: orbitals and probability. Bohr's model works well for hydrogen and explains periodic trends, but it fails in detail for atoms with many electrons. In the 1920s Erwin Schrödinger treated the electron as a wave and produced the quantum mechanical model. In it you cannot say where an electron is or how it moves; you can only describe a region of space around the nucleus where there is a high probability of finding it, called an orbital or electron cloud. This is the currently accepted model. Its mathematics is reserved for later study; what matters here is its picture of energy levels and orbitals, which is the basis of electron configurations. Notice the shape of the story: the model changed each time because of new evidence, and a technology, high-voltage electricity, opened each new door. Older models were not thrown away; they are still used where they are good enough.

Levels, sublevels and orbitals. In the quantum mechanical model, energy levels are numbered 1, 2, 3 and so on outward from the nucleus. Each level is divided into sublevels labeled s, p, d and f. Level 1 has only an s sublevel; level 2 has s and p; level 3 has s, p and d. Each sublevel has a characteristic number of orbitals, and each orbital holds at most two electrons: an s sublevel has one orbital (2 electrons), a p sublevel has three orbitals (6 electrons) and a d sublevel has five orbitals (10 electrons). The s orbitals are spherical; each p orbital has two lobes along one axis.

Three rules for filling. Electrons in a ground-state atom occupy the lowest available energies, filling in the order 1s, 2s, 2p, 3s, 3p for the first three periods, which is as far as this course goes. (The usual filling guide continues 4s and then 3d, but that is a simplified sequence with exceptions, because the relative energies of 4s and 3d depend on the atom and on how many electrons it has.) An orbital holds at most two electrons, and they must have opposite spins, drawn as up and down arrows. Within a sublevel, electrons enter empty orbitals singly, all with the same spin, before any orbital gets a second electron. An orbital diagram shows a box or line for each orbital with its arrows; an electron configuration lists sublevels with superscripts giving the number of electrons in each, so the superscripts add up to the atomic number for a neutral atom.

A worked model: phosphorus, atomic number 15. Fill in order: 1s holds 2 (total 2), 2s holds 2 (total 4), 2p holds 6 (total 10), 3s holds 2 (total 12), and the remaining 3 go into 3p (total 15). Extended notation: 1s²2s²2p⁶3s²3p³. The orbital diagram puts one up-arrow in each of the three 3p boxes, not two in the first, because of the rule about filling singly first. Noble gas notation replaces the filled inner core with the previous noble gas in brackets: neon is 1s²2s²2p⁶, so phosphorus is [Ne]3s²3p³. The valence electrons are the electrons in the highest occupied level, here the five electrons in 3s and 3p, and they are the ones that decide how phosphorus bonds. The same method gives sodium [Ne]3s¹, chlorine [Ne]3s²3p⁵ and argon [Ne]3s²3p⁶, a full level 3s and 3p, which is why argon is unreactive.

Ions. A sodium atom that loses its single 3s electron becomes Na⁺ with the configuration of neon; a chlorine atom that gains one electron becomes Cl⁻ with the configuration of argon. A configuration therefore explains why sodium forms a 1+ ion and not 2+ in ordinary chemistry, and why chlorine forms 1−. Group 2 metals lose two s electrons, and their reactivity increases down the group because the outer electrons are farther from the nucleus and more easily removed.

Isotopes and mass. Most elements are mixtures of isotopes with the same atomic number but different mass numbers. The atomic mass on the periodic table is a weighted average. For chlorine, 75.76 percent of atoms have mass 34.97 u (atomic mass units) and 24.24 percent have mass 36.97 u, so the average is 0.7576 × 34.97 + 0.2424 × 36.97 = 26.49 + 8.96 = 35.45 u. The mass sits closer to 35 than to 37, which tells you before any calculation that chlorine-35 is the more abundant isotope.

Investigation: see the lines. Materials: a diffraction grating or a plastic spectroscope card, and several light sources: an incandescent or halogen bulb, a compact fluorescent lamp, a white LED and, if one is nearby at night, a neon shop sign or an orange sodium street lamp. Steps: look at each source through the grating from a few meters away, never at the sun, and sketch what you see beside each source's name. Record whether the spread of colors is continuous or broken into separate bright lines, and the colors and rough positions of any lines. A result that supports the claim: the hot filament gives a continuous band, while the fluorescent lamp and the gas signs show separate bright lines at fixed positions that differ from source to source. What the model cannot show: the grating spreads out light already emitted; it cannot show the electron jumps themselves, and it cannot tell you the energy of a level, only differences between levels. Flame tests with metal salts show the same idea but involve an open flame and must be done only by or with a supervising adult.

Beyond this course. Quantum numbers, the equations of de Broglie, Schrödinger and Planck, and configurations of elements beyond the third period are reserved for later study. You do need to be aware that the quantum mechanical model is the accepted one and why.

Worked example 1 — which model explains which result? A student lists three observations: (a) the same negative particle comes out of every element in a cathode-ray tube; (b) a few alpha particles bounce back from gold foil; (c) hydrogen gas glows with four visible lines, not a rainbow. Match each to the model it forced. (a) shows atoms contain electrons and are divisible, which ended Dalton's indivisible sphere and led to Thomson's plum pudding. (b) cannot happen if positive charge is spread thinly; it needs a small dense positive nucleus, Rutherford's model. (c) needs allowed energy levels with fixed gaps, Bohr's model, since a continuous range of electron energies would give a continuous spectrum. None of the three requires the quantum mechanical model, which was adopted because Bohr's orbits fail for atoms beyond hydrogen.

Worked example 2 — configuration and orbital diagram for sulfur, Z = 16. Fill: 1s² (2), 2s² (4), 2p⁶ (10), 3s² (12), 3p⁴ (16). Extended: 1s²2s²2p⁶3s²3p⁴; noble gas: [Ne]3s²3p⁴. Orbital diagram for 3p: three boxes; place one up-arrow in each first (three electrons), then the fourth electron pairs in the first box with a down-arrow. So two boxes hold single electrons and one holds a pair. Valence electrons: 2 + 4 = 6. Sulfur gains two electrons to reach [Ar], which is why its common ion is S²⁻.

Worked example 3 — from the table to the configuration. Aluminum is in period 3 and is the first element of the p-block in that period. Period 3 means the outer electrons are in level 3; being in the first p-block column means one electron in 3p. So aluminum is [Ne]3s²3p¹, and its atomic number must be 10 (neon) + 3 = 13, which matches. Reading the table this way works for any element in the first three periods: the period number gives the level, the block gives s or p, and the position within the block gives the superscript.

Worked example 4 — an excited atom. A sodium atom absorbs a photon and its 3s electron moves to a 3p orbital. Its configuration is now [Ne]3p¹, an excited state. Within a tiny fraction of a second the electron returns to 3s and emits a photon whose energy equals the 3p to 3s gap; for sodium that photon is yellow-orange, the color of a sodium street lamp. A photon with slightly more or less energy than the gap would not have been absorbed at all; that is what discrete means.

Representation task: a timeline with evidence. Make a timeline with five entries: Dalton, Thomson, Rutherford, Bohr, Schrödinger. For each, write the experiment or evidence in one line, a sketch of the model, and one sentence saying what the previous model could not explain. Then write electron configurations in both extended and noble-gas notation, plus an orbital diagram, for magnesium, silicon and chlorine, and circle the valence electrons in each.

Independent performance: spectra report and configuration set. Carry out the diffraction-grating investigation with at least three light sources and write a one-page report. Success criteria: a sketch for each source that says whether the spectrum is continuous or a line spectrum; an explanation, using the words ground state, excited state, photon and energy level, of why a gas gives lines and a hot filament does not; a statement of what the observation cannot tell you; and a table of configurations for all eight elements of period 3 (sodium to argon) in noble-gas notation with the number of valence electrons. A reviewer should be able to check every superscript against the atomic number.