Difficulty: Introductory to Intermediate | Prerequisites: Basic understanding of the periodic table layout (groups, periods)
Everything in chemistry comes back to atoms and how their electrons are arranged. This topic covers how to count subatomic particles, how to write electron configurations, and how the periodic table encodes trends in atomic size, ionisation energy, and more. If you can work confidently with electron configurations and periodic trends, you will find bonding, molecular geometry, and even thermochemistry easier to navigate. This material spans roughly the first third of CHM 11500 and shows up in some form on every exam.
Atoms are defined by their protons (atomic number), and ions gain or lose electrons. Electron configurations follow the Aufbau principle. Atomic radius decreases across a period and increases down a group. The electromagnetic spectrum links wavelength, frequency, and photon energy.
Atomic number (Z)
The number of protons in the nucleus of an atom. It defines the element.
In simple terms, the atomic number is the element's identity card.
Mass number (A)
The total number of protons plus neutrons in the nucleus.
In simple terms, mass number = protons + neutrons. It is the superscript in isotope notation.
Ion
An atom or group of atoms with a net electric charge due to the gain or loss of electrons. Cations are positive (lost electrons); anions are negative (gained electrons).
Isoelectronic
Two species are isoelectronic if they have the same number of electrons and therefore the same electron configuration.
In simple terms, isoelectronic species are "electron twins," different elements or ions that happen to share the same electron count.
Electron configuration
The distribution of electrons among the orbitals of an atom, written in order of increasing energy sublevel (1s, 2s, 2p, 3s, 3p, 4s, 3d, etc.).
Core electrons
The inner electrons that are not in the outermost (valence) shell. They match the electron configuration of the nearest noble gas of lower atomic number.
Valence electrons
The electrons in the outermost shell (highest principal quantum number). These are the electrons involved in chemical bonding.
In simple terms, valence electrons are the ones that "do the chemistry."
Electromagnetic radiation
Energy that travels through space as waves with characteristic wavelength (λ) and frequency (ν). The full spectrum runs from long-wavelength radio waves to short-wavelength gamma rays.
Photon
A discrete packet (quantum) of electromagnetic energy. The energy of a single photon is E = hν = hc/λ.
Atomic radius
The distance from the nucleus to the outermost electron shell. It decreases across a period (more protons pulling electrons inward) and increases down a group (additional electron shells).
Effective nuclear charge (Z_eff)
The net positive charge experienced by a valence electron after accounting for the shielding effect of inner electrons.
In simple terms, Z_eff is how strongly the nucleus "pulls" on the outer electrons once you subtract the screening from the core electrons.
For any isotope written as ᴬ_Z X:
Protons = Z (atomic number)
Neutrons = A - Z
Electrons = Z for a neutral atom; subtract the charge for cations, add for anions
Worked example: The iodide anion ¹²⁷₅₃I⁻
Protons = 53
Neutrons = 127 - 53 = 74
Electrons = 53 + 1 = 54 (gained one electron to become I⁻)
To find a species isoelectronic with an ion, look for a neutral atom or another ion with the same electron count.
I⁻ has 54 electrons. The noble gas with 54 electrons is xenon (Xe), so I⁻ is isoelectronic with Xe.
Sr²⁺ has 38 - 2 = 36 electrons. The noble gas with 36 electrons is krypton (Kr), so Sr²⁺ is isoelectronic with Kr.
Follow the Aufbau order: 1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → 5s → 4d → 5p → 6s → 4f → 5d → 6p → 7s → 5f → 6d → 7p.
Each sublevel holds a maximum number of electrons: s = 2, p = 6, d = 10, f = 14.
Worked example: Neutral sulfur (S), Z = 16
Full configuration: 1s² 2s² 2p⁶ 3s² 3p⁴
That accounts for 2 + 2 + 6 + 2 + 4 = 16 electrons.
Worked example: Strontium ion Sr²⁺, Z = 38, charge = +2, so 36 electrons
Full configuration: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶
This is identical to the electron configuration of krypton, confirming that Sr²⁺ is isoelectronic with Kr.
Magnesium (Mg), Z = 12: configuration is 1s² 2s² 2p⁶ 3s²
Core electrons = those in the 1s, 2s, and 2p sublevels = 2 + 2 + 6 = 10
Valence electrons = those in the outermost shell (n = 3) = 2
The number of valence electrons for main-group elements equals their group number (using the old IUPAC numbering) or can be read from the last shell in the configuration.
Across a period (left to right): Atomic radius decreases. Each element has one more proton and one more electron in the same shell. The increased nuclear charge pulls the electron cloud inward because shielding stays roughly constant.
Down a group (top to bottom): Atomic radius increases. Each row adds a new principal energy level, putting the outermost electrons farther from the nucleus.
Worked example: Rank K, Al, P, N, O from largest to smallest atomic radius.
K is in period 4, group 1: largest radius (furthest left, lowest in the table among this set).
Al is in period 3, group 13.
P is in period 3, group 15 (smaller than Al, same period but further right).
N is in period 2, group 15.
O is in period 2, group 16 (smaller than N, same period but further right).
Order: K > Al > P > N > O.
The electromagnetic spectrum, from lowest to highest energy per photon: radio waves, microwaves, infrared, visible light, ultraviolet, X-rays, gamma rays.
Key relationships:
c = λν (speed of light = wavelength x frequency)
E = hν = hc / λ (energy of a photon)
Where c = 2.998 x 10⁸ m/s and h = 6.626 x 10⁻³⁴ J·s.
Shorter wavelength means higher frequency and higher energy per photon. Gamma rays have the shortest wavelength and the greatest energy per photon.
Calculate the energy of a photon with a wavelength of 517 nm.
Convert nm to m: 517 nm = 517 x 10⁻⁹ m = 5.17 x 10⁻⁷ m
E = hc / λ = (6.626 x 10⁻³⁴ J·s)(2.998 x 10⁸ m/s) / (5.17 x 10⁻⁷ m)
E = 1.987 x 10⁻²⁵ / 5.17 x 10⁻⁷ = 3.84 x 10⁻¹⁹ J
Neutrons = A - Z
Electrons in an ion = Z - charge (positive charge means fewer electrons; negative charge means more)
c = λν
E = hν = hc / λ
c = 2.998 x 10⁸ m/s
h = 6.626 x 10⁻³⁴ J·s
Electron configurations determine how elements bond, which controls everything from drug design to semiconductor behaviour. The photon energy equation is the basis of spectroscopy: chemists shine light of known wavelength at a sample and measure what is absorbed, identifying unknown compounds by their characteristic absorption patterns. This is how astronomers determine the composition of distant stars.
Students often forget to add an electron for anions or subtract one for cations when counting electrons in an ion. I⁻ has 54 electrons, not 53.
Students sometimes confuse mass number with atomic mass. Mass number is an integer (protons + neutrons for one isotope). Atomic mass on the periodic table is a weighted average of all isotopes.
Students commonly assume that atomic radius increases across a period because "more electrons means bigger." The increased nuclear charge more than compensates, pulling the shell inward.
When ranking atomic size, students forget that going down a group adds an entirely new shell, which is a larger effect than moving across a period.
⚠️ Subatomic particle counts for ions (especially the electron count) are a favourite multiple-choice question. Double-check the charge.
⚠️ Electron configuration questions often include distractor options that put the wrong number of electrons in a sublevel (e.g., 3p⁶ for sulfur instead of the correct 3p⁴). Count your electrons against Z.
⚠️ Periodic trend ranking questions require you to know both the across-a-period and down-a-group directions. Draw a quick arrow on your scratch paper if it helps.
⚠️ Gamma rays have the highest energy per photon. This is a straightforward recall question that appears regularly.
True or false: A cation has more electrons than protons. False (a cation has fewer electrons than protons).
Fill in the blank: The electron configuration of neon is ___. 1s² 2s² 2p⁶.
True or false: Atomic radius increases from left to right across a period. False (it decreases).
Fill in the blank: The energy of a photon is inversely proportional to its ___. Wavelength.
True or false: Sr²⁺ is isoelectronic with Xe. False (Sr²⁺ has 36 electrons, the same as Kr, not Xe).
Q: How many protons, electrons, and neutrons are in the iodide ion ¹²⁷₅₃I⁻?
A: 53 protons, 54 electrons, 74 neutrons.
Q: Which noble gas is isoelectronic with I⁻?
A: Xenon (Xe). Both have 54 electrons.
Q: Write the full electron configuration for neutral sulfur (S, Z = 16).
A: 1s² 2s² 2p⁶ 3s² 3p⁴.
Q: Rank these elements from largest to smallest atomic radius: O, N, P, K, Al.
A: K > Al > P > N > O.
Q: A neutral Mg atom has 12 electrons. How many are core electrons and how many are valence electrons?
A: 10 core electrons and 2 valence electrons. The core electrons fill 1s² 2s² 2p⁶; the valence electrons are the 3s².
Q: Calculate the energy of a photon with wavelength 517 nm.
A: E = hc/λ = (6.626 x 10⁻³⁴)(2.998 x 10⁸) / (5.17 x 10⁻⁷) = 3.84 x 10⁻¹⁹ J.
Electron configurations feed directly into chemical bonding: the number of valence electrons determines how many bonds an atom forms and its Lewis structure. Periodic trends in atomic radius connect to trends in ionisation energy and electronegativity, which in turn determine whether a bond is ionic, polar covalent, or nonpolar covalent. Photon energy calculations reappear in topics on atomic emission spectra and the Bohr model.
atomic number, mass number, isotope notation, protons, neutrons, electrons, cation, anion, isoelectronic, electron configuration, Aufbau principle, orbital filling order, core electrons, valence electrons, noble gas core, periodic trends, atomic radius, effective nuclear charge, shielding, electromagnetic spectrum, photon energy, wavelength, frequency, Planck's constant, speed of light, E = hc/λ, gamma rays, UV, IR, CHM 11500, Purdue, general chemistry