Difficulty: Intermediate | Prerequisites: Matter and Properties (Ch. 1–2)
This block of material answers the question "What is everything made of at the smallest scale?" You move from Dalton's idea that matter is made of indivisible atoms, through Rutherford's discovery of the nucleus, to the modern quantum mechanical model with orbitals and electron configurations. Understanding atomic structure is essential for the periodic table, bonding and every reaction you will study later. If the periodic table makes no sense to you yet, this is the chapter to revisit.
Atoms consist of protons and neutrons in a dense nucleus, with electrons occupying orbitals around it. The atomic number (protons) defines the element; the mass number (protons + neutrons) identifies the isotope. Electrons fill orbitals from lowest to highest energy (aufbau principle), and their arrangement determines an element's chemical behaviour. Light is electromagnetic radiation, and its interaction with atoms produces emission spectra that reveal electron energy levels.
Atom
The smallest particle of an element that retains the chemical properties of that element.
Proton
A positively charged subatomic particle found in the nucleus. Its number defines which element the atom is.
Neutron
A subatomic particle with no charge, found in the nucleus. Its mass is nearly equal to that of a proton.
Electron
A negatively charged subatomic particle found outside the nucleus. It is far lighter than protons or neutrons and determines chemical behaviour.
Nucleus
The small, dense, positively charged central core of the atom, composed of protons and neutrons.
Atomic number (Z)
The number of protons in the nucleus of an atom. It uniquely identifies the element.
Mass number (A)
The total number of protons and neutrons in the nucleus. In the notation krypton-84, the 84 is the mass number.
Isotope
Atoms of the same element that have different numbers of neutrons (and therefore different mass numbers). They share the same atomic number.
Atomic mass unit (amu)
The unit used to express weighted average atomic mass. One amu is defined as 1/12 the mass of a carbon-12 atom.
Bohr model
An early model of the atom in which electrons orbit the nucleus in fixed energy levels. Each orbit corresponds to a specific, quantised energy.
Quantum mechanical model
The modern model of the atom, developed by Schrodinger, in which electrons occupy orbitals (probability regions) rather than fixed orbits.
Orbital
A three-dimensional region around the nucleus where an electron is most likely to be found. Orbitals have characteristic shapes: s (sphere), p (dumbbell), d (cloverleaf), f (complex).
Energy sublevel
A subdivision of a principal energy level, labelled s, p, d or f. The sublevel determines the shape of the orbital.
Aufbau principle
Electrons fill orbitals starting from the lowest available energy level. Think of it as building up the electron configuration from the ground floor.
Pauli exclusion principle
No more than two electrons may occupy a single orbital, and they must have opposite spins.
Hund's rule
When filling orbitals of equal energy, electrons occupy them singly (with parallel spins) before pairing up.
Electron configuration
The arrangement of electrons in an atom's orbitals, written as a sequence like 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹.
Electromagnetic spectrum
The full range of electromagnetic radiation, from radio waves (long wavelength, low frequency) to gamma rays (short wavelength, high frequency). Visible light is a narrow band in the middle.
Wavelength (λ)
The distance between successive crests of a wave. Measured in metres (or nanometres for light).
Frequency (ν)
The number of wave crests passing a point per second. Measured in hertz (Hz). Frequency and wavelength are inversely proportional.
Emission spectrum
The set of specific wavelengths of light emitted by an element when its electrons drop from higher to lower energy levels. Each element has a unique emission spectrum.
Heisenberg uncertainty principle
It is impossible to know both the exact position and the exact velocity of a moving particle at the same time.
Dalton proposed that all elements are composed of atoms, atoms of the same element are identical, atoms combine in simple whole-number ratios, and atoms are neither created nor destroyed in chemical reactions. The statement "atoms are always in motion" is NOT part of Dalton's theory (it belongs to kinetic molecular theory).
Three subatomic particles matter here:
Protons: positive charge, found in the nucleus, mass ≈ 1 amu.
Neutrons: no charge, found in the nucleus, mass ≈ 1 amu (nearly equal to a proton).
Electrons: negative charge, found outside the nucleus, mass ≈ 1/1836 amu (negligible compared to protons/neutrons).
The mass of a neutron nearly equals the mass of a proton. Electrons are the lightest subatomic particle, not the heaviest.
Rutherford's gold foil experiment revealed that the atom is mostly empty space, with a small, dense, positively charged nucleus at the centre. Electrons are distributed around the nucleus and occupy almost all the volume of the atom. The nucleus is composed of protons and neutrons (not electrons).
Atomic number (Z) = number of protons. This defines the element. All atoms of the same element have the same number of protons.
In a neutral atom, protons = electrons. An element with atomic number 76 has 76 protons and 76 electrons.
Mass number (A) = protons + neutrons. In "krypton-84," the 84 is the mass number.
Number of neutrons = mass number – atomic number.
Isotopes are atoms of the same element with different numbers of neutrons. They have the same atomic number but different mass numbers.
To identify isotopes, look for symbols with the same atomic number (same bottom number) but different mass numbers (different top numbers).
Example: ¹⁹F has 9 protons and 10 neutrons (19 – 9 = 10), not 0 neutrons. ²³⁸U has 92 protons and 146 neutrons (238 – 92 = 146).
The atomic mass listed on the periodic table is a weighted average of all naturally occurring isotopes, measured in amu. It accounts for both the mass and the relative abundance of each isotope.
In the Bohr model, electrons move in fixed circular orbits around the nucleus. Each orbit has a fixed energy. When an electron moves closer to the nucleus, its energy decreases. When it moves further away, its energy increases.
Erwin Schrodinger developed the quantum mechanical model, which replaces Bohr's fixed orbits with orbitals (probability clouds). Key points:
Principal energy level (n): n = 1, 2, 3, etc. The number of sublevels in a principal energy level equals n. So the second principal energy level (n = 2) has 2 sublevels: s and p.
Sublevels and orbital shapes:
s orbitals are spherical. 1 orbital per sublevel, holds 2 electrons.
p orbitals are dumbbell-shaped. 3 orbitals per sublevel, holds 6 electrons.
d orbitals have cloverleaf shapes. 5 orbitals per sublevel, holds 10 electrons.
f orbitals have complex shapes. 7 orbitals per sublevel, holds 14 electrons.
The shape of an electron cloud is determined by the energy sublevel (s, p, d, f). The size is determined by the principal quantum number (n).
In the notation 4p³, the "4" is the principal energy level, the "p" indicates the orbital shape, and the "3" is the number of electrons in that sublevel.
Aufbau principle: electrons enter orbitals of lowest energy first. The filling order is 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p.
Pauli exclusion principle: a maximum of two electrons per orbital, with opposite spins.
Hund's rule: electrons fill orbitals of equal energy singly before pairing.
Potassium (K, Z = 19): 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹. The next orbital after 3p in the filling order is 4s, not 3d.
Sulfur (S, Z = 16): 1s² 2s² 2p⁶ 3s² 3p⁴. The 3p sublevel has 3 orbitals; with 4 electrons, two orbitals are singly occupied and one is paired, giving 2 unpaired electrons.
All electromagnetic radiation travels at the speed of light (c ≈ 3.00 x 10⁸ m/s) in a vacuum. Visible light and gamma rays travel at the same speed.
Wavelength and frequency are inversely proportional: c = λν. Higher frequency means shorter wavelength.
The EM spectrum from longest wavelength to shortest: radio waves, microwaves, infrared, visible light, ultraviolet, X-rays, gamma rays.
Gamma rays have the highest frequencies and shortest wavelengths.
Within visible light, violet has the shortest wavelength and red has the longest.
A wavelength of 10⁻⁷ m falls in the visible light / near-UV range.
When an electric discharge passes through a gas (e.g. sodium vapour), the atoms emit light at specific wavelengths. This produces an emission spectrum (a set of discrete coloured lines), not a continuous spectrum or white light.
If you know the position of a moving particle precisely, you cannot simultaneously know its velocity precisely (and vice versa). The quantities that trade off are position and velocity (not mass, charge or spin).
Number of neutrons = mass number – atomic number (A – Z)
Speed of light: c = λν (wavelength x frequency)
Wavelength and frequency are inversely proportional: as one increases, the other decreases.
Maximum electrons per sublevel: s = 2, p = 6, d = 10, f = 14
Maximum electrons per energy level n: 2n²
Filling order: 1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → 5s → 4d → 5p → 6s → 4f → 5d → 6p → 7s → 5f → 6d → 7p
Students often think "atoms are always in motion" is part of Dalton's theory. It is not; that idea comes from kinetic molecular theory.
Electrons are NOT the heaviest subatomic particle. They are the lightest by a large margin.
The nucleus does not contain electrons. It contains only protons and neutrons.
The number after an element name (e.g. krypton-84) is the mass number, not the atomic number.
The next orbital after 3p is 4s, not 3d. This catches students who assume the principal quantum number always increases sequentially.
⚠️ Know how to calculate the number of neutrons from mass number and atomic number.
⚠️ Be able to write electron configurations and identify unpaired electrons.
⚠️ The filling order (especially 4s before 3d) is heavily tested.
⚠️ Visible light, gamma rays and all EM radiation travel at the same speed in a vacuum.
⚠️ Violet has the shortest wavelength of visible light; gamma rays have the highest frequency of all EM radiation.
⚠️ Frequency and wavelength are inversely proportional. This relationship comes up repeatedly.
⚠️ Heisenberg: position and velocity are the pair you cannot know simultaneously.
True or False: The nucleus of an atom contains protons, neutrons and electrons.
Fill in the blank: The number of protons in an atom is called the ______ number.
True or False: Isotopes of the same element have different numbers of protons.
Fill in the blank: In the Bohr model, an electron in an orbit has a fixed ______.
True or False: Gamma rays travel faster than visible light in a vacuum.
Q: Which of the following is NOT part of Dalton's atomic theory: (a) all elements are composed of atoms, (b) atoms are always in motion, (c) atoms of the same element are identical, (d) atoms combine in whole-number ratios?
A: (b) Atoms are always in motion. This belongs to kinetic molecular theory.
Q: The particles found in the nucleus of an atom are ______.
A: Protons and neutrons.
Q: An element has atomic number 76. How many protons and electrons does a neutral atom have?
A: 76 protons and 76 electrons.
Q: What is the next atomic orbital in the filling series 1s, 2s, 2p, 3s, 3p?
A: 4s (not 3d).
Q: How many unpaired electrons does a sulfur atom (Z = 16) have?
A: 2. Sulfur's configuration ends 3p⁴: three p orbitals hold 4 electrons as ↑↓ ↑ ↑, leaving 2 unpaired.
Q: Which colour of visible light has the shortest wavelength?
A: Violet.
Q: According to the Heisenberg uncertainty principle, if you know the position of a particle, which quantity cannot be known?
A: Its velocity.
Electron configurations link directly to the periodic table (Ch. 6): the table's shape reflects the filling order of s, p, d and f orbitals. Valence electrons, which you will study next, determine how atoms bond (Ch. 7–8) and explain periodic trends like electronegativity and ionisation energy.
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