Difficulty: Introductory | Prerequisites: None (this is foundational material for general chemistry).
This is the starting point for all of chemistry. Atomic structure explains what matter is made of at its smallest scale, how elements differ from one another, and why atoms behave the way they do. If you are coming in cold, start here: everything from bonding to reactions builds on these ideas. You should be comfortable with basic algebra and scientific notation before diving in.
Atoms are built from three subatomic particles: protons (positive, in the nucleus), neutrons (neutral, in the nucleus), and electrons (negative, orbiting outside). The number of protons defines the element; the number of neutrons defines the isotope. The Bohr model describes electrons sitting in quantised energy levels around the nucleus, absorbing or emitting photons when they jump between levels, which is the basis for atomic spectra.
Proton
A positively charged subatomic particle found in the nucleus, with a mass of approximately 1 atomic mass unit (amu). The number of protons defines which element an atom is.
In simple terms, protons are what give each element its identity. Hydrogen has one, helium has two, and so on.
Neutron
An uncharged subatomic particle found in the nucleus, with a mass slightly larger than that of a proton.
Think of neutrons as the glue that helps hold the nucleus together. They add mass but no charge.
Electron
A negatively charged particle that occupies energy levels (shells) around the nucleus. Its mass is roughly 1/2000 of a proton's mass.
In simple terms, electrons are the tiny, fast-moving particles on the outside of the atom that determine how it bonds and reacts.
Atomic number (Z)
The number of protons in the nucleus of an atom. This number uniquely identifies the element.
Think of it as an element's ID number on the periodic table.
Mass number (A)
The total number of protons plus neutrons in the nucleus (A = p + n).
In simple terms, mass number tells you how heavy the nucleus is in whole-number terms.
Isotope
Atoms of the same element (same atomic number) that have different mass numbers because they contain different numbers of neutrons.
Think of isotopes as siblings: same element, same number of protons, but different weights because the neutron count varies.
Atomic mass
The mass of a single atom, nearly equal to its mass number (p + n), measured in atomic mass units (amu).
In simple terms, it is the actual mass of one particular atom.
Atomic weight
The weighted average of the atomic masses of all naturally occurring isotopes of an element, measured in amu.
Think of it as the "average weight" of an element's atoms as they exist in nature. This is the number you see on the periodic table.
Avogadro's number
6.02 x 10^23. One mole of any substance contains this many particles (atoms, ions, or molecules).
In simple terms, a mole is a chemist's counting unit, the way a dozen means 12. A mole means 6.02 x 10^23.
Quantum (plural: quanta)
A discrete bundle of energy. Energy emitted as electromagnetic radiation comes in these fixed packets, not as a continuous stream.
Think of it as energy delivered in parcels rather than poured from a tap.
Planck relation
The equation E = hf, which relates the energy of a photon to its frequency. h is Planck's constant (6.63 x 10^-34 J·s).
In simple terms, higher frequency means more energy per photon.
Bohr model
A model of the atom in which electrons orbit the nucleus at fixed energy levels. When an electron moves between levels, it absorbs or emits a photon whose energy matches the gap.
Think of it as a ladder: electrons sit on rungs, and they can only jump whole rungs, never hover between them.
Ground state
The lowest energy state of an electron in an atom (n = 1 for hydrogen).
In simple terms, ground state is where the electron "wants" to be when nothing is pushing it up.
Emission spectrum
The set of specific wavelengths (colours) of light that an element emits when its excited electrons drop back to lower energy levels.
Think of it as an element's unique light fingerprint.
Absorption spectrum
The set of specific wavelengths of light that an element absorbs when its electrons are promoted to higher energy levels. It appears as dark lines against a continuous spectrum.
In simple terms, the absorption spectrum is the reverse of the emission spectrum: where emission shows bright lines, absorption shows dark ones.
Lyman series
Hydrogen emission lines produced when electrons transition from higher energy levels down to n = 1. These fall in the ultraviolet range.
Balmer series
Hydrogen emission lines produced when electrons transition from n ≥ 3 down to n = 2. These fall in the visible range.
Paschen series
Hydrogen emission lines produced when electrons transition from n ≥ 4 down to n = 3. These fall in the infrared range.
Located in the nucleus; carry a positive charge (+1)
Mass of approximately 1 amu
The number of protons equals the atomic number (Z), which defines the element
Located in the nucleus; carry no charge
Mass is slightly larger than a proton (roughly 1 amu)
Adding protons and neutrons gives the mass number: A = p + n
Isotopes arise when atoms of the same element differ in their neutron count, producing different mass numbers
Written as the element symbol with mass number (A) as a superscript and atomic number (Z) as a subscript
Example: Helium-4 is written with A = 4 and Z = 2
Found outside the nucleus, distributed across energy levels (shells)
Mass is roughly 1/2000 of a proton's mass, so electrons contribute negligibly to the atom's total mass
Electrons closer to the nucleus are at lower energy; those farther away are at higher energy and experience less nuclear pull
Valence electrons (the outermost electrons) determine an atom's chemical reactivity
The mass of a single atom, nearly equal to its mass number (p + n)
Because isotopes have different neutron counts, different atoms of the same element can have different atomic masses
The weighted average of the atomic masses of all naturally occurring isotopes of an element
This is the number reported on the periodic table
Example: carbon has an atomic weight of 12 amu. One mole (6.02 x 10^23 atoms) of carbon has a combined mass of 12 grams
Avogadro's number = 6.02 x 10^23
A mole is simply this many particles (atoms, ions, or molecules)
It bridges the gap between individual atoms (too small to weigh) and laboratory-scale quantities you can measure on a balance
Energy emitted as electromagnetic radiation comes in discrete bundles called quanta, not as a continuous flow
The energy of a single quantum (photon) is given by the Planck relation: E = hf = hc/λ
h = 6.63 x 10^-34 J·s (Planck's constant)
f = frequency of the photon
c = 3 x 10^8 m/s (speed of light in a vacuum)
λ = wavelength of the photon
Key relationships to commit to memory:
Increasing frequency increases energy
Increasing wavelength decreases energy
Frequency and wavelength are inversely related (c = fλ)
Bohr proposed that electrons orbit the nucleus at fixed distances in quantised energy levels
This model is historically important but ultimately incorrect: electrons do not follow neat circular paths (the quantum mechanical model replaced it)
Angular momentum of an electron in orbit: L = mvr = nh/2π
As the quantum number n increases, angular momentum increases
Energy of an electron in a Bohr orbit: E = -R_H / n²
R_H (Rydberg constant) = 2.179 x 10^-18 J/atom = 13.6 eV/atom = 1312 kJ/mol
n = 1, 2, 3, ... (the principal quantum number)
As n increases, the electron's energy increases (becomes less negative) and the electron sits farther from the nucleus
Ground state: the lowest energy level (n = 1 for hydrogen)
When an electron is excited, it jumps to a higher energy level
Most atoms sit in the ground state at room temperature
When heated, electrons absorb energy, jump to higher levels, then fall back down, emitting photons in the process
The AHED mnemonic describes what happens when an electron moves to a higher level:
Absorbs light
Higher potential energy
Excited state
Distant from the nucleus
Combining Bohr's model with Planck's equation gives: E = hc/λ = -R_H (1/n_i² - 1/n_f²)
Negative energy result: the electron went up in energy level and absorbed light
Positive energy result: the electron went down in energy level and emitted a photon
Lyman series: transitions down to n = 1 (ultraviolet)
Balmer series: transitions from n ≥ 3 down to n = 2 (visible light)
Paschen series: transitions from n ≥ 4 down to n = 3 (infrared)
Every element has a unique absorption spectrum, which acts as a fingerprint for identifying that element
E = hf = \frac{hc}{\lambda}where h = 6.63 x 10^-34 J·s, f = frequency, c = 3 x 10^8 m/s, λ = wavelength.
E = -\frac{R_H}{n^2}where R_H = 2.179 x 10^-18 J/atom, n = principal quantum number (1, 2, 3, ...).
L = mvr = \frac{nh}{2\pi}E = \frac{hc}{\lambda} = -R_H \left( \frac{1}{n_i^2} - \frac{1}{n_f^2} \right)If E is negative, the electron moved to a higher level (absorbed light). If E is positive, the electron dropped to a lower level (emitted a photon).
A = p + n\text{Max electrons} = 2n^2The element symbol is written with the mass number (A) as a left superscript and the atomic number (Z) as a left subscript. For helium-4: A = 4, Z = 2.
Atomic emission and absorption spectra are used in astronomy to identify the composition of distant stars, since each element produces a unique spectral fingerprint. Isotopes are the basis of radiocarbon dating (carbon-14) and medical imaging (technetium-99m in PET scans), where different neutron counts give an element different physical behaviours while keeping its chemistry the same.
Students often think atomic mass and atomic weight are the same thing. They are not. Atomic mass is the mass of a single atom; atomic weight is the weighted average across all naturally occurring isotopes.
Students often think the Bohr model is the current accepted model of the atom. It is not. The quantum mechanical model replaced it. The Bohr model is a useful stepping stone but incorrectly depicts electrons in fixed circular orbits.
Students often confuse mass number with atomic weight. Mass number (A) is a whole number count of protons plus neutrons in one specific atom. Atomic weight is a decimal average across isotopes.
Students often think that when wavelength increases, energy increases. It does not. Energy and wavelength are inversely related: longer wavelength means lower energy.
⚠️ You will almost certainly be asked to calculate the energy of a photon given its wavelength or frequency using E = hf = hc/λ. Know the constants by heart.
⚠️ Expect a question on the relationship between wavelength, frequency, and energy. Remember: high frequency = short wavelength = high energy.
⚠️ Be ready to use the Bohr-Planck equation to calculate the energy emitted or absorbed during an electron transition: E = -R_H (1/n_i² - 1/n_f²). Know the sign convention: negative = absorption, positive = emission.
⚠️ Know the three hydrogen spectral series by name and which n level they transition to: Lyman (n = 1), Balmer (n = 2), Paschen (n = 3).
⚠️ Be able to distinguish between atomic mass (single atom) and atomic weight (weighted average of isotopes). This is a common exam trap.
True or false: Neutrons have a positive charge. (False. Neutrons are uncharged.)
Fill in the blank: The atomic number (Z) equals the number of ______ in the nucleus. (Protons.)
True or false: As wavelength increases, the energy of a photon increases. (False. Energy decreases as wavelength increases.)
Fill in the blank: The Lyman series corresponds to electron transitions down to n = ______. (1.)
True or false: Atomic weight and mass number are the same thing. (False. Mass number is for a single atom; atomic weight is a weighted average of isotopes.)
Q: Calculate the energy of a photon with a wavelength of 500 nm.
A: E = hc/λ = (6.63 x 10^-34 J·s)(3 x 10^8 m/s) / (500 x 10^-9 m) = 3.98 x 10^-19 J.
Q: What is the difference between atomic mass and atomic weight?
A: Atomic mass is the mass of a single atom, roughly equal to its mass number (protons + neutrons). Atomic weight is the weighted average of the masses of all naturally occurring isotopes of that element.
Q: An electron in a hydrogen atom transitions from n = 3 to n = 1. Is energy absorbed or emitted? Which spectral series does this belong to?
A: Energy is emitted (the electron drops to a lower level, releasing a photon). This transition belongs to the Lyman series (transitions to n = 1).
Q: Two isotopes of carbon are carbon-12 and carbon-14. How do they differ?
A: Both have 6 protons (atomic number 6), but carbon-12 has 6 neutrons and carbon-14 has 8 neutrons, giving them different mass numbers.
Q: Using the Bohr model, what happens to an electron's energy as n increases?
A: The energy increases (becomes less negative), and the electron moves farther from the nucleus.
This material connects directly to quantum numbers and electron configuration (covered in Part 2 of these notes), because the Bohr model's energy levels are the foundation for understanding how electrons fill orbitals. It also connects to periodic trends: the number of valence electrons determines an element's group on the periodic table and its bonding behaviour. Finally, emission and absorption spectra link to spectroscopy techniques used throughout analytical chemistry and physics.
Atomic structure, subatomic particles, protons, neutrons, electrons, atomic number, mass number, isotopes, atomic mass, atomic weight, Avogadro's number, mole, quanta, quantum, photon energy, Planck constant, Planck relation, E = hf, Bohr model, Bohr atom, energy levels, ground state, excited state, emission spectrum, absorption spectrum, spectral lines, Lyman series, Balmer series, Paschen series, Rydberg constant, hydrogen atom, electromagnetic radiation, wavelength, frequency, CHEM 101, general chemistry, introductory chemistry, MSU chemistry.