Difficulty: Intermediate | Prerequisites: Chapter 4 (atomic structure basics, protons/neutrons/electrons)
Source: Chem 1 Final Study Guide, General Chemistry (Purdue University, Chhabra)
Tags: electromagnetic radiation, wavelength, frequency, quantum, photon, photoelectric effect, atomic emission spectrum, Bohr model, quantum mechanical model, orbitals, electron configuration, aufbau principle, Pauli exclusion, Hund's rule, periodic table, periodic law, atomic radius, ionization energy, electronegativity, electron affinity, alkali metals, halogens, noble gases, transition metals
Once you know atoms have electrons, the next question is: where are those electrons, and how do they behave? Chapter 5 takes you from the wave nature of light through Bohr's hydrogen model to the full quantum mechanical picture with orbitals and electron configurations. Chapter 6 then shows how the periodic table is organised around those configurations, and Chapter 7 walks through each group's distinctive properties. These chapters are the bridge between atomic structure and chemical bonding. If electron configurations do not make sense to you, bonding (Ch. 8-9) and reactions will be significantly harder.
Light behaves as both a wave and a particle. Electrons in atoms occupy orbitals described by quantum numbers, and their arrangement follows the aufbau principle, Pauli exclusion principle, and Hund's rule. The periodic table is organised by electron configuration, and periodic trends (atomic radius, ionisation energy, electronegativity) follow predictable patterns across periods and down groups.
Electromagnetic radiation (EMR)
A form of energy that exhibits wavelike behaviour as it travels through space. Visible light is one type. In simple terms, it is energy travelling as waves.
Wavelength (λ)
The shortest distance between equivalent points on a continuous wave (e.g. crest to crest). Measured in metres or nanometres.
Frequency (ν)
The number of waves that pass a given point per second. Measured in hertz (Hz, or s⁻¹).
Amplitude
The height of a wave from its origin (resting position) to its crest. Related to the intensity or brightness of light.
Speed of light (c)
The product of wavelength and frequency: c = λν. Approximately 3.00 x 10⁸ m/s.
Electromagnetic spectrum
The full range of EMR, from radio waves (long wavelength, low frequency) to gamma rays (short wavelength, high frequency). Visible light sits in a narrow band in the middle.
Quantum
The minimum amount of energy that can be gained or lost by an atom. Proposed by Max Planck.
Photon
A particle of EMR with no mass that carries a quantum of energy. Think of it as a tiny "packet" of light energy.
Photoelectric effect
The emission of electrons from a metal surface when light of a sufficient frequency strikes it. Einstein explained this by proposing light has particle-like properties.
Atomic emission spectrum
The set of specific frequencies of light emitted by atoms of a given element when excited. Each element has a unique emission spectrum, like a fingerprint.
Ground state
The lowest energy state of an atom, where the electron occupies the lowest available energy level.
Excited state
A higher-energy state reached when an atom absorbs energy and an electron moves to a higher energy level.
Bohr model
Niels Bohr's model of the hydrogen atom with electrons in fixed circular orbits at specific energy levels. Correctly predicts hydrogen's emission spectrum but fails for multi-electron atoms.
Balmer series
The four electron transitions that produce visible lines in hydrogen's emission spectrum (n = 6, 5, 4, 3 all dropping to n = 2, producing violet, blue, green, and red light respectively).
de Broglie's hypothesis
Since light (normally a wave) can behave as a particle, electrons (normally particles) can behave as waves. All moving particles have wave characteristics.
Heisenberg uncertainty principle
It is impossible to know both the exact velocity and the exact position of an electron at the same time. The act of measuring one disturbs the other.
Quantum mechanical model
The modern model of the atom in which electrons are treated as waves. It describes the probability of finding an electron in a region of space (an orbital) rather than tracing a definite path.
Principal quantum number (n)
Describes the main energy level of an electron. Any positive integer (1, 2, 3...). The larger the number, the higher the energy and the further from the nucleus.
Angular momentum quantum number (ℓ)
Describes the shape of the orbital. Ranges from 0 to (n – 1). ℓ = 0 is s (spherical), ℓ = 1 is p (dumbbell), ℓ = 2 is d, ℓ = 3 is f.
Magnetic quantum number (mₗ)
Describes the orientation of the orbital in space. Ranges from –ℓ to +ℓ.
Electron configuration
The arrangement of electrons in an atom's orbitals, written using sublevel notation (e.g. 1s² 2s² 2p⁶).
Aufbau principle
Each electron occupies the lowest energy orbital available. In simple terms, fill from the bottom up.
Pauli exclusion principle
A maximum of two electrons may occupy a single orbital, and they must have opposite spins.
Hund's rule
Within a set of equal-energy orbitals, each orbital gets one electron (all with the same spin) before any orbital gets a second. Think of it as "one in each seat before anyone doubles up."
Valence electrons
Electrons in the atom's outermost (highest principal energy level) orbitals. These determine chemical behaviour.
Electron-dot structure (Lewis dot)
A representation showing an element's symbol surrounded by dots for its valence electrons.
Periodic law
The statement that the chemical and physical properties of elements repeat periodically when arranged by increasing atomic number.
Atomic radius
Half the distance between the nuclei of two bonded identical atoms. Decreases across a period (more protons pulling electrons in) and increases down a group (more electron shells).
Ionic radius
The radius of an ion. Cations are smaller than their parent atoms (lost electrons, fewer shells or less shielding). Anions are larger (gained electrons, more repulsion).
Ionisation energy
The energy required to remove an electron from a gaseous atom or ion. Increases across a period and decreases down a group.
Electronegativity
The relative ability of an atom to attract electrons in a chemical bond. Increases across a period and decreases down a group. Fluorine is the most electronegative element.
Electron affinity
The ability of an atom to hold onto electrons (related to the energy change when an electron is added).
Shielding (screening)
Inner electrons partially block the nuclear charge felt by outer electrons. Shielding is constant across a period but increases down a group.
Wave nature of light
EMR travels as waves characterised by wavelength, frequency, and amplitude
c = λν (speed of light = wavelength x frequency)
The electromagnetic spectrum spans radio waves through gamma rays
Particle nature of light
Planck: energy is quantised; E = hν (where h = Planck's constant, 6.626 x 10⁻³⁴ J·s)
Einstein: the photoelectric effect proves light has particle-like behaviour; photons are discrete packets of energy
Light has dual wave-particle nature
Atomic emission spectra
Excited atoms emit light at specific frequencies, producing a line spectrum unique to each element
This can identify unknown elements in a sample
Bohr model (hydrogen)
Electrons occupy fixed circular orbits with quantised energy levels (n = 1, 2, 3...)
Ground state: n = 1 (lowest energy, no radiation emitted)
Excited state: electron absorbs energy and jumps to a higher orbit
When the electron drops back down, it emits a photon of specific frequency
Balmer series (visible): transitions ending at n = 2
n = 6 to 2: violet
n = 5 to 2: blue
n = 4 to 2: green
n = 3 to 2: red
Lyman series (UV): transitions ending at n = 1
Paschen series (IR): transitions ending at n = 3
Works for hydrogen only; breaks down for multi-electron atoms
Quantum mechanical model
de Broglie: electrons have wave properties
Heisenberg: you cannot simultaneously know an electron's position and velocity
Schrödinger: wave equation describes electron behaviour; solutions give orbitals (probability regions)
This model does not attempt to describe the electron's path
Orbitals and quantum numbers
Principal quantum number (n): energy level (1, 2, 3...)
Angular momentum quantum number (ℓ): orbital shape (0 = s, 1 = p, 2 = d, 3 = f)
n = 1 has 1 sublevel (s), n = 2 has 2 (s, p), n = 3 has 3 (s, p, d), etc.
Magnetic quantum number (mₗ): orbital orientation within a sublevel
Each orbital holds a maximum of 2 electrons (opposite spins)
s: 1 orbital (2e⁻), p: 3 orbitals (6e⁻), d: 5 orbitals (10e⁻), f: 7 orbitals (14e⁻)
Electron configurations
Aufbau principle: fill lowest energy orbitals first (follow the aufbau diagram)
Pauli exclusion: max 2 electrons per orbital, opposite spins
Hund's rule: one electron in each equal-energy orbital before pairing
Notable exceptions: chromium is [Ar]4s¹3d⁵ (not 4s²3d⁴) and copper is [Ar]4s¹3d¹⁰ (not 4s²3d⁹), because half-filled and fully filled d subshells are especially stable
Valence electrons
Only electrons in the outermost principal energy level
Example: sulfur has 16 electrons total, but only 6 are valence (in the 3s and 3p orbitals)
History
Dmitri Mendeleev organised elements by increasing atomic mass into columns with similar properties
He left gaps for undiscovered elements
Modern periodic law: properties repeat when elements are arranged by increasing atomic number
Organisation by electron configuration
Elements in the same group share the same number of valence electrons and similar chemical properties
The energy level of valence electrons indicates the period
The group number (for representative elements) equals the number of valence electrons
Blocks of the periodic table
s-block: groups 1A and 2A (valence electrons in s orbitals); 2 groups wide
p-block: groups 3A through 8A (valence electrons in p orbitals); 6 groups wide
d-block: transition metals (filling d orbitals of energy level n – 1); 10 groups wide
f-block: inner transition metals (lanthanides and actinides, filling 4f and 5f orbitals); 14 columns wide
Periodic trends
Atomic radius: decreases left to right across a period (more protons, same shielding); increases top to bottom down a group (more shells)
Ionisation energy: increases across a period; decreases down a group
Electronegativity: increases across a period; decreases down a group (fluorine is highest)
Electron affinity: generally increases across a period
Shielding: constant across a period; increases down a group
Group 1A: Alkali metals
Highly reactive, 1 valence electron
Reactivity increases down the group (Rb, Cs, Fr most reactive; Li least reactive)
Na and K are the most abundant
Li has a diagonal relationship with Mg
Group 2A: Alkaline earth metals
Harder and less reactive than alkali metals
Be has a diagonal relationship with Al
Ca is widely found in nature; Mg is abundant and easily shaped
Group 3A: Boron group
Found as oxides in Earth's crust
Boron is a metalloid; aluminium is the most abundant metal (3rd most abundant element in the crust)
Gallium can melt in your hand
Group 4A: Carbon group
Metallic character increases going down the group
Carbon: organic chemistry, has allotropes (diamond, graphite)
Silicon: second most abundant element in the crust, commonly as SiO₂
Group 5A: Nitrogen group
5 valence electrons
Nitrogen: 78% of Earth's atmosphere
Phosphorus: 3 solid allotropes
Group 6A: Oxygen group (chalcogens)
6 valence electrons, tend to form 2- ions
Oxygen: most abundant element in Earth's crust, has 2 allotropes (O₂ and O₃)
Sulfur: 10 allotropes
Group 7A: Halogens
7 valence electrons, form 1- ions
Very reactive nonmetals; always found combined in nature
Fluorine: most electronegative and most reactive element
Chlorine: reacts with nearly all elements
Group 8A: Noble gases
8 valence electrons (except He with 2), extremely stable and unreactive
Among the last naturally occurring elements discovered (colourless, odourless)
He: lightest noble gas; Ar: most abundant noble gas in the atmosphere
Transition metals
Electrical conductivity, lustre, malleability
Can lose s electrons and sometimes d electrons, forming 2+ or 3+ ions
Magnetism: diamagnetic (all electrons paired, slightly repelled by magnetic field), paramagnetic (unpaired electron, attracted), ferromagnetic (strong attraction, ions stay aligned)
Inner transition metals
Lanthanides: silvery, high melting points, hard to separate
Actinides: all radioactive; only 3 exist in nature, the rest are transuranium elements
Speed of light: c = λν (3.00 x 10⁸ m/s)
Energy of a quantum: E = hν (h = 6.626 x 10⁻³⁴ J·s)
Hydrogen visible emission (Balmer series): n = 6→2 (violet), 5→2 (blue), 4→2 (green), 3→2 (red)
Periodic trend arrows:
Atomic radius: ← across period, ↓ down group (increases)
Ionisation energy: → across period, ↑ up group (increases)
Electronegativity: → across period, ↑ up group (increases)
Atomic emission spectra are used in forensic science and astronomy to identify elements at a distance, including the composition of distant stars.
Electron configurations explain why certain elements are metals, nonmetals, or metalloids, and predict how they will bond.
Electronegativity differences determine whether a bond is ionic, polar covalent, or nonpolar covalent, which is fundamental to understanding materials science and drug design.
Students often think the Bohr model applies to all atoms. It works only for hydrogen. Multi-electron atoms require the quantum mechanical model.
A common mistake is writing chromium as [Ar]4s²3d⁴. The correct configuration is [Ar]4s¹3d⁵, because a half-filled d subshell is unusually stable. The same logic applies to copper ([Ar]4s¹3d¹⁰).
Students sometimes confuse "energy level" with "orbital." An energy level (n = 2) contains sublevels (2s, 2p), and each sublevel contains one or more orbitals.
"Valence electrons" does not mean "outermost electrons in any sublevel." For main-group elements, valence electrons are only those in the highest principal energy level.
⚠️ Electron configurations are tested directly and are essential for predicting bonding, periodic trends, and magnetism.
⚠️ Know the exceptions: chromium and copper. These appear on nearly every exam.
⚠️ Be able to identify the block (s, p, d, f) of any element from its position on the periodic table.
⚠️ Periodic trend questions are common. Remember the direction of each trend and be able to explain why (nuclear charge vs shielding vs number of shells).
⚠️ Understand the difference between the Bohr model and the quantum mechanical model, and why the Bohr model is limited.
True or False: An electron in an excited state has more energy than one in the ground state.
Fill in the blank: Hund's rule says that electrons fill equal-energy orbitals __________ before pairing up.
True or False: Ionisation energy decreases as you move across a period from left to right.
Fill in the blank: The angular momentum quantum number ℓ = 1 corresponds to a __________ orbital.
True or False: Noble gases have 8 valence electrons, including helium.
Q: Write the electron configuration of iron (Fe, atomic number 26).
A: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁶. Or in noble gas shorthand: [Ar] 4s² 3d⁶.
Q: Why is the atomic radius of sodium larger than the atomic radius of chlorine, even though both are in period 3?
A: Chlorine has more protons (17 vs 11), so its nucleus pulls the same-shell electrons inward more strongly. Shielding is approximately the same across a period, so the effective nuclear charge increases and the radius shrinks.
Q: Which element has the highest electronegativity?
A: Fluorine.
Q: What colour of light is emitted when an electron in hydrogen drops from n = 4 to n = 2?
A: Green (part of the Balmer series).
Q: Explain why chromium's electron configuration is [Ar]4s¹3d⁵ rather than [Ar]4s²3d⁴.
A: A half-filled d subshell (3d⁵) is unusually stable. Chromium "borrows" one electron from the 4s orbital to achieve this more stable arrangement.
Electron configurations are the foundation for ionic bonding (Ch. 8) and covalent bonding (Ch. 9). You cannot draw Lewis structures or predict molecular shapes without knowing valence electron counts.
Periodic trends explain reactivity patterns in chemical reactions (Ch. 10) and the activity series used in single-replacement reactions.
The concept of quantised energy connects to thermochemistry (Ch. 16) and the energy changes in chemical reactions.
electromagnetic radiation, wavelength frequency amplitude, speed of light c = lambda nu, electromagnetic spectrum, Planck quantum, photon, photoelectric effect Einstein, atomic emission spectrum line spectrum, Bohr model hydrogen, ground state excited state, Balmer series Lyman Paschen, de Broglie wave-particle duality, Heisenberg uncertainty principle, Schrödinger wave equation, quantum mechanical model, orbitals s p d f, principal quantum number, angular momentum quantum number, magnetic quantum number, aufbau principle, Pauli exclusion principle, Hund's rule, electron configuration notation, chromium copper exceptions, valence electrons, Lewis dot structure, periodic law Mendeleev, periodic table blocks s p d f, atomic radius trend, ionisation energy trend, electronegativity trend, electron affinity, shielding, alkali metals, alkaline earth metals, halogens, noble gases, transition metals, lanthanides actinides, diamagnetic paramagnetic ferromagnetic, allotropes