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Atomic Structure

Chapter 1: Atomic Structure

5 min read Updated Mar 26, 2026
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🎯 Diagnostic: Test Your Starting Level 24 questions (2 per section). No prior reading required, see what you already know.

Aim to answer every question before checking. Missed questions point you to the sections you need most.

1. (1.1) Which subatomic particle has approximately unit mass and no charge?
A. Proton: +1, ~1 amu. Neutron: 0, ~1 amu. Electron: -1, ~11836\frac{1}{1836} amu. Positrons are positive electrons (antimatter).
2. (1.1) The mass of an electron compared to a proton is approximately:
D. The electron is so light that atomic mass is essentially due to protons and neutrons.
3. (1.2) The atomic number (Z) equals:
B. Z defines the element. In a neutral atom, Z also equals the number of electrons. Mass number A = Z + N (neutrons).
4. (1.2) Isotopes of an element share the same:
C. Same element, different neutrons. Chemistry is essentially identical; only mass and nuclear stability change.
5. (1.3) The Bohr model proposed that:
D. Bohr quantized angular momentum as nh/2π, giving stable orbits. The model works only for one-electron systems (H, He⁺), not multi-electron atoms.
6. (1.3) The energy of a hydrogen electron at level n is:
A. Negative energy means the electron is bound. As n rises, energy approaches zero (ionization at n = ∞).
7. (1.4) The quantum mechanical model differs from the Bohr model because it:
C. Schrödinger's wave equation gives orbitals (regions of high electron probability) instead of precise orbits.
8. (1.4) Heisenberg's uncertainty principle states:
B. Δx · Δp ≥ ħ/2. The more precisely we pin down position, the less we know about momentum. This is fundamental, not a measurement limitation.
9. (1.5) The principal quantum number (n) specifies:
D. Higher n = higher energy and average radius. Shape is set by l, orientation by mₗ, and spin by mₛ.
10. (1.5) The azimuthal quantum number (l) can take values:
A. For n = 3, l can be 0 (s), 1 (p), or 2 (d). The magnetic quantum number mₗ runs from -l to +l.
11. (1.6) Ground-state electron configurations obey:
C. Aufbau fills from lowest energy upward; Pauli says ≤2 electrons per orbital with opposite spin; Hund says maximize parallel spins in degenerate orbitals.
12. (1.6) The ground-state configuration of carbon (Z = 6) is:
B. Fill 1s, 2s, then 2p (one electron in each of 2 separate p orbitals per Hund's rule).
13. (1.7) The shape of an s orbital is:
A. s orbitals are spherically symmetric. Electron density depends only on radial distance.
14. (1.7) The three p orbitals in a given shell are:
D. Three p orbitals (pxp_{x}, pyp_{y}, pzp_{z}) are mutually perpendicular. d orbitals (four cloverleaf + one dumbbell with donut) follow.
15. (1.8) Valence electrons are:
C. Valence electrons are the chemistry: the octet rule, Lewis structures, and reactivity all depend on them.
16. (1.8) Chlorine (Z = 17) has how many valence electrons?
B. Group 17 (halogen) = 7 valence electrons (3s² 3p⁵). Cl⁻ gains one to reach 8 and the Ar configuration.
17. (1.9) When an atom becomes a cation:
D. Cations are smaller than their parent atoms. The nucleus is untouched; only the electron cloud contracts.
18. (1.9) For transition-metal cations, electrons are removed first from:
A. Classic example: Fe is [Ar]4s² 3d⁶, but Fe²⁺ is [Ar]3d⁶. The 4s electrons go first.
19. (1.10) The photoelectric effect demonstrated that light behaves:
B. Einstein's explanation (1905, Nobel 1921) tied electron ejection to photon frequency rather than intensity, establishing wave-particle duality for light.
20. (1.10) The kinetic energy of an emitted photoelectron is:
C. KE_max = hν - φ. Intensity sets how many electrons; frequency sets how energetic each one is.
21. (1.11) Line emission spectra arise when:
A. Each line corresponds to an electron transition. The hydrogen Balmer series is the prototypical example.
22. (1.11) Absorption spectra differ from emission spectra in that they show:
D. The absorption lines occur at the same wavelengths as emission lines, but they appear as dark gaps in a continuous spectrum.
23. (1.12) Fluorescence is best described as:
C. Excited singlet relaxes back to ground state on a nanosecond scale. The emitted photon is usually lower energy than the absorbed one (Stokes shift).
24. (1.12) Phosphorescence is characterized by:
B. Intersystem crossing traps the electron in the triplet state. The long-lived glow of glow-in-the-dark items uses this mechanism.

You are made of atoms. So is this screen, the air you are breathing, and the coffee keeping you awake right now. Every reaction you will study in general chemistry, organic chemistry, and biochemistry starts with atoms and their electrons. If you understand how atoms are built and how their electrons behave, you hold the key to everything else in this book.

This chapter is not just “Chapter 1 material” to check off a list. The MCAT tests atomic structure directly - and it also hides atomic structure inside passages about spectroscopy, bonding, periodic trends, and even drug design. A passage might hand you an emission spectrum and ask you to calculate the energy of a transition. Another might describe an experiment and ask why a magnetic field splits spectral lines. The concepts you build here will pay dividends in every chapter that follows.

The Atom is a Building with Strict Rules

Think of an atom as a high-rise building with a very unusual elevator. The nucleus sits in the basement - tiny, dense, and holding almost all of the building’s mass. Electrons live on specific floors above, and here is the key: there are no stairs. Electrons cannot exist between floors. They must absorb exactly the right amount of energy to jump up a floor, and they release exactly that amount of energy (as light) when they drop back down.

Each floor has a limited number of rooms (orbitals), and each room holds at most two tenants (electrons) who must have opposite spins. The rules about which floors fill first, how many rooms exist per floor, and how tenants choose rooms - those rules are the Aufbau principle, Hund’s rule, and the Pauli exclusion principle. Master those rules, and you can write the electron configuration of any element on the periodic table.

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