Atomic Structure
All chemistry rests on atoms. Understanding how atoms are built — their subatomic particles, electron arrangements, and position in the periodic table — is the foundation for predicting chemical behavior. This skill covers atomic models from Dalton to quantum mechanics, electron configuration rules, periodic table organization, periodic trends, and isotopes with radioactivity.
Agent affinity: mendeleev (periodic/inorganic chemistry, primary), curie-m (nuclear/radiochemistry, for isotope and decay topics)
Concept IDs: chem-atomic-structure, chem-periodic-table-organization, chem-periodic-trends, chem-isotopes-radioactivity
Historical Atomic Models
| # | Model | Year | Key idea | Limitation | |---|---|---|---|---| | 1 | Dalton | 1803 | Indivisible solid spheres | No internal structure | | 2 | Thomson | 1897 | "Plum pudding" — electrons in positive matrix | No nucleus | | 3 | Rutherford | 1911 | Dense positive nucleus, electrons orbit | Classical orbits radiate energy, collapse | | 4 | Bohr | 1913 | Quantized circular orbits, energy levels | Only works for hydrogen | | 5 | Quantum mechanical | 1926 | Probability orbitals (Schrodinger equation) | Full model — currently accepted |
Each model was displaced by experimental evidence the previous model could not explain. Rutherford's gold foil experiment demolished Thomson's model. Bohr's model explained hydrogen line spectra but failed for multi-electron atoms. The quantum mechanical model, based on the Schrodinger equation, treats electrons as probability clouds (orbitals) rather than particles on fixed paths.
Subatomic Particles
| Particle | Symbol | Charge | Mass (amu) | Location | |---|---|---|---|---| | Proton | p+ | +1 | 1.0073 | Nucleus | | Neutron | n0 | 0 | 1.0087 | Nucleus | | Electron | e- | -1 | 0.00055 | Electron cloud |
Atomic number (Z): Number of protons. Defines the element. Carbon is always Z = 6.
Mass number (A): Protons + neutrons. Written as superscript-left: 12-C or carbon-12.
Charge: Protons minus electrons. Neutral atom: protons = electrons.
Quantum Numbers and Orbitals
Each electron in an atom is described by four quantum numbers:
| Quantum number | Symbol | Values | Describes | |---|---|---|---| | Principal | n | 1, 2, 3, ... | Energy level (shell) | | Angular momentum | l | 0 to n-1 | Orbital shape (s, p, d, f) | | Magnetic | m_l | -l to +l | Orbital orientation | | Spin | m_s | +1/2 or -1/2 | Electron spin direction |
Orbital shapes: s = spherical (l=0), p = dumbbell (l=1), d = cloverleaf (l=2), f = complex multi-lobed (l=3).
Capacity per subshell: 2(2l + 1). So s holds 2, p holds 6, d holds 10, f holds 14.
Electron Configuration Rules
Three rules govern how electrons fill orbitals:
1. Aufbau Principle. Electrons fill the lowest-energy orbitals first. The filling order follows increasing (n + l), with lower n breaking ties:
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p
2. Pauli Exclusion Principle. No two electrons in the same atom can have all four quantum numbers identical. Each orbital holds at most 2 electrons with opposite spins.
3. Hund's Rule. Within a subshell of equal-energy orbitals, electrons occupy empty orbitals first (one per orbital, same spin) before pairing.
Worked Example: Electron Configuration of Iron (Z = 26)
Problem. Write the full electron configuration and orbital diagram for iron.
Solution. Fill orbitals in Aufbau order, distributing 26 electrons:
1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^6
Check: 2 + 2 + 6 + 2 + 6 + 2 + 6 = 26. Correct.
Noble gas shorthand: [Ar] 4s^2 3d^6, where [Ar] = 1s^2 2s^2 2p^6 3s^2 3p^6 (18 electrons).
Orbital diagram for 3d subshell (Hund's rule):
3d: [up/down] [up/down] [up] [up] [up] — this is wrong. Let me re-count. 3d^6 means 6 electrons in 5 d orbitals. By Hund's rule, fill each of the 5 orbitals with one electron first (5 electrons, all spin-up), then the 6th pairs in the first orbital:
3d: [up/down] [up] [up] [up] [up]
Iron has 4 unpaired electrons, making it paramagnetic.
Worked Example: Exception — Chromium (Z = 24)
Problem. Predict and correct the electron configuration of chromium.
Expected by Aufbau: [Ar] 4s^2 3d^4
Actual: [Ar] 4s^1 3d^5
Why. A half-filled d subshell (d^5) has extra stability due to exchange energy — the quantum mechanical stabilization from having all five d orbitals singly occupied with parallel spins. Chromium "steals" one electron from 4s to achieve this. Copper (Z = 29) does the same: [Ar] 4s^1 3d^10 rather than 4s^2 3d^9, preferring the fully filled d^10.
Periodic Table Organization
Periodic Law. When elements are arranged by increasing atomic number, their physical and chemical properties repeat periodically. This is the single most powerful organizing principle in chemistry.
Structure:
| Feature | Description | |---|---| | Period (row) | Elements in the same period have the same highest principal quantum number n | | Group (column) | Elements in the same group have the same valence electron configuration pattern | | s-block | Groups 1-2 + He: filling s orbitals | | p-block | Groups 13-18: filling p orbitals | | d-block | Groups 3-12: filling d orbitals (transition metals) | | f-block | Lanthanides + actinides: filling f orbitals (inner transition metals) |
Group names worth knowing: Group 1 = alkali metals, Group 2 = alkaline earth metals, Group 17 = halogens, Group 18 = noble gases.
Metals, nonmetals, metalloids. Metals dominate the left and center (lose electrons, conduct). Nonmetals cluster in the upper right (gain electrons, insulate). Metalloids straddle the staircase line (B, Si, Ge, As, Sb, Te) — semiconducting properties.
Periodic Trends
Five trends follow directly from nuclear charge and electron shielding:
Trend 1 — Atomic Radius
Across a period (left to right): Decreases. More protons pull the same-shell electrons inward.
Down a group: Increases. Each period adds a new shell farther from the nucleus.
Worked example. Rank Na, Mg, K in order of increasing atomic radius.
Na and Mg are in period 3; K is in period 4. Down group 1: K > Na. Across period 3: Na > Mg (more protons, same shell). Order: Mg < Na < K.
Trend 2 — Ionization Energy (IE)
Definition. Energy required to remove the highest-energy electron from a gaseous atom. Always positive (endothermic).
Across a period: Increases (electrons held more tightly as Z_eff grows).
Down a group: Decreases (valence electrons farther from nucleus, more shielding).
Worked example. Why is the first ionization energy of oxygen (1314 kJ/mol) lower than that of nitrogen (1402 kJ/mol), despite oxygen being further right?
Nitrogen has a half-filled 2p^3 configuration (exchange energy stabilization). Oxygen's 2p^4 has one paired electron — electron-electron repulsion in the paired orbital makes that electron easier to remove. This is the classic "dip" in the IE trend at Group 16.
Trend 3 — Electronegativity
Definition. Tendency of an atom in a bond to attract shared electrons. Pauling scale: F = 4.0 (highest), Cs = 0.79 (lowest of main group).
Across a period: Increases. Down a group: Decreases.
Practical consequence. Electronegativity differences predict bond type: less than 0.4 = nonpolar covalent, 0.4 to 1.7 = polar covalent, greater than 1.7 = ionic (approximate guidelines).
Trend 4 — Electron Affinity
Definition. Energy change when a gaseous atom gains an electron. More negative = more favorable.
General trend. Becomes more negative (more exothermic) moving right across a period and up a group. Halogens have the most negative electron affinities. Noble gases have near-zero or positive (unfavorable) electron affinities — their shells are already full.
Trend 5 — Metallic Character
Across a period: Decreases. Down a group: Increases.
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