Atomic Structure and Properties — AP Chemistry
1. Subatomic Particles, Isotopes, and Mass Spectrometry ★☆☆☆☆ ⏱ 4 min
All atoms are composed of three core subatomic particles that differ in charge, mass, and location within the atom. The identity of an element is defined by its number of protons, while variation in neutron number creates isotopes of the same element.
- **Protons**: Positively charged ($+e$), ~1 amu, located in the nucleus; atomic number $Z$ = number of protons
- **Neutrons**: Neutral charge (0), ~1 amu, located in the nucleus; neutron count varies between isotopes
- **Electrons**: Negatively charged ($-e$), ~1/1840 amu, located in the electron cloud outside the nucleus; electron count determines atomic charge
A = Z + N
Mass spectrometry is an analytical technique that measures mass-to-charge ratio ($m/z$) of ionized particles to determine isotopic masses and their relative natural abundances. This data is used to calculate the average atomic mass found on the periodic table.
\text{Average Atomic Mass} = \sum (\text{Isotopic Mass}_i \times \text{Relative Abundance}_i)
2. Coulomb's Law and Effective Nuclear Charge ★★☆☆☆ ⏱ 3 min
Coulomb's Law describes the electrostatic force between two charged particles, and it governs all interactions between the positive nucleus and negative electrons in an atom.
F = k \frac{q_1 q_2}{r^2}
Opposite charges produce an attractive force, while like charges produce a repulsive force. The magnitude of the force increases with larger charge values and decreases as the distance between particles increases.
Z_{eff} = Z - S
3. Electron Configurations and Orbital Diagrams ★★★☆☆ ⏱ 4 min
Electrons occupy discrete quantized energy levels (shells, labeled by principal quantum number $n$). Each shell contains subshells (s, p, d, f), which each hold multiple orbitals: regions of space with 90% probability of containing an electron. Each orbital can hold a maximum of 2 electrons with opposite spin.
- **Aufbau Principle**: Fill lower-energy orbitals before higher-energy orbitals. Order of filling: $1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p...$
- **Pauli Exclusion Principle**: No two electrons can have the same set of four quantum numbers, so each orbital holds max 2 electrons with opposite spin.
- **Hund's Rule**: For degenerate (equal energy) orbitals, fill each orbital with one parallel-spin electron before pairing electrons, to minimize repulsion.
Electron configuration notation lists subshells with the number of electrons as a superscript. Noble gas shorthand replaces core electrons with the nearest preceding noble gas in brackets. Common exceptions for period 4 transition metals: chromium ($[Ar]4s^13d^5$) and copper ($[Ar]4s^13d^{10}$) due to extra stability of half-filled and fully filled d subshells.
4. Periodic Trends ★★☆☆☆ ⏱ 3 min
Periodic trends are predictable changes in atomic properties across periods (rows) and down groups (columns) of the periodic table, driven by changes in $Z_{eff}$ and principal quantum number $n$.
5. AP-Style Concept Check ★★★☆☆ ⏱ 2 min
Common Pitfalls
Why: Rushing calculations leads to forgetting to divide percentages by 100, resulting in an answer 100 times larger than the correct value
Why: Students incorrectly assume lower subshell number equals lower energy for ions, which is not true
Why: Students forget that core electrons shield valence electrons from the full pull of the nucleus
Why: Rushing leads to ignoring Hund's rule, which minimizes electron-electron repulsion
Why: Students memorize only the general trend and forget the extra stability of half-filled (p³, d⁵) or fully filled (p⁶, d¹⁰) subshells