CHEM Unit 1 Exam 2 Review: Atomic Structure and Periodic Trends

Vocabulary practice flashcards covering quantum numbers, orbital shapes, electron configuration rules, magnetism, effective nuclear charge, shielding, and periodic trends from CH 301 Unit 1 Exam 2 review notes.

Quantum Mechanics

A theoretical framework developed by Schrodinger to explain the wave behavior of electrons, defining their most probable locations using quantum rules rather than exact paths.

Principal Quantum Number ($n$)

The quantum number representing the electron shell, energy level, distance from the nucleus, and size of the atomic orbital ($n = 1, 2, 3, \text{...}$).

Angular Momentum Quantum Number ($l$)

The quantum number defining the subshell and the 3D shape of an energy level, taking integer values from $l = 0$ up to $n - 1$.

s Orbital Shape

A spherical atomic orbital corresponding to angular momentum quantum number $l = 0$, having only one possible spatial orientation ($m_l = 0$) and holding up to $2$ electrons.

p Orbital Shape

A dumbbell-shaped atomic orbital corresponding to angular momentum quantum number $l = 1$, possessing three spatial orientations ($m_l = -1, 0, +1$) and holding up to $6$ total electrons in its subshell.

d Orbital Shape

A multi-lobed atomic orbital corresponding to angular momentum quantum number $l = 2$, possessing five spatial orientations ($m_l = -2, -1, 0, +1, +2$) and holding up to $10$ total electrons in its subshell.

Magnetic Quantum Number ($m_l$)

The quantum number determining the 3D spatial orientation of an orbital, ranging in integer values from $-l$ through $0$ to $+l$.

Spin Quantum Number ($m_s$)

The quantum number describing the intrinsic spin direction of an individual electron within an orbital, which can only equal $+\frac{1}{2}$ or $-\frac{1}{2}$.

Aufbau's Rule

The electronic configuration principle stating that electrons must completely fill orbitals starting from the lowest energy level before occupying higher energy levels (e.g., filling $3s$ before $3p$).

Hund's Rule

The orbital filling rule stating that within degenerate orbitals (such as the $p$ or $d$ subshells), electrons spread out singly with parallel spins before doubling up in any single orbital.

Pauli Exclusion Principle

The quantum rule dictating that no two electrons in an atom can have the exact same set of four quantum numbers, which translates to a maximum capacity of two electrons with opposite spins per orbital.

Shorthand Notation

An electron configuration format where the electron core is replaced with the symbol of the preceding noble gas in brackets (e.g., $[\text{Ne}] 3s^2 3p^5$ for Cl or $[\text{Ar}] 4s^2 3d^2$ for Ti) to avoid writing all occupied inner subshells.

Valence Electron Configuration Notation

A condensed notation that lists only the electrons residing in the outermost valence subshells (e.g., $s^2 p^4$ for sulfur or $s^2 p^3$ for arsenic).

Isoelectronic

A term describing chemical species (atoms or ions) that contain identical numbers of electrons and share the same ground-state electron configuration (e.g., $\text{Cr}^{3+}$ and $\text{Sc}$, or $\text{Na}^+$ and $\text{Ne}$).

Paramagnetic

Atoms or ions containing one or more unpaired electrons in their orbital diagrams, rendering them attracted to magnetic fields (the more unpaired electrons, the stronger the magnet).

Diamagnetic

Atoms or ions that have all of their electrons paired within their orbitals, resulting in no net magnetic attraction.

Effective Nuclear Charge (ENC)

The net positive nuclear charge experienced by an electron, calculated using the formula: $\text{ENC} = \text{\# protons} - \text{\# shielding } e^-$.

Shielding Effect

The electrostatic shielding of outer valence electrons from the full positive nuclear charge caused by inner-shell core electrons, allowing valence electrons to feel a reduced effective nuclear charge.

Atomic Radius Trend

The periodic trend wherein atomic radius increases down a column due to additional shielding electron shells, and decreases left-to-right across a period as increasing effective nuclear charge pulls the valence electrons closer to the nucleus.

Isoelectronic Size Trend

The trend among isoelectronic series where species with higher nuclear charges (more protons) exert a larger effective nuclear charge, pulling electrons more tightly and resulting in smaller ionic radii (e.g., $\text{H}^- > \text{He} > \text{Li}^+ > \text{Be}^{2+} > \text{B}^{3+}$).

Ionization Energy (I.E.)

The energy required to pull a valence electron away from a given neutral atom in the gas phase, forming a positive cation.

Subshell Repulsion Exception in Ionization Energy

An anomaly where a $p^4$ configuration atom (e.g., oxygen or sulfur) exhibits a lower first ionization energy than a preceding $p^3$ configuration atom (e.g., nitrogen or phosphorus) because electron-electron repulsion within the doubly occupied $p$ orbital makes the fourth electron easier to remove.

Electronegativity (EN)

A relative measure of an atom's ability to attract shared electrons in a bond, which increases diagonally up and to the right across the periodic table, peaking at fluorine ($\text{EN} = 4.0$).

Electronegativity Difference ($\Delta\text{EN}$)

The value obtained by subtracting the smaller electronegativity value of a bonded element from the larger one (e.g., for HF: $4.0 - 2.1 = 1.9$).