CH Unit 1 Exam 1 Review Flashcards Part 2 (5-10)

Vocabulary flashcards covering atomic structure, quantum numbers, electron configuration rules, periodic table classifications, and the Rydberg equation.

Rydberg Equation

A mathematical formula ($\nu = R \left(\frac{1}{n_1^2} - \frac{1}{n_2^2}\right)$) used to calculate the frequency of light emitted or absorbed when an electron jumps or falls between energy levels in a hydrogen atom.

Rydberg Constant

A fundamental constant equal to the ionization energy of a hydrogen atom ($3.3 \times 10^{15}\,\text{Hz}$) that serves as a bridge between experiment and quantum theory.

Hydrogen Energy Levels

Quantized energy states in an atom where energy levels get closer together as the electron shells get further from the proton nucleus.

Aufbau Principle

An electron configuration rule stating that electrons must fill lower energy orbitals completely before occupying higher energy orbitals.

Hund's Rule

An electron configuration rule stating that in subshells with degenerate orbitals (like p and d orbitals), electrons spread out to occupy each orbital individually before doubling up.

Pauli Exclusion Principle

A principle stating that no two electrons in an atom can have the exact same set of four quantum numbers, requiring electrons in the same orbital to have opposite spins ($m_s = +\frac{1}{2}$ or $-\frac{1}{2}$).

Potassium (K)

A highly reactive alkali metal and main group element whose single valence electron resides in the $n=4$ shell ($l=0$), causing it to react readily with non-metals to form salts.

Principal Quantum Number ($n$)

The quantum number ($n = 1, 2, 3, \dots$) that specifies the shell, overall energy level, size of an atomic orbital, and distance of the electron from the nucleus.

Angular Momentum Quantum Number ($l$)

The quantum number ($l = 0, 1, \dots, n-1$) that defines the subshell type and the three-dimensional shape of an atomic orbital.

Magnetic Quantum Number ($m_l$)

The quantum number ($m_l = -l, \dots, +l$) that designates the specific atomic orbital and its spatial orientation within a subshell.

Spin Quantum Number ($m_s$)

The quantum number indicating the spin state of an electron within an orbital, which can only take values of $+\frac{1}{2}$ or $-\frac{1}{2}$.

s Orbital Shape

A spherical orbital shape corresponding to an angular momentum quantum number of $l=0$.

p Orbital Shape

A dumbbell-shaped atomic orbital corresponding to an angular momentum quantum number of $l=1$.

d Subshell Capacity

A subshell characterized by $l=2$ containing 5 individual orbitals, accommodating a maximum total of 10 electrons ($5 \times 2 = 10$).

Schrödinger

The physicist who developed quantum mechanics and wave equations to describe the wave behavior and energy levels of electrons in atoms.

Rydberg Equation

A mathematical formula ($\nu = R \left(\frac{1}{n_1^2} - \frac{1}{n_2^2}\right)$) used to calculate the frequency of light emitted or absorbed when an electron jumps or falls between energy levels in a hydrogen atom.

Rydberg Constant

A fundamental constant equal to the ionization energy of a hydrogen atom ($3.3 \times 10^{15}\,\text{Hz}$) that serves as a bridge between experiment and quantum theory.

Hydrogen Energy Levels

Quantized energy states in an atom where energy levels get closer together as the electron shells get further from the proton nucleus.

Aufbau Principle

An electron configuration rule stating that electrons must fill lower energy orbitals completely before occupying higher energy orbitals.

Hund's Rule

An electron configuration rule stating that in subshells with degenerate orbitals (like p and d orbitals), electrons spread out to occupy each orbital individually before doubling up.

Pauli Exclusion Principle

A principle stating that no two electrons in an atom can have the exact same set of four quantum numbers, requiring electrons in the same orbital to have opposite spins ($m_s = +\frac{1}{2}$ or $-\frac{1}{2}$).

Potassium (K)

A highly reactive alkali metal and main group element whose single valence electron resides in the $n=4$ shell ($l=0$), causing it to react readily with non-metals to form salts.

Principal Quantum Number ($n$)

The quantum number ($n = 1, 2, 3, \dots$) that specifies the shell, overall energy level, size of an atomic orbital, and distance of the electron from the nucleus.

Angular Momentum Quantum Number ($l$)

The quantum number ($l = 0, 1, \dots, n-1$) that defines the subshell type and the three-dimensional shape of an atomic orbital.

Magnetic Quantum Number ($m_l$)

The quantum number ($m_l = -l, \dots, +l$) that designates the specific atomic orbital and its spatial orientation within a subshell.

Spin Quantum Number ($m_s$)

The quantum number indicating the spin state of an electron within an orbital, which can only take values of $+\frac{1}{2}$ or $-\frac{1}{2}$.

s Orbital

A spherical orbital shape corresponding to an angular momentum quantum number of $l=0$.

p Orbital Shape

A dumbbell-shaped atomic orbital corresponding to an angular momentum quantum number of $l=1$.

d Subshell Capacity

A subshell characterized by $l=2$ containing 5 individual orbitals, accommodating a maximum total of 10 electrons ($5 \times 2 = 10$).

Schrödinger

The physicist who developed quantum mechanics and wave equations to describe the wave behavior and energy levels of electrons in atoms.

Scenario: An electron is placed in a higher-energy $2p$ orbital before the lower-energy $2s$ orbital is completely filled.

Violates the Aufbau Principle. Electrons must fill lower-energy orbitals completely before occupying orbitals of higher energy levels.

Scenario: An orbital diagram shows two electrons paired in one $2p$ orbital while an adjacent $2p$ orbital in the same subshell remains completely empty.

Violates Hund's Rule. Electrons in subshells with degenerate orbitals (orbitals of equal energy) must occupy each orbital singly with parallel spins before doubling up.

Scenario: An orbital diagram shows two electrons in the same $1s$ orbital drawn with identical upward-pointing spins ($m_s = +\frac{1}{2}$ for both).

Violates the Pauli Exclusion Principle. Electrons sharing the same orbital must have opposite spins ($m_s = +\frac{1}{2}$ and $m_s = -\frac{1}{2}$) so that no two electrons share the exact same set of four quantum numbers.