Chemistry: Electronic Structure of Atoms and Periodic Properties

Vocabulary flashcards covering electronic structure, quantum mechanics, electron configurations, and periodic trends based on Chapters 6 and 7.

Wavelength ($\lambda$)

The distance between corresponding points on adjacent waves.

Frequency ($\nu$)

The number of complete waves passing a given point per unit of time.

Speed of Light ($c$)

The universal constant speed at which all electromagnetic radiation travels through a vacuum, equal to $3.00 \times 10^8\,\text{m/s}$.

Quantum

A fixed, discrete packet of energy that can be absorbed or released by matter.

Photon

A stream of particle-like quantized energy that has no mass and carries a quantum of energy calculated by $E = h\nu$.

Photoelectric Effect

The phenomenon in which electrons are ejected from the surface of a metal when light of at least a minimum threshold frequency strikes it.

Planck's Constant ($h$)

The physical constant relating the energy of a photon to its frequency, equal to $6.626 \times 10^{-34}\,\text{J}\cdot\text{s}$.

Line Spectrum

A spectrum containing light of only discrete, specific wavelengths emitted by electronically excited atoms.

Ground State

The lowest energy state of an electron in an atom.

Excited State

Any state of an electron in an atom that has a higher potential energy than its ground state.

Heisenberg Uncertainty Principle

The principle stating that it is impossible to simultaneously know both the exact momentum ($mv$) and position ($x$) of a particle with absolute precision.

Electron Density

The probability of finding an electron at a particular location in an atom, given by the square of the wave function ($\psi^2$).

Principal Quantum Number ($n$)

The quantum number that designates the main energy level or shell on which an orbital resides, taking integer values $n \ge 1$.

Angular Momentum Quantum Number ($l$)

The quantum number that defines the shape of an orbital, with allowed integer values ranging from $0$ to $n - 1$.

Magnetic Quantum Number ($m_l$)

The quantum number that describes the three-dimensional orientation of an orbital in space, taking integer values from $-l$ to $l$.

Spin Quantum Number ($m_s$)

The quantum number describing the intrinsic spin direction and magnetic field of an electron, having allowed values of $+\frac{1}{2}$ or $-\frac{1}{2}$.

Node

A region in an orbital where the wave function passes through zero and there is zero probability of finding an electron.

Degenerate Orbitals

Orbitals that belong to the same sublevel and possess identical energy levels.

Aufbau Principle

The principle stating that electrons fill the lowest-energy orbitals available in the ground state of an atom before filling higher-energy orbitals.

Pauli Exclusion Principle

The principle stating that no two electrons in the same atom can have identical sets of all four quantum numbers ($n, l, m_l, m_s$).

Hund's Rule

The rule stating that when filling degenerate orbitals, the lowest energy arrangement is achieved when electrons occupy orbitals singly with parallel spins before pairing up.

Valence Electrons

The electrons in the outermost occupied shell ($n$) of an atom that participate in chemical bonding.

Core Electrons

Inner-shell electrons of an atom corresponding to filled noble-gas configurations and completed $d$ or $f$ sublevels.

p Orbitals

Dumbbell-shaped orbitals with angular momentum quantum number $l = 1$, existing as three degenerate orientations ($p_x, p_y, p_z$).

d Orbitals

Orbitals with angular momentum quantum number $l = 2$, four of which have four lobes and one which resembles a dumbbell surrounded by a donut.

Periodicity

The repetitive, predictable pattern of chemical and physical properties exhibited by elements when arranged by increasing atomic number.

Effective Nuclear Charge ($Z_{\text{eff}}$)

The net positive charge felt by a valence electron, calculated using $Z_{\text{eff}} = Z - S$.

Screening Constant ($S$)

A numerical factor accounting for the nuclear shielding caused by core electrons repelling valence electrons, usually close to the number of inner electrons.

Bonding Atomic Radius

Half of the distance between the nuclei of two bonded atoms in a molecule.

Nonbonding Atomic Radius

Half of the shortest distance separating the nuclei of two unbonded atoms when they collide.

Isoelectronic Series

A group of ions and atoms containing the same total number of electrons, such as $N^{3-}, O^{2-}, F^{-}, Ne, Na^{+}, Mg^{2+},$ and $Al^{3+}$.

First Ionization Energy ($I_1$)

The minimum energy required to remove the first electron from the ground state of a neutral, gaseous atom.

Electron Affinity

The energy change accompanying the addition of an electron to a gaseous atom to form an anion.

Electronegativity

A scale ranging from 0 to 4.0 that measures the attraction an atom has for shared bonding electrons within a compound.

Allotropes

Different elemental forms of the same element in the same physical state, such as dioxygen ($O_2$) and ozone ($O_3$).

de Broglie Wavelength ($\lambda$)

The wavelength associated with a moving particle of matter, calculated by dividing Planck's constant by momentum ($\lambda = \frac{h}{mv}$).

Rydberg Constant ($R_H$)

An empirical physical constant appearing in the mathematical formula used to predict the discrete wavelengths of spectral lines emitted by hydrogen atoms.

Continuous Spectrum

An unbroken sequence of radiation containing all visible wavelengths without gaps, produced when white light passes through a prism.

Alkali Metals

Group 1A soft metallic solids with low densities, low ionization energies, and violent exothermic reactivity with water.

Halogens

Group 7A typical nonmetals with highly negative electron affinities that readily gain electrons to form $-1$ anions.

Noble Gases

Group 8A unreactive monatomic gases with completely filled valence $s$ and $p$ sublevels and positive electron affinities.

Metal Oxides

Basic ionic compounds formed between metals and oxygen that react with acids to form a salt and water.

Nonmetal Oxides

Acidic molecular compounds formed between nonmetals and oxygen that react with bases.

Electromagnetic Radiation

Radiated energy composed of perpendicular electrical and magnetic waves traveling through space at the speed of light.

Wavelength ($\lambda$)

The distance between corresponding points on adjacent waves.

Frequency ($\nu$)

The number of complete wave cycles passing a given point per unit of time, measured in units of $\text{s}^{-1}$ or Hertz ($\text{Hz}$).

Speed of Light Equation

The formula $c = \lambda \nu$, which relates speed of light ($3.00 \times 10^8\,\text{m/s}$) to wavelength and frequency.

Visible Light Wavelength Range

The region of the electromagnetic spectrum detectable by the human eye, spanning approximately $400\,\text{nm}$ to $700\,\text{nm}$.

Electromagnetic Spectrum Order (Increasing Wavelength)

Gamma rays, X-rays, Ultraviolet, Visible light, Infrared, Microwaves, Radio waves.

Blackbody Radiation

The emission of light from hot, glowing objects, which classical wave theory could not explain without quantum mechanics.

Quantum

The minimum discrete packet of energy that can be absorbed or released by matter.

Planck's Constant ($h$)

The proportionality constant relating photon energy to frequency, equal to $6.626 \times 10^{-34}\,\text{J}\cdot\text{s}$.

Photon Energy Equation

The formula $E = h\nu$, describing the energy carried by a single photon of electromagnetic radiation.

Photoelectric Effect

The ejection of electrons from a metal surface when light exceeding a minimum threshold frequency strikes it.

Flame Test Color of Copper

Green flame emission.

Flame Test Color of Sodium

Yellow flame emission.

Flame Test Color of Strontium

Red flame emission.

Flame Test Color of Calcium

Orange-red flame emission.

Flame Test Color of Potassium

Purple flame emission.

Flame Test Color of Lithium

Purple-red flame emission.

Flame Test Color of Barium

Yellow-green flame emission.

Continuous Spectrum

An unbroken rainbow of radiation containing all visible wavelengths without gaps, produced by white light sources.

Line Spectrum

A spectrum containing light of only discrete, specific wavelengths emitted by excited gaseous atoms.

Bohr Model Postulate for Hydrogen

Electrons orbit the nucleus only in specific permitted circular radii corresponding to allowed energy states without radiating energy.

Ground State

The lowest possible potential energy state of an electron in an atom ($n = 1$).

Excited State

Any allowed electronic state with higher potential energy than the ground state ($n > 1$).

Sign of Positive Energy Change ($\Delta E > 0$) in Electronic Transitions

Indicates energy absorption by the atom when an electron moves to a higher energy level ($n_f > n_i$).

Sign of Negative Energy Change ($\Delta E < 0$) in Electronic Transitions

Indicates energy release as a photon when an electron drops to a lower energy level ($n_f < n_i$).

Limitations of the Bohr Model

It accurately predicts line spectra only for single-electron systems like hydrogen and incorrectly assumes circular orbital paths.

de Broglie Wavelength Equation

The formula $\lambda = \frac{h}{mv}$, describing the wavelength associated with a moving particle of mass $m$ and velocity $v$.

Heisenberg Uncertainty Principle

The fundamental limit stating that position ($x$) and momentum ($mv$) of a subatomic particle cannot both be known precisely: $(\Delta x)(\Delta mv) \ge \frac{h}{4\pi}$.

Electron Density ($\psi^2$)

The probability of finding an electron at a specific location in space, given by the square of Schrödinger's wave function.

Principal Quantum Number ($n$)

Designates the main energy level or shell of an orbital, taking positive integer values $n \ge 1$.

Angular Momentum Quantum Number ($l$)

Defines the shape of an orbital or subshell, taking integer values from $0$ to $n - 1$.

Magnetic Quantum Number ($m_l$)

Describes the 3D orientation of an orbital in space, taking integer values from $-l$ to $+l$.

Spin Quantum Number ($m_s$)

Describes the intrinsic spin direction and magnetic field of an electron, taking values of $+\frac{1}{2}$ or $-\frac{1}{2}$.

Sublevel Letter Designations for $l$ Values

$l=0$ corresponds to $s$, $l=1$ to $p$, $l=2$ to $d$, and $l=3$ to $f$.

Total Number of Orbitals in Shell $n$

Calculated by $n^2$.

Maximum Electron Capacity of Shell $n$

Calculated by $2n^2$.

Radial Node

A spherical surface where the probability of finding an electron drops to zero ($\psi^2 = 0$).

Number of Radial Nodes in an $ns$ Orbital

Calculated by $n - 1$.

Degenerate Orbitals

Orbitals belonging to the same subshell that possess identical energy levels.

Aufbau Principle

Electrons fill the lowest available energy orbitals in the ground state before occupying higher energy levels.

Pauli Exclusion Principle

No two electrons in the same atom can have an identical set of all four quantum numbers ($n, l, m_l, m_s$).

Hund's Rule

When filling degenerate orbitals, electrons occupy them singly with parallel spins before pairing up to minimize electron-electron repulsion.

Valence Electrons

Electrons in the outermost occupied principal energy shell ($n$) involved in chemical bonding.

Core Electrons

Inner-shell electrons corresponding to filled noble gas configurations and completed inner $d$ or $f$ sublevels.

Ground-State Electron Configuration of Chromium ($Z = 24$)

$[\text{Ar}]\,4s^1\,3d^5$, anomalous due to extra stability of half-filled subshells.

Ground-State Electron Configuration of Copper ($Z = 29$)

$[\text{Ar}]\,4s^1\,3d^{10}$, anomalous due to extra stability of a completely filled $d$ subshell.

Mendeleev's Development of the Periodic Table

Arranged elements by atomic mass and chemical reactivity, predicting undiscovered elements like germanium.

Moseley's Contribution to the Periodic Table

Established atomic number (proton count) through X-ray spectroscopy as the true fundamental basis for periodic organization.

Periodicity

The repetitive, predictable pattern of physical and chemical properties exhibited by elements when arranged by atomic number.

Effective Nuclear Charge ($Z_{\text{eff}}$)

The net positive nuclear charge experienced by a valence electron, calculated as $Z_{\text{eff}} = Z - S$.

Screening Constant ($S$)

A numerical value representing inner-shell core electron shielding that repels valence electrons from the nucleus.

Periodic Trend in Effective Nuclear Charge ($Z_{\text{eff}}$)

Increases significantly from left to right across a period and increases slightly down a group.

Bonding Atomic Radius

Half the distance between the nuclei of two identical atoms joined by a covalent bond.

Nonbonding Atomic Radius (van der Waals Radius)

Half the shortest distance separating the nuclei of two unbonded identical atoms when they collide.

Periodic Trend in Atomic Radius Across a Period

Decreases from left to right because increasing $Z_{\text{eff}}$ pulls valence electrons closer to the nucleus.

Periodic Trend in Atomic Radius Down a Group

Increases from top to bottom because valence electrons enter higher principal quantum levels ($n$) farther from the nucleus.

Relative Size of Cations vs Parent Atoms

Cations are smaller than their parent neutral atoms due to electron loss, reduced repulsions, and increased nuclear pull.