CHEM 211 Unit Exam 2: Electronic Structure & Periodic Properties Vocabulary

Vocabulary practice flashcards covering electronic structure, quantum mechanics, electron configurations, and periodic trends based on the Unit Exam 2 study guide for CHEM 211.

Electromagnetic Radiation

Energy carried through space by oscillating electric and magnetic fields.

Wave

A disturbance that transfers energy through space.

Amplitude

The height of a wave from its center line to its crest or trough, where greater amplitude corresponds to greater intensity or brightness.

Wavelength ($\lambda$)

The distance between identical points on consecutive waves, such as crest to crest, commonly measured in meters ($\text{m}$) or nanometers ($\text{nm}$) where $1\text{ nm} = 1 \times 10^{-9}\text{ m}$.

Frequency ($\nu$)

The number of complete wave cycles that pass a point each second, measured in Hertz ($\text{Hz}$) or $\text{s}^{-1}$.

Speed of Light ($c$)

The constant speed at which electromagnetic radiation travels in a vacuum, represented by the equation $c = \lambda \nu$.

Visible Light

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

Electromagnetic Spectrum

The full range of electromagnetic radiation ordered from lowest energy/longest wavelength to highest energy/shortest wavelength: Radio, Microwave, Infrared, Visible, Ultraviolet, X-ray, and Gamma.

Constructive Interference

The phenomenon occurring when waves overlap in phase, combining their amplitudes to create a larger wave.

Destructive Interference

The phenomenon occurring when waves overlap out of phase, causing them to partially or completely cancel.

Interference Pattern

A pattern of alternating areas produced by constructive and destructive interference.

Diffraction

The spreading or bending of waves as they pass through an opening or around an obstacle.

Blackbody Radiation

Electromagnetic radiation emitted by an object because of its temperature.

Planck's Explanation

Max Planck's proposal that matter can emit or absorb energy only in specific, discrete amounts called quanta, given by $E = n h \nu$.

Quantum

A discrete or fixed amount of energy that can be emitted or absorbed.

Photon

A discrete packet of electromagnetic radiation whose energy is $E = h \nu$ or $E = \frac{h c}{\lambda}$.

Photoelectric Effect

The ejection of electrons from a metal surface when light with a frequency at or above a required threshold frequency strikes the metal.

Atomic Spectroscopy

The study of electromagnetic radiation absorbed or emitted by atoms.

Line Emission Spectrum

A series of specific wavelengths of light emitted when excited electrons fall from higher energy levels to lower energy levels.

Bohr Model

An atomic model proposing that electrons occupy specific, quantized energy levels around the nucleus and cannot exist between allowed levels.

Absorption

The process in which an electron absorbs energy and moves from a lower energy level to a higher energy level ($\Delta E > 0$).

Emission

The process in which an electron falls from a higher energy level to a lower energy level, releasing energy as a photon ($\Delta E < 0$).

de Broglie Wavelength

The wavelength associated with a particle such as an electron, calculated using $\lambda = \frac{h}{m v}$.

Heisenberg Uncertainty Principle

The principle stating that it is impossible to simultaneously know both the exact position and exact momentum of an electron.

Quantum Mechanical Model

An atomic model that describes electron locations using probability distributions called orbitals rather than fixed circular paths.

Orbital

A three-dimensional region around the nucleus where there is a high probability of finding an electron.

Principal Quantum Number ($n$)

The quantum number describing the main energy level, general size, and general energy of an orbital, with allowed values $n = 1, 2, 3, 4\dots$

Angular Momentum Quantum Number ($l$)

The quantum number describing the subshell and orbital shape, with possible integer values from $0$ to $n - 1$.

Magnetic Quantum Number ($m_l$)

The quantum number describing the orientation of an orbital in space, with values ranging from $-l$ through $+l$.

Spin Quantum Number ($m_s$)

The quantum number describing an electron's spin, with allowed values of $+\frac{1}{2}$ or $-\frac{1}{2}$.

s Orbital

A spherical orbital corresponding to $l = 0$, containing $1$ orbital per subshell and a maximum capacity of $2$ electrons.

p Orbitals

Dumbbell-shaped orbitals corresponding to $l = 1$, containing $3$ orbitals per subshell and a maximum capacity of $6$ electrons.

d Orbitals

Mostly cloverleaf-shaped orbitals corresponding to $l = 2$, containing $5$ orbitals per subshell and a maximum capacity of $10$ electrons.

f Orbitals

Complex/multilobed orbitals corresponding to $l = 3$, containing $7$ orbitals per subshell and a maximum capacity of $14$ electrons.

Subshell Quantum Capacities Chart

A table organizing subshell labels, angular momentum quantum numbers ($l$), number of orbitals, and maximum electron capacities.

Degenerate Orbitals

Orbitals that possess the exact same energy level.

Electron Configuration

A notation showing how electrons are distributed among an atom's orbitals.

Aufbau Principle

The principle stating that electrons occupy the lowest-energy available orbitals first.

Pauli Exclusion Principle

The principle stating that no two electrons in an atom can have the same four quantum numbers, limiting each orbital to a maximum of $2$ electrons with opposite spins.

Hund's Rule

The rule stating that electrons fill degenerate orbitals one at a time with parallel spins before pairing.

Orbital Diagram

A visual representation of electron distributions using boxes or lines for orbitals and arrows for electron spins.

Core Electrons

Inner electrons that generally do not participate directly in chemical bonding.

Valence Electrons

Electrons in the outermost principal energy level available for chemical bonding.

Noble-Gas Configuration

An abbreviated electron configuration in which inner core electrons are represented by the bracketed symbol of the preceding noble gas.

Transition-Metal Cation Formation Rule

The rule stating that when forming transition-metal cations, outer s electrons are removed before d electrons.

Atomic Radius

One-half the distance between the nuclei of two identical bonded atoms.

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

The net positive attraction experienced by an electron from the nucleus after accounting for shielding, given approximately by $Z_{\text{eff}} = Z - S$.

Cations

Positively charged ions formed by losing electrons, which are always smaller than their neutral parent atoms.

Anions

Negatively charged ions formed by gaining electrons, which are always larger than their neutral parent atoms.

Isoelectronic

Atoms or ions that have the same number of electrons and identical electron configurations.

First Ionization Energy

The energy required to remove the outermost electron from a neutral gaseous atom.

Electron Affinity

The energy change associated with adding an electron to a gaseous atom.

Periods

Horizontal rows on the periodic table, also known as series.

Groups

Vertical columns on the periodic table, also known as families, containing elements with similar chemical properties.

Metals

Elements located primarily on the left and center of the periodic table that conduct heat and electricity, are malleable, and tend to lose electrons.

Nonmetals

Elements located primarily on the upper-right side of the periodic table (plus Hydrogen) that tend to gain or share electrons.

Metalloids

Elements located along the stair-step boundary between metals and nonmetals that possess properties intermediate between metals and nonmetals.

Main-Group Elements

Elements located in the s and p blocks of the periodic table (Groups 1–2 and 13–18).

Transition Metals

Elements located in the d block of the periodic table (Groups 3–12).

Inner Transition Metals

Elements located in the f block of the periodic table, comprising the lanthanides and actinides.

Alkali Metals

Group 1 elements (excluding Hydrogen) that are highly reactive metals and typically form $+1$ ions.

Alkaline Earth Metals

Group 2 elements that are reactive metals and typically form $+2$ ions.

Chalcogens

Group 16 elements, including Oxygen and Sulfur, that often form $-2$ ions.

Halogens

Group 17 reactive nonmetals that often form $-1$ ions.

Noble Gases

Group 18 elements with completely filled valence shells that are generally very unreactive.