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Home :Glossary

Glossary: Electrons in atoms

Glossary
absorption spectrum. absorption spectra. Compare with absorption spectroscopy*.
A plot that shows how much radiation a substance absorbs at different wavelengths*. Absorption spectra are unique for each element and compound and they are often used as chemical "fingerprints" in analytical chemistry. The spectrum can represented by a plot of either absorbance* or transmittance* versus wavelength, frequency*, or wavenumber*.

angular momentum quantum number. (ell) azimuthal quantum number; orbital angular momentum quantum number.
A quantum number that labels the subshells* of an atom. Sometimes called the orbital angular momentum quantum number, this quantum number dictates orbital shape. ell can take on values from 0 to n-1 within a shell* with principal quantum number* n.

atomic orbital.
A wavefunction* that describes the behavior of an electron in an atom.

aufbau principle. aufbau construction; building-up principle.
An approximate procedure for writing the ground state* electronic configuration* of atoms. The configuration of an atom is obtained by inserting one electron into the configuration of the atom immediately to its left on the periodic table. The electron is inserted into the subshell indicated by the element's period* and block*.

Balmer series. Balmer lines.
A series of lines in the emission spectrum* of hydrogen* that involve transitions to the n=2 state from states with n>2.

band spectrum. band spectra. Compare with line spectrum* and continuous spectrum*.
An emission spectrum* that contains groups of sharp peaks that are so close together that they are not distinguishable separately, but only as a "band".

basis function.
A mathematical function that can be used to build a description of wavefunctions* for electrons in atoms or molecules.

basis set.
A set of mathematical functions that are combined to approximate the wavefunctions* for electrons in atoms and molecules.

Bohr atom. Bohr's theory; Bohr's atomic theory; Bohr model.
A model of the atom that explains emission and absorption of radiation as transitions between stationary electronic states in which the electron orbits the nucleus at a definite distance. The Bohr model violates the Heisenberg uncertainty principle, since it postulates definite paths and momenta for electrons as they move around the nucleus. Modern theories usually use atomic orbitals* to describe the behavior of electrons in atoms.

continuous spectrum. Compare with line spectrum* and band spectrum*.
A plot of the relative absorbance or intensity of emitted light vs. wavelength or frequency that shows a smooth variation, rather than a series of sharp peaks or bands.

core electron. Compare with valence electron*.
Electrons occupying completely filled shells* under the valence shell*.

degenerate. degenerate orbital.
A set of orbitals are said to be degenerate if they all have the same energy. This degeneracy can sometimes be "lifted" by external electric or magnetic fields.

diamagnetism. diamagnetic. Compare with paramagnetism*.
Diamagnetic materials are very weakly repelled by magnetic fields. The atoms or molecules of diamagnetic materials contain no unpaired spins*.

effective nuclear charge. (Zeff) Compare with atomic number*.
The nuclear charge experienced by an electron when other electrons are shielding the nucleus.

electron configuration. electronic configuration.
A list showing how many electrons are in each orbital or subshell*. There are several notations. The subshell notation lists subshells in order of increasing energy, with the number of electrons in each subshell indicated as a superscript. For example, 1s2 2s2 2p3 means "2 electrons in the 1s subshell, 2 electrons in the 2s subshell, and 3 electrons in the 2p subshell.

emission spectrum. emission spectra. Compare with absorption spectrum*.
A plot of relative intensity of emitted radiation* as a function of wavelength* or frequency*.

excited state. Compare with ground state*.
An atom or molecule which has absorbed energy is said to be in an excited state. Excited states tend to have short lifetimes; they lose energy either through collisions or by emitting photons* to "relax" back down to their ground states*.

f orbital. f-orbital.
An orbital* with angular momentum quantum number* ell = 2. The f orbitals generally have 3 nuclear nodes and rather complex shapes.

ground state. Compare with excited state*.
The lowest energy state for an atom or molecule. When an atom is in its ground state, its electrons fill the lowest energy orbitals completely before they begin to occupy higher energy orbitals, and they fill subshells in accordance with Hund's rule* (usually!)

Hund's rule. rule of maximum multiplicity.
A rule of thumb stating that subshells* fill so that the number of unpaired spins* is maximized, or "spread them out and line them up."

isoelectronic.
Refers to a group of atoms or ions having the same number of electrons. For example, F-, Ne, and Na+ are isoelectronic.

line spectrum. line spectra; line emission spectrum. Compare with band spectrum* and continuous spectrum*.
A emission spectrum* that contains very sharp peaks, corresponding to transitions between states in free atoms. For example, the line spectrum of hydrogen* contains 4 sharp lines in the visible part of the spectrum.

magnetic quantum number. (mell)
Quantum number that labels different orbitals within a subshell*. mell can take on values from -ell to +ell. The number of orbitals in a subshell is the same as the number of possible mell values.

noble gas core. ([X], where X is the symbol of an inert gas element) core configuration. Compare with valence shell*.
All completely filled shells* underneath the valence shell*.

orbital.
A wavefunction* that describes what an electron with a given energy is doing inside an atom or molecule.

paramagnetism. paramagnetic. Compare with diamagnetism* and ferromagnetism*.
Paramagnetic materials are attracted to a magnetic field due to the presence of least one unpaired spin* in their atoms or molecules.

Pauli principle. exclusion principle; Pauli exclusion; Pauli exclusion principle.
No two electrons in an atom can have the same set of 4 quantum numbers. Because the n, ell, and mell quantum numbers address a particular orbital, and because the ms quantum number has only two possible values, the Pauli principle says that a maximum of two electrons can occupy an atomic orbital- and these electrons must have opposite spins.

penetration. Compare with shielding*.
Electrons in penetrating orbitals can reach the nucleus. The n and ell quantum numbers determine how well an orbital penetrates. Lower n and lower ell values mean better penetration. A low n value means the orbital is small. A low ell value means the orbital has fewer nuclear nodes (planes that pass through the nucleus where the probability of locating the electron is zero).

In order of decreasing penetration, the subshells are s > p > d > f. A 1s orbital penetrates better than a 2s orbital.

principal quantum number. (n)
The quantum number that determines the size and (in hydrogen atoms) the energy of an orbital*. n is used to label electron shells*. n may take on integer values from 1 to infinity.

pseudocore.
Electrons in d or f subshells which are outside the noble gas core*.

shell. Compare with subshell*.
A set of electrons with the same principal quantum number*. The number of electrons permitted in a shell is equal to 2n2. A shell contains n2 orbitals*, and n subshells*.

shielding. Compare with penetration*.
Electrons in orbitals with high penetration* can shield the nucleus from less penetrating electrons. Because they are closer to the nucleus on average, they repel those farther away and lessen the effective nuclear charge* for the more distant electrons.

spectrophotometry. spectrophotometric.
Determination of the concentration of a material in a sample by measurement of the amount of light the sample absorbs.

spectroscopy. spectrometry; spectroscopic.
Spectroscopy is analysis of the interaction between electromagnetic radiation and matter. Different types of radiation interact in characteristic ways with different samples of matter; the interaction is often unique and serves as a diagnostic "fingerprint" for the presence of a particular material in a sample. Spectroscopy is also a sensitive quantitative technique that can determine trace concentrations of substances.

spectrum.
1. A sequence of colors produced by passing light through a prism or diffraction grating. 2. A range of wavelengths* of electromagnetic radiation. 3. A plot that shows how some intensity-related property of a beam of radiation or particles depends on another property that is related to dispersal of the beam by a prism, a magnet, or some other device. For example, a plot of light absorbance vs. wavelength is an absorption spectrum; a plot of ion abundance vs. mass is a mass spectrum.

spin.
Electrons have an intrinsic angular momentum that is similar to what would be observed if they were spinning. Electron spin is sometimes called a "twoness" property because it can have two values, referred to as "spin up" and "spin down". Nuclei can have spins of their own.

spin pair. (spinpair) paired spins; electron pair; paired electrons. Compare with unpaired spin*.
Two electrons with opposite spins*, usually occupying the same orbital.

subshell. sublevel.
A set of electrons with the same azimuthal quantum number*. The number of electrons permitted in a subshell is equal to 2ell + 1.

unpaired spin. (spinup) unpaired electron. Compare with paired spin*.
A single electron occupying an orbital*.

valence electron.
Electrons that can be actively involved in chemical change; usually electrons in the shell with the highest value of n. For example, sodium's ground state* electron configuration is 1s2 2s2 2p6 3s1; the 3s electron is the only valence electron in the atom. Germanium (Ge) has the ground state electron configuration 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p2; the 4s and 4p electrons are the valence electrons.

valence shell.
The shell* corresponding to the highest value of principal quantum number* in the atom. The valence electrons* in this shell are on average farther from the nucleus than other electrons; they are often directly involved in chemical reaction.



General Chemistry Online! Electrons in atoms

Copyright © 1997-2010 by Fred Senese
Comments & questions to fsenese@frostburg.edu
Last Revised 02/15/10.URL: http://antoine.frostburg.edu/chem/senese/101/electrons/glossary.shtml