Unit 3 Chemistry: Electron Configurations and Periodic Trends

Comprehensive vocabulary flashcards covering atomic structure, electron configurations, electromagnetic radiation, and periodic trends based on Unit 3 lecture review materials.

Electron

The negative subatomic particle in the electron cloud of an atom.

Electromagnetic radiation

A form of energy that exhibits wavelike behavior as it travels through space.

Wavelength

A unit of distance for measuring waves.

Frequency

The number of waves that pass a given point in a certain time.

Emission spectrum

The spectrum of light released from excited atoms of an element.

Heisenberg uncertainty principle

The principle stating that it is impossible to determine the position and velocity of an electron or other particle simultaneously.

Valence electrons

Electrons located in the outermost energy level of the electron cloud.

Octet rule

The tendency of atoms to prefer to have 8 valence electrons.

Ions

Atoms or molecules that have an electric charge due to having lost or gained electrons.

Lewis structures

A convenient and simple way to represent an element and its valence electrons.

Quantum Theory

A mathematical description for the wave properties of electrons and other tiny particles.

Orbital

A 3D region around the nucleus that shows where an electron probably is.

Electron configuration

The most stable arrangement of electrons in an atom’s electron cloud.

Periodic trends

Specific patterns in the periodic table that make it possible to make predictions about an element’s properties.

Coulomb’s law

A law stating that as the atomic number increases, the nuclear attraction for electrons increases, thus pulling the electrons in more tightly to the nucleus.

Atomic radius

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

Reactivity

A measure of how readily an atom will react.

Shielding effect

An effect where electrons in the inner energy levels block the attraction of the nucleus to the valence electrons.

Electronegativity

A measure of the ability of an atom in a chemical bond to attract electrons.

Ionization energy

The energy required to remove one electron from a neutral atom of an element.

Oxidation numbers

The charge of an ion.

Ionic radius

The radius of a monatomic ion in an ionic compound.

Electromagnetic spectrum

The full range of radiant energy—ordered from radiowave to gamma ray—displaying longer wavelengths, lower frequencies, and lower energies at the radiowave end, and shorter wavelengths, higher frequencies, and higher energies at the gamma ray end.

Ground state vs. Excited state transitions

An electron in its ground state absorbs energy to jump energy levels (A), and when returning to its ground state (B), it releases that energy in the form of electromagnetic radiation (C).

Cations

Atoms, usually metals, that lose electrons in order to become stable, resulting in a positive charge.

Anions

Atoms, usually nonmetals, that gain electrons in order to become stable, resulting in a negative charge.

Bohr model for Sulfur

A diagram depicting a sulfur nucleus containing $16\text{p}$ and $16\text{n}$, surrounded by three concentric electron energy levels holding 2, 8, and 6 electrons respectively.

Lewis structure for Sulfur

A Lewis dot diagram representing a neutral sulfur atom with the chemical symbol S surrounded by 6 valence electrons.

Aufbau principle

A fundamental rule stating that electrons fill the lowest energy orbital first.

Hund’s rule

A rule stating that orbitals of equal energy each get 1 electron before any orbital gets a second electron pair.

Pauli exclusion principle

A principle stating that no two electrons in the same atom can have the same set of quantum numbers.

Periodic trend: Atomic radius

The periodic pattern showing that atomic radius increases moving down a group and increases moving from right to left across a period.

Periodic trend: Reactivity

The periodic trend showing that reactivity increases down and to the left for metals, and increases up and to the right for nonmetals.

Periodic trend: Electronegativity

The periodic trend showing that electronegativity increases from left to right across a period and increases moving up a group.

Periodic trend: Ionization energy

The periodic trend showing that ionization energy increases from left to right across a period and increases from bottom to top up a group.

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

A measure used to predict bond type: if $\Delta\text{EN} > 2$, an ionic bond forms; if intermediate, a polar covalent bond forms; and if $\Delta\text{EN} < 0.5$, a covalent bond forms with mostly equal electron sharing.