Electron Configuration, Elements, and Periodic Trends

Practice vocabulary flashcards covering electron configurations, orbital diagrams, the historical development of the periodic table, element classifications, and periodic trends.

Electron Configuration

A notation using an element's symbol along with a sequence of numbers and letters to show how electrons are arranged and distributed in an atom or ion.

Orbitals

Specific regions within sublevels of an energy level where electrons exist, with each orbital capable of holding a maximum of $2$ electrons with opposite spins.

Aufbau Principle

The building-up principle stating that electrons fill the lowest energy orbitals first before occupying higher energy orbitals.

Pauli Exclusion Principle

The rule stating that an orbital can hold a maximum of two electrons, and these electrons must have opposite (spin-up and spin-down) spins.

Hund's Rule

The rule stating that every orbital in a given subshell is occupied by one electron before any orbital is doubly occupied with a second electron.

Valence Electrons

The electrons found in the highest energy level (outermost shell) of an atom that govern its chemical properties and bonding behavior.

Orbital Diagram

A visual representation of electron distribution where orbitals within a sublevel are drawn as squares or circles and electrons are represented as upward or downward arrows.

$s$ Sublevel

A sublevel containing $1$ orbital that can accommodate a maximum of $2$ electrons.

$p$ Sublevel

A sublevel containing $3$ orbitals that can accommodate a maximum of $6$ electrons.

$d$ Sublevel

A sublevel containing $5$ orbitals that can accommodate a maximum of $10$ electrons.

$f$ Sublevel

A sublevel containing $7$ orbitals that can accommodate a maximum of $14$ electrons.

Chemical Symbol

A shorthand representation of an element's name consisting of one or two letters, officially assigned by the International Union of Pure and Applied Chemistry (IUPAC).

Law of Triads

An organizational scheme proposed in 1817 by Johann Wolfgang Döbereiner where elements with similar chemical properties are grouped in sets of three, with the middle element's atomic weight roughly averaging the other two.

Telluric Screw

A 3D spiral cylinder model created in 1862 by Alexandre-Émile Béguyer de Chancourtois, marking the first arrangement of elements by increasing atomic weight.

Alexandre-Émile Béguyer de Chancourtois

The French scientist who in 1862 was the first to arrange chemical elements in order of increasing atomic weight around a cylinder known as the telluric screw.

Law of Octaves

An 1863 concept proposed by John Newlands stating that when elements are arranged by increasing atomic weight, similar physical and chemical properties recur every eighth element.

Dmitri Mendeleev

The Russian chemist known as the father of the periodic table who arranged 60 known elements by increasing atomic mass and left blank spaces predicting undiscovered elements.

Lothar Meyer

The German chemist who independently produced an arrangement of elements by increasing atomic mass in 1869, very similar to Mendeleev's periodic table.

Henry Moseley

The English scientist who in 1913 experimentally discovered that elements should be organized by increasing atomic number rather than atomic weight, revising the Periodic Law.

Periods

The seven horizontal rows of the periodic table, where each row number matches the principal energy level ($n$) of the valence electrons.

Groups

The 18 vertical columns of the periodic table containing elements that possess the same number of valence electrons and exhibit similar properties.

Alkali Metals

Group 1 elements characterized by soft textures, silvery appearance, low melting and boiling points, high reactivity, and one valence electron that readily forms $+1$ cations.

Alkaline Earth Metals

Group 2 soft and silvery metallic elements that have two valence electrons, commonly used in applications like cement (calcium) and alloys or fireworks (magnesium).

Transition Metals

Elements in Groups 3 through 12 ($d$-block and $f$-block) that display high melting points, high electrical conductivity, multiple oxidation states, colored compounds, and effective catalytic properties.

Boron Group

Group 13 elements with three valence electrons and a general valence electron configuration of $ns^2 np^1$, containing the metalloid boron alongside metals.

Carbon Group

Group 14 elements possessing four valence electrons with a general electron configuration of $ns^2 np^2$, including carbon which serves as the backbone of organic compounds.

Nitrogen Group

Group 15 elements with five valence electrons and a general configuration of $ns^2 np^3$, featuring nitrogen, phosphorus, and antimony.

Chalcogens

Group 16 elements (also known as the oxygen group) possessing six valence electrons with a general electron configuration of $ns^2 np^4$.

Halogens

Highly reactive Group 17 nonmetallic elements that possess seven valence electrons ($ns^2 np^5$) and react with metals to produce salts.

Noble Gases

Group 18 elements that are odorless, colorless, and unreactive due to their stable electron configuration of eight valence electrons (except helium, which has two).

Periodic Trends

Specific recurring patterns observed across periods and down groups in the periodic table describing properties such as atomic size, ionization energy, and electronegativity.

Atomic Radius

One-half the distance between the nuclei of two adjacent identical atoms, which decreases from left to right across a period and increases from top to bottom down a group.

Metallic Character

A measure of the tendency of an element to lose electrons and form cations, increasing down a group and decreasing across a period.

Ionization Energy

The quantity of energy needed to remove an electron from a neutral atom in its gaseous state, which increases across a period from left to right and decreases down a group.

Electronegativity

The relative ability of an atom to attract bonding electrons, increasing from left to right across a period and increasing from bottom to top within a group.

Electron Affinity

The amount of energy released when a gaseous atom gains an electron to form a negative ion, increasing across a period and decreasing down a group.