Physical Science Chapter 4: The Periodic Table and Atomic Structure

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

Robert Boyle

Began the transition from alchemy to chemistry, believed it was possible to change one element into another, rejected the Greek concept of matter as mixtures of the five basic elements, and published The Skeptical Chemyst (1661) calling for experimentation to be the basis of science.

Antoine Lavoisier

Demonstrated that combustion requires oxygen and therefore that fire cannot be an element (1779), and compiled one of the first lists of known elements, which included 33 elements known in his day.

Jacob Berzelius

Developed a system of symbols for naming elements using 1–2 letters of each element's Latin name to ease information sharing, combined element symbols for compounds, and used superscripts to represent atom quantities (where subscripts are used today).

Modern Element Symbols

Symbols made up of one or two letters where the first letter is always capitalized and the second letter is always lower-case; they sometimes originate from other languages, mostly Latin (e.g., W for tungsten from the German name wolfram).

Johann Döbereiner

Noticed that when sorted by their masses, elements appear in groups of three with similar properties (which he called triads), setting the stage for discovering periodicity.

Triads

Groups of three elements with similar properties identified by Johann Döbereiner when elements were sorted by their masses.

Periodicity

The repetition of the properties of the elements in a regular or periodic fashion when the elements are ordered by atomic mass or atomic number.

John Newlands

Arranged elements in order by atomic mass, discovered similar pairs of elements positioned every eight elements, developed the law of octaves (1865), and is credited with the earliest form of the periodic law.

Law of Octaves

A principle formulated by John Newlands in 1865 stating that every eighth element exhibits similar properties when elements are ordered by atomic mass.

Dmitri Mendeleev

Organized elements in a table in columns by increasing atomic masses (1869) and accurately predicted future discoveries by leaving blank spaces where expected elements were missing according to the periodic law.

Henry Moseley

Discovered a way to determine the atomic number for each element (1914), leading to the arrangement of elements by increasing atomic number in rows rather than atomic mass, an updated periodic law, and the modern periodic table.

Periodic Law

The law that states that the properties of the elements vary with their atomic numbers in a regular, repeated pattern.

Periodic Table of the Elements

A table of the chemical elements arranged on the basis of atomic structure, which displays their periodic properties in relation to their atomic numbers.

Period

A row in the periodic table of the elements; also called a series. Designates the energy level of valence electrons.

Family

A column of elements in the periodic table having similar valence electron arrangement, resulting in similar chemical and physical properties; also known as a group.

Atomic Number ($Z$)

The number of protons and electrons in a neutral atom.

Ground State

The lowest and most stable energy state an atom can have.

Octet Rule

The arrangement rule stating that the maximum number of electrons an atom can have in its outermost energy level is eight, which is usually the most stable arrangement of electrons.

$2n^2$ Rule

The formula used to determine the maximum number of electrons that can occupy a given energy level $n$ (level 1 holds 2, level 2 holds 8, level 3 holds 18, level 4 holds 32).

Valence Electron

An electron in the outermost energy level of a neutral atom.

Metal

An element that is typically dense, solid, ductile, malleable, highly conductive, lustrous, and chemically reactive (especially with nonmetals); located to the left of the heavy stair-step line, making up almost 80% of the elements and typically possessing one or two valence electrons.

Metalloid

An element with characteristics between those of metals and nonmetals (also called semi-conductors) whose conductivity increases with heat; a brittle solid with metallic luster located along the stair-step line.

Nonmetal

An element that typically has four or more valence electrons, does not exhibit the general properties of metals, and exists as a gas, liquid, or dull, brittle solid; located to the right of, but not touching, the stair-step line.

Alkali Metals

Elements in Group 1 of the periodic table having one valence electron that they easily lose to form a $1+$ cation; they are extremely reactive, the most reactive of all metals, and never found in nature in their pure form.

Alkaline-Earth Metals

Elements in Group 2 of the periodic table having two valence electrons that tend to be lost easily to become a $2+$ cation, making them very reactive.

Transition Metals

Any elements in Groups 3–12 of the periodic table, typically having one or two valence electrons which are easily lost, resulting in cations with charges of $1+$ or $2+$.

Inner Transition Metals

Elements from either of the two rows usually placed below the periodic table to save space; members of the lanthanide or actinide series that typically have two valence electrons.

Mixed Groups

Any of Groups 13–16 in the periodic table, so named because they contain metals, nonmetals, and metalloids; often named for the first element in the family.

Halogen

A nonmetal in Group 17 of the periodic table having seven valence electrons; easily gains an electron to form a $1-$ anion, which causes them to be highly reactive.

Noble Gas

An element in Group 18 on the periodic table having eight valence electrons filling the outer energy level (except helium, which has two); with a full outer energy level, each noble gas is inert (nonreactive).

Atomic Radius

The distance from the center of an atom's nucleus to its outermost energy level; gets smaller moving across a period to the right due to opposite charge attraction, and gets larger with more energy levels down a group.

Electronegativity

A measure of an element's ability to attract and hold electrons when bonded to other atoms; increases as atomic radius decreases because of the strong attraction exerted on the electrons by the nucleus.

Electron Dot Notation

A representation of an element's ability to attract and hold electrons when bonded to other atoms due to their valence electrons and reactivity.