Periodic Table and Periodic Properties

Flashcards covering the historical background, main features, groups, families, and periodic properties of elements in the modern periodic table, including bonding and reactions of sodium and magnesium.

Antoine Lavoisier

The scientist who made the first classification of elements in the 18th century, attempting to group them as metals and nonmetals.

Triads

A classification developed by Döbereiner in 1829 where elements were grouped in threes with similar properties, such that the atomic weight of the middle element was roughly the average of the other two.

John Newlands

The English chemist who, in 1864, first observed periodicity by arranging 62 elements in increasing order of atomic masses, noting that every eighth element resembled the first.

Dmitri Mendeleev

Known as the father of the Periodic Table, he arranged 63 elements by increasing atomic mass in 1869, leaving gaps for undiscovered elements and accurately predicting their properties.

Moseley

In 1913, he determined the exact atomic numbers of elements using X-ray emission, leading to the arrangement of the Periodic Table by atomic numbers instead of atomic masses.

Modern Periodic Law

States that the physical and chemical properties of elements are periodic functions of their atomic numbers.

Periods

The seven horizontal rows in the modern periodic table.

Groups

The eighteen vertical columns in the modern periodic table.

Metals

Elements which tend to lose electrons to form positive ions, such as iron, copper, gold, and silver.

Non-metals

Elements which tend to gain electrons to form negative ions, such as chlorine, sulfur, and phosphorous.

Metalloids

Also known as semimetals, these elements exhibit properties of both metals and non-metals and are located along a "stair-step line" starting at boron ($B$) and extending to polonium ($Po$).

s-block

Elements in the first two groups of the periodic table whose valence electrons are in the "s" subshells.

p-block

The elements in groups 13 to 18, including the inert gases, characterized by valence electrons in the "p" subshells.

d-block

Transition elements located in the middle of the periodic table.

f-block

Comprised of the Lanthanides and Actinides located in the two rows at the bottom of the periodic table.

Alkali metals

Group 1 elements ($Li$, $Na$, $K$, $Rb$, $Cs$, $Fr$) which have one valence electron and produce alkalis when they react with water.

Alkaline earth metals

Group 2 elements ($Be$, $Mg$, $Ca$, $Sr$, $Ba$, $Ra$) which have two electrons in their valence shell and are primarily found in the earth.

Chalcogens

The group 16 elements ($O$, $S$, $Se$, $Te$, $Po$, $Lv$), named because most ores of copper are oxides or sulfides.

Halogens

The group 17 elements ($F$, $Cl$, $Br$, $I$, $At$, $Ts$), which means "salt-former" because they react with metals to form salts.

Noble Gases

The unreactive elements in Group 18 ($He$, $Ne$, $Ar$, $Kr$, $Xe$, $Rn$, $Og$) that have complete outermost shells.

Atomic Radius

Half of the distance between two identical atoms bonded together, typically measured in picometer ($pm$) or Angstrom ($\text{Å}$).

Shielding Effect

The effect where inner electrons reduce the attraction between the nucleus and the outermost electrons, contributing to an increase in atomic size down a group.

Ionic Radius

The measure of the size of an ion in a crystal lattice, defined as the distance from the nucleus to the outermost electron shell.

Ionization Energy ($ΔH_{i1}$)

The energy needed to remove one electron from each atom in one mole of atoms of an element in the gaseous state.

Electron Affinity ($ΔH_{ea}$)

The enthalpy change involved when 1 mole of electrons is added to 1 mole of gaseous atoms to form 1 mole of gaseous uni-negative ions.

Electronegativity

The power of an atom to attract a shared pair of electrons toward itself in a molecule.

Metallic Character

The tendency of elements to lose electrons; this character decreases across a period and increases down a group.

Basic oxide

An oxide that gives off an alkali when combined with water, usually formed by metals in Groups 1 and 2, such as $Na_2O$ and $CaO$.

Acidic oxide

An oxide that gives off an acid when combined with water, typically formed by non-metals, such as $SO_2$ or $P_2O_3$.

Amphoteric oxides

Oxides that can react with both acids and bases, such as aluminum oxide ($Al_2O_3$).

Hydrolysis

The process by which certain chlorides, like $AlCl_3$ or $SiCl_4$, react with water to form acidic solutions.

Oxidation Number

The formal charge on an atom in a molecule, representing the charge that appears on ions in ionic compounds.