Classification of Elements and Periodicity in Properties

Comprehensive vocabulary flashcards covering the historical development, structural organization, blocks, and periodic trends of the chemical elements based on the provided lecture transcript.

Chemical Elements

The basic units of all types of matter and the fundamental building blocks of chemistry.

Law of Triads

A relationship noted by Johann Dobereiner in 1829 where the middle element of a group of three had an atomic weight about halfway between the other two, and properties in between them.

Law of Octaves

Proposed by John Alexander Newlands in 1865, stating that when elements are arranged in increasing order of atomic weights, every eighth element has properties similar to the first.

Mendeleev’s Periodic Law

The principle stating that the properties of the elements are a periodic function of their atomic weights.

Eka-Aluminium

The name given by Mendeleev to the then-undiscovered element later identified as Gallium.

Eka-Silicon

The name given by Mendeleev to the then-undiscovered element later identified as Germanium.

Modern Periodic Law

The principle stating that the physical and chemical properties of the elements are periodic functions of their atomic numbers.

Atomic Number ($Z$)

A fundamental property of an element equal to the nuclear charge ($\text{number of protons}$) or the number of electrons in a neutral atom.

Periods

The horizontal rows in the Periodic Table, where the period number corresponds to the highest principal quantum number ($n$) of the elements.

Groups

The vertical columns in the Periodic Table (also called families) containing elements with similar outer electronic configurations.

IUPAC Numerical Roots (0-9)

The roots used for systematic nomenclature of elements with $Z > 100$: 0=nil, 1=un, 2=bi, 3=tri, 4=quad, 5=pent, 6=hex, 7=sept, 8=oct, 9=enn.

Electronic Configuration

The distribution of electrons into orbitals ($s, p, d, f$) of an atom.

s-Block Elements

Elements of Group 1 ($\text{alkali metals}$) and Group 2 ($\text{alkaline earth metals}$) with outermost electronic configurations of $ns^1$ and $ns^2$.

p-Block Elements

Elements belonging to Groups 13 to 18 characterized by the filling of p orbitals, with an outermost configuration varying from $ns^2 np^1$ to $ns^2 np^6$.

Representative Elements

The elements comprising the s-block and p-block (except Group 18 in some contexts), also known as Main Group Elements.

d-Block Elements (Transition Elements)

Elements of Group 3 to 12 characterized by the filling of inner d orbitals, with the general outer configuration $(n-1)d^{1-10} ns^{0-2}$.

f-Block Elements (Inner-Transition Elements)

The Lanthanoids and Actinoids located at the bottom of the Periodic Table, characterized by the filling of f orbitals and a general configuration of $(n-2)f^{1-14} (n-1)d^{0-1} ns^2$.

Transuranium Elements

The chemical elements with atomic numbers greater than 92 (the atomic number of uranium).

Metalloids (Semi-metals)

Elements such as silicon, germanium, arsenic, antimony, and tellurium that Border the zig-zag line in the periodic table and show properties of both metals and non-metals.

Covalent Radius

Half the distance between two atoms when they are bound together by a single bond in a covalent molecule.

Metallic Radius

Half the internuclear distance separating the metal cores in a metallic crystal.

Isoelectronic Species

Atoms and ions that contain the same number of electrons, such as $O^{2-}$, $F^-$, $Na^+$, and $Mg^{2+}$.

Ionization Enthalpy ($\triangle_i H$)

The energy required to remove an electron from an isolated gaseous atom in its ground state, expressed in $kJ \text{ mol}^{-1}$.

Shielding (Screening) Effect

The reduction in the effective nuclear charge experienced by valence electrons due to the presence of intervening core electrons.

Electron Gain Enthalpy ($\triangle_{eg}H$)

The enthalpy change that occurs when an electron is added to a neutral gaseous atom to convert it into a negative ion.

Electronegativity

A qualitative measure of the ability of an atom in a chemical compound to attract shared electrons to itself.

Valence

The combining capacity of an element, usually equal to the number of electrons in the outermost orbitals or eight minus that number.

Diagonal Relationship

The similarity in properties between the first element of a group (like lithium or beryllium) and the second element of the following group (like magnesium or aluminium).

Amphoteric Oxides

Oxides that behave as acidic with bases and as basic with acids, such as $Al_2O_3$ and $As_2O_3$.

Neutral Oxides

Oxides that have no acidic or basic properties, such as $CO$, $NO$, and $N_2O$.