Atomic Structure and Periodicity

Flashcards on Atomic Structure and Periodicity

Democritus

Proposed that the world was made of empty space and fine, indivisible particles called atomos

Aristotle

Proposed that matter is a continuum composed of earth, air, fire, and water.

Ludacris

Described matter as bodies composed of empty space that allows movement

John Dalton

Refined the atomic view of matter through his atomic theories

Dalton's Atomic Theory (1)

Each element is composed of extremely small particles called atoms

Dalton's Atomic Theory (2)

All atoms of a given element are identical, but differ from other elements.

Dalton's Atomic Theory (3)

Atoms of one element cannot be changed into atoms of a different element.

Dalton's Atomic Theory (4)

A given compound always has the same relative number and kind of atoms.

Law of Definite Composition

States that in a given compound, the kinds and relative numbers of atoms are constant.

Law of Conservation of Mass

States that the total mass of the materials present after a chemical reaction is the same as the total mass before the reaction.

Law of Multiple Proportions

States that when two or more elements combine to form more than one compound, they combine in a ratio of small whole numbers.

Protons

Discovered in 1896 by Eugene Goldstein.

Electrons

Discovered in 1897 by JJ Thomson.

Neutrons

Discovered in 1932 by James Chadwick.

X-rays

Discovered in 1895 by Wilhelm Roentgen.

Radioactivity

Discovered in 1896 by Antoine Henri Becquerel.

Plum-pudding Atomic Model

Proposed by JJ Thomson, states that an atom is made up of negatively-charged electrons embedded in a nebulous cloud of positive charges.

Nuclear Atomic Model

Proposed by Ernest Rutherford, states an atom has a dense center of positive charge (the nucleus) from which electrons move around.

Planetary Model of the Atom

Proposed by Neils Bohr, suggests that electrons move in a path of definite amount of energy around the nucleus.

Quantum Mechanical Model

Developed by Erwin Schrodinger, Werner Heisenberg, and Louis de Broglie, states that electrons move at various energy levels with definite amount of energy or quanta.

Atomic number (Z)

Fingerprint of an atom; gives the element’s unique number of protons.

Mass number (A)

Gives the total number of protons and neutrons.

Isotopes

Atoms with the same number of protons but different number of neutrons.

Atomic Mass

The average atomic mass of each element is the sum of the individual isotopes and their corresponding abundance.

Quantum Numbers

Describes the designation of how electrons are distributed among various orbitals in principal shells and subshells.

Shell

Each division of space around the nucleus where electrons travel (main energy levels).

Orbital

A particular region in space around the nucleus where the probability of finding the electron is greatest.

Principal Quantum Number (n)

Gives the main energy and size of an orbital.

Azimuthal Quantum Number (l)

Defines the shape of the atomic orbital and comprises the sublevels of the principal quantum number.

Magnetic Quantum Number (ml)

Describes the orientation of the degenerate orbitals.

Spin Quantum Number (ms)

Defines the orientation of the electron in an orbital (+1/2 or -1/2).

Electronic Configuration

A designation of how orbitals are filled with electrons.

Valence electrons

The outermost electrons or those on the highest energy level.

Core electrons

The rest of the electrons other than the valence electrons.

Expanded form (electronic configuration)

Identifies all the electrons of the atom.

Abbreviated form (electronic configuration)

Shows only the noble gas element that is isoelectronic with the configuration of the core electrons and the valence electrons.

Aufbau Principle

In filling orbitals, orbitals with the lowest energy are filled first.

Hund's Rule of Maximum Multiplicity

In filling degenerate orbitals, each orbital is half-filled with one electron before any are filled with two.

Pauli’s Exclusion Principle

Each orbital is filled with electrons of opposite spins; no two electrons can have the same set of quantum numbers.

Triads

Groups of three elements with similar properties

John Newlands

Expanded the group into eight elements (or octaves) with similar properties

Dmitri Mendeleev

Arrangement based on increasing atomic weight.

Henry Moseley

Proposed that physical and chemical properties of elements vary periodically with increasing atomic number

Alkali Metals

Group 1A

Alkaline-Earth Metals

Group 2A

Noble Gasses

Halogens

Atomic Radius

Half the distance between two bonding nuclei.

Ionic Radius

Radius when an atom loses electron/s to become a cation or accepts electron/s to become an anion.

Ionization Energy (IE)

Minimum energy required to release an electron from a gaseous atom or ion to become a cation.

Electron Affinity (EA)

Energy change associated with the addition of an electron to a gaseous atom.

Electronegativity

Ability of an atom to attract a shared pair of electrons to itself.