Unit 1: Development of the Atomic Theory

Vocabulary practice flashcards based on Unit 1: Development of the Atomic Theory notes.

Aristotle's View of Matter

The philosophy that all physical objects combine matter (potentiality) and form (identity), and matter is composed of four elements (earth, air, water, fire) defined by opposing qualities (hot/cold and wet/dry).

Democritus's Atomic Theory

The hypothesis that the universe consists of tiny, uncuttable, indestructible, and eternal particles called atoms (atomos) and void (empty space), moving constantly and differing in shape, size, and mass.

Dalton's Atomic Theory

A 1803-1807 theory stating that elements are composed of identical atoms, atoms of different elements differ in properties, chemical reactions rearrange atoms without creating or destroying them, and compounds form in fixed ratios of atoms.

Cathode Ray Experiment

An 1897 experiment by J. J. Thomson where rays moved toward positive charges, leading to the discovery of the electron and the calculation of its charge-to-mass ratio.

Sir Joseph John "J. J." Thomson

Scientist credited with the discovery of the electron in 1897, calculating its charge-to-mass ratio, and proposing the Plum-Pudding Model.

Plum-Pudding Model

Atomic model proposed by Thomson picturing the atom as a uniform positive sphere of matter with negative electrons embedded inside.

Oil-drop Experiment

An experiment (1908-1917) conducted by Robert Millikan that determined the electric charge of an electron, allowing its mass to be calculated.

Gold Foil Experiment

A 1910 experiment by Ernest Rutherford where alpha particles were shot at gold foil; most passed through but some deflected, proving the existence of a dense positive nucleus.

Ernest Rutherford

Scientist credited with discovering the nucleus in 1910 via the Gold Foil Experiment and later characterizing and naming the proton in 1919.

Eugen Goldstein

Scientist who discovered the existence of a positive charge in the atom in 1886 using a modified cathode ray tube.

Sir James Chadwick

Scientist who discovered neutrons in 1932 by bombarding a Beryllium plate with alpha particles, releasing neutral particles with mass slightly greater than a proton.

Niel's Bohr

Scientist who proposed (1911-1918) that electrons revolve in permitted circular orbits with definite energies, emitting or absorbing photons when transitioning between states.

Groups

Vertical columns on the periodic table containing elements with similar physical and chemical properties.

Periods

Horizontal rows on the periodic table across which physical and chemical properties change, indicating the number of energy levels in an atom.

Alkali Metals

Elements located in Group 1 of the periodic table (excluding hydrogen).

Alkaline Earth Metals

Elements located in Group 2 of the periodic table.

Transition Metals

Elements located in Groups 3 through 12 of the periodic table.

Chalcogens

Elements located in Group 16 of the periodic table.

Halogens

Elements located in Group 17 of the periodic table.

Noble Gases

Elements located in Group 18 of the periodic table.

Representative Elements

Elements in Groups 1, 2, and 13 through 18 on the periodic table.

Lanthanide Series

Series of elements on the periodic table with atomic numbers 58 through 71.

Actinide Series

Series of elements on the periodic table with atomic numbers 90 through 103.

Proton

Subatomic particle located in the nucleus with a positive ($+1$) electric charge and an approximate mass of $1.0073\,amu$.

Neutron

Subatomic particle located in the nucleus with no electric charge (neutral) and an approximate mass of $1.0087\,amu$.

Electron

Subatomic particle located in the electron cloud with a negative ($-1$) electric charge and a negligible mass of $0.0005486\,amu$.

Atomic Number (Z)

The whole number on the periodic table representing the total number of protons in an atom's nucleus.

Mass Number (A)

A whole number equal to the sum of protons and neutrons in an atom ($A = \text{protons} + \text{neutrons}$).

Isotopes

Two or more atoms of the same element that have equal numbers of protons but different numbers of neutrons.

Ion

An atom or former atom that has gained or lost electrons, resulting in a net positive or negative electric charge.

Cation

A positively charged ion formed when an atom loses electrons.

Anion

A negatively charged ion formed when an atom gains electrons.

Mass Spectrometry

An analytical technique that separates ions by mass-to-charge ratio to produce spectra graphing isotopic mass versus relative abundance.

Average Atomic Mass

The weighted average mass of all naturally occurring isotopes of an element, calculated using $\text{Average Atomic Mass} = \frac{\sum (\text{mass}_i \times \text{abundance}_i)}{100}$.

Carbon-12

The standardized reference isotope assigned an exact mass of $12\,amu$, against which all other atomic masses are compared.