Chemistry: Atomic Theory, Isotopes, and Atomic Mass

Vocabulary flashcards covering key terms from Dalton’s theory through isotopes and atomic mass.

Law of Definite Proportions

A given compound is composed of fixed, definite proportions by mass of its elements.

Proportion by Mass

The ratio of the mass of a specific element to the total mass of the compound.

Percent by Mass

Proportion by mass expressed as a percentage (element mass divided by total mass × 100%).

Law of Multiproportions

For compounds formed from the same elements, the masses of one element that combine with a fixed mass of the other are in small whole-number ratios.

Dalton’s Atomic Theory

1) Matter is made of atoms; 2) Atoms of a given element have the same mass; 3) Atoms combine to form molecules in small whole-number ratios; 4) Atoms of different elements can combine in different small whole-number ratios to form compounds.

Subatomic Particles

Protons, neutrons, and electrons—the fundamental particles that make up atoms.

Nucleus

The small, dense center of an atom containing protons and neutrons; contains most of the atom’s mass.

Electron

A negatively charged subatomic particle with very small mass that orbits the nucleus.

Proton

A positively charged subatomic particle located in the nucleus; mass about 1 amu.

Neutron

A neutrally charged subatomic particle located in the nucleus; mass about 1 amu.

Atomic Number Z

The number of protons in an atom’s nucleus; defines the element’s identity.

Mass Number A

The total number of protons and neutrons in an atom’s nucleus.

Isotopes

Atoms of the same element (same Z) with different numbers of neutrons and therefore different masses.

Atomic Mass Unit (amu/u)

Relative unit for atomic masses; 1 amu is defined in relation to a standard isotope (1/12 the mass of a C-12 atom in the modern scale).

C-12 Standard

In 1961, carbon-12 was designated as the standard with a mass of exactly 12 u.

Weighted Average Atomic Mass

The atomic mass shown on the periodic table, the weighted average of the masses of naturally occurring isotopes based on their abundances.

Rutherford’s Gold Foil Experiment

Experiment showing most alpha particles pass through a thin gold foil with some deflections, indicating a tiny, dense nucleus surrounded by mostly empty space.

Cathode Ray Tube

Device used by Thomson to produce a beam of electrons; demonstrated electron properties via deflection in electric and magnetic fields.

Plum-Pudding Model

Thomson’s model of the atom: electrons embedded in a positively charged sphere.

Rutherford’s Nuclear Model

Model proposing a small, dense nucleus containing protons (and later neutrons) with electrons scattered in the surrounding mostly empty space.

Millikan Oil-Drop Experiment

Experiment used to determine the charge-to-mass ratio of the electron, enabling calculation of electron mass.

Chadwick (Neutrons)

In 1932, James Chadwick demonstrated the existence of neutrons in the nucleus.