Chapter 4: Atomic Structure

Vocabulary practice flashcards covering Chapter 4: Atomic Structure, including historical atomic models, subatomic particles, atomic notation, ions, and isotopes.

Democritus

A Greek philosopher (460–370 BC) who was the first to use the word atom and believed atoms were indivisible and indestructible, though he lacked scientific support.

Aristotle

A Greek philosopher (384–322 BC) who rejected the theory of atoms and believed matter was composed of four core elements: fire, air, water, and earth.

Law of Conservation of Matter

A fundamental law proposed by Antoine Lavoisier stating that matter is not created or destroyed during a chemical reaction.

Law of Definite Proportions

A chemical law established by Joseph Proust stating that a compound is always composed of the same elements in the same proportion by mass.

Law of Multiple Proportions

A law formulated by John Dalton stating that when two elements form more than one compound, they combine in ratios of small whole numbers.

Dalton's Atomic Theory

A theory stating that all elements are composed of indivisible 'solid sphere' atoms, all atoms of a given element are identical, atoms of different elements differ in mass, and atoms combine in simple whole-number ratios to form compounds.

Subatomic Particles

The three smaller components into which divisible atoms are split: protons, neutrons, and electrons.

Cathode-Ray Tube

A vacuum tube apparatus where electrons travel as a ray from the negative metal disk (cathode) to the positive metal disk (anode), used by J.J. Thomson to identify the electron.

Electron

The first identified subatomic particle, discovered by J.J. Thomson in 1897; has the symbol $e^-$, a charge of $-1$, and a mass of approximately $0\,\text{amu}$.

Plum Pudding Model

J.J. Thomson's atomic model (also described as the chocolate chip cookie model) proposing an atom composed of negative electrons embedded within a spherical positive matrix.

Proton

A subatomic particle located in the nucleus with the symbol $p^+$, a charge of $+1$, and a mass of $1\,\text{amu}$.

Neutron

A neutral subatomic particle located in the nucleus with the symbol $n^0$, a charge of $0$, and a mass of $1\,\text{amu}$.

Gold Foil Experiment

Ernest Rutherford's 1911 experiment directing alpha particles at gold foil, revealing that atoms are mostly empty space with a dense, positively charged nucleus.

Nucleus

The tiny, dense, positively charged center of an atom containing protons and neutrons, which accounts for most of the atom's mass.

Atomic Number

The number of protons in an atom's nucleus, written as a subscript, which determines an element's unique identity.

Nuclear Charge

The total amount of positive charge in an atom's nucleus, which is numerically equivalent to the atomic number (number of protons).

Mass Number

The total mass of the nucleus, calculated as the sum of protons and neutrons ($A = Z + N$), written as a superscript or hyphenated after the element symbol.

Isotope Notation

A symbolic representation written in the format $^A_Z X$, where $X$ is the chemical symbol, $A$ is the mass number, and $Z$ is the atomic number.

Ion

A charged atom produced when an atom gains or loses electrons during a chemical reaction, resulting in an unequal number of protons and electrons.

Cation

A positively charged ion formed when an atom loses one or more electrons via oxidation (e.g., $\text{Na}^{1+}$, $\text{Ca}^{2+}$, $\text{Al}^{3+}$).

Anion

A negatively charged ion formed when an atom gains one or more electrons via reduction (e.g., $\text{F}^{1-}$, $\text{O}^{2-}$, $\text{P}^{3-}$).

Oxidation

A chemical reaction process in which an atom loses electrons, represented by the mnemonic 'LEO' (Lose Electrons Oxidation).

Reduction

A chemical reaction process in which an atom gains electrons, decreasing in numerical oxidation state, represented by the mnemonic 'GER' (Gain Electrons Reduction).

Isotope

Naturally occurring atoms of the same element that possess identical numbers of protons but different numbers of neutrons and different masses.

Average Atomic Mass

The weighted average mass of all naturally occurring isotopes of an element, calculated using the formula: $\sum (\text{abundance} \times \text{mass})$ and expressed in units of $\text{amu}$.