Chem 201 Fall 2025 - Worksheet 1 Key Terms (VOCABULARY Flashcards)

Vocabulary flashcards covering key concepts from the notes: subatomic particles, atomic structure, isotopes, ions, historical models, Avogadro’s number, mole, molar mass, amu, isotopic abundances, and representative calculations.

Particulate model of matter

Idea that matter is composed of small particles (atoms/molecules) whose arrangement and motion create the different states of matter (solids, liquids, gases).

Phase of matter

One of the states of matter (solid, liquid, or gas) characterized by particle arrangement and energy level.

Proton

Positively charged subatomic particle located in the nucleus; mass ≈ 1 amu; its number determines the element (atomic number).

Neutron

Electrically neutral subatomic particle in the nucleus; mass ≈ 1 amu; contributes to the atom’s mass but not its charge.

Electron

Negatively charged subatomic particle surrounding the nucleus; mass is about 1/1836 of a proton; determines chemical behavior and charge balance in atoms.

Atomic number (Z)

Number of protons in an atom’s nucleus; defines the element and equals the number of electrons in a neutral atom.

Mass number (A)

Total number of protons and neutrons in the nucleus; A = Z + N.

Isotope

Atoms of the same element (same Z) that have different mass numbers due to different numbers of neutrons.

Ion

Atom or molecule with a net electric charge due to loss or gain of electrons.

Nuclear atom (Rutherford model)

Model in which a dense, positively charged nucleus contains most of the atom’s mass, surrounded by electrons.

Plum pudding model

Thomson’s early atomic model where electrons were embedded in a uniformly positive sphere.

Oil drop experiment

Experiment by Millikan to measure the charge of the electron (and determine e/m).

Charge-to-mass ratio (e/m) of the electron

The ratio of the electron’s charge to its mass, used to determine properties of the electron.

Avogadro’s number

6.022 × 10^23; the number of particles in one mole of a substance.

Mole

Amount of substance containing Avogadro’s number of particles (symbol: mol).

Molar mass

Mass of one mole of a substance (units: g/mol); numerically equal to the atomic/molecular mass in amu.

Atomic mass unit (amu)

1/12 the mass of a carbon-12 atom; unit used to express atomic and molecular masses.

Isotopes of carbon (12C and 13C)

Two common carbon isotopes differing in neutron number but with the same proton number (Z).

12C+ and 13C-

Ionized forms of carbon-12 and carbon-13 (positive for 12C+ and negative for 13C-).

Dozen

A quantity equal to 12 items.

Dozen vs mole (conceptual)**

A dozen is 12 items; a mole is 6.022 × 10^23 items, linking mass to particles via Avogadro’s number.

Molar masses (examples)

He: ~4.00 g/mol; K: ~39.10 g/mol; Cl: ~35.45 g/mol; H: ~1.008 g/mol—the mass per mole of each element.

Chlorine: atomic number and average mass

Atomic number 17 (protons/electrons in neutral Cl); average atomic mass ≈ 35.453 amu (mass per mole ≈ 35.453 g).

Rhenium isotopes

Natural Re contains isotopes 185Re and 187Re; their abundances determine the average atomic mass.

Boron average atomic mass

Weighted average mass of boron’s isotopes (e.g., 10B, 11B, 12B) based on natural abundances, yielding ~10.81 amu.

Neon isotopes

Neon has significant amounts of 20Ne and 22Ne; 20Ne mass ≈ 19.9924 amu with a large natural abundance (≈90.5%).

Chlorine natural isotopes and abundance

Natural chlorine consists mainly of 35Cl and 37Cl with characteristic abundances; contributes to the average atomic mass 35.453 amu.

Mass of one mole of helium atoms

Approximately 4.00 g (helium’s molar mass).

Molar mass vs atomic mass unit (amu)

Molar mass (g/mol) is numerically equal to the atomic/molecular mass in amu for a given substance.

Atomic mass concept (weighted average)

The average atomic mass is a weighted average of isotopic masses according to their natural abundances.

Isotopic abundance and average mass relationship

Different isotopes with different masses contribute to a weighted average that appears on the periodic table.

Example computation: 25.0 g Cl to iodine mass

To find a mass of I with the same number of atoms as 25.0 g Cl, use m_I = (mass Cl / M(Cl)) × M(I); apply Avogadro’s concept.

Isotope mass numbers (e.g., 185Re, 187Re)

Mass numbers of isotopes differ by neutrons; used with abundances to calculate average atomic mass.

Symbolic isotope notation

Notation like 12C, 13C, 12C+, 13C- indicates isotope (mass) and charge state (ion).