Chapter 2 Review41 Total Terms
Chapter 2 Key Terms & Vocabulary
Comprehensive index of all scientific terms, definitions, and examples across Chapter 2: Atoms, Molecules, and Ions.
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From Section 2.1: Early Ideas in Atomic Theory
Dalton’s atomic theoryAtomic Theory
A scientific model proposed by John Dalton stating that matter is composed of indestructible atoms that combine in whole-number ratios to form compounds.
Example: Explains why chemical reactions conserve mass.
From Section 2.1: Early Ideas in Atomic Theory
atomMatter & Structure
The smallest unit of an element that can participate in a chemical change.
Example: A single copper atom (Cu).
From Section 2.1: Early Ideas in Atomic Theory
elementMatter & Structure
A substance consisting of only one type of atom, with a characteristic mass and identical chemical properties.
Example: Pure gold (Au).
From Section 2.1: Early Ideas in Atomic Theory
law of definite proportionsChemical Laws
The principle that all samples of a pure compound contain the exact same elements in the exact same proportion by mass.
Example: Pure water is always 11.2% hydrogen and 88.8% oxygen by mass.
From Section 2.1: Early Ideas in Atomic Theory
law of constant compositionChemical Laws
Another name for the law of definite proportions, highlighting that a pure compound
Example: Carbon dioxide always contains carbon and oxygen in a fixed mass ratio.
From Section 2.1: Early Ideas in Atomic Theory
law of multiple proportionsChemical Laws
The principle that when two elements react to form more than one compound, a fixed mass of one element will react with masses of the other element in a ratio of small, whole numbers.
Example: Carbon and oxygen combining to form CO and CO₂.
From Section 2.2: Evolution of Atomic Theory
ThomsonHistorical Figures
British physicist who discovered the electron and proposed the plum pudding model of the atom.
Example: J.J. Thomson used cathode ray tubes in 1897.
From Section 2.2: Evolution of Atomic Theory
cathode rayExperimental Apparatus
A beam of electrons emitted from the cathode in a evacuated glass tube under high voltage.
Example: Deflected by electric and magnetic fields.
From Section 2.2: Evolution of Atomic Theory
electronSubatomic Particles
A negatively charged, subatomic particle with a mass more than one thousand times less than an atom.
Example: Found orbiting or surrounding the atomic nucleus.
From Section 2.2: Evolution of Atomic Theory
MillikanHistorical Figures
American physicist who measured the exact charge of a single electron using oil-drop experiments.
Example: Robert Millikan
From Section 2.2: Evolution of Atomic Theory
NagaokaHistorical Models
Japanese physicist who proposed a Saturn-like atomic model featuring a positive sphere surrounded by electron rings.
Example: An early alternative to Thomson
From Section 2.2: Evolution of Atomic Theory
RutherfordHistorical Figures
New Zealand physicist who discovered the atomic nucleus via alpha particle gold foil scattering.
Example: Ernest Rutherford established the nuclear model of the atom.
From Section 2.2: Evolution of Atomic Theory
alpha particles (α particles)Radiation & Particles
High-speed, positively charged particles consisting of two protons and two neutrons.
Example: Emitted by radioactive decay of radium.
From Section 2.2: Evolution of Atomic Theory
GeigerHistorical Figures
German physicist who assisted Rutherford in conducting the gold foil scattering experiments.
Example: Hans Geiger later co-invented the Geiger counter.
From Section 2.2: Evolution of Atomic Theory
MarsdenHistorical Figures
British physicist who worked with Rutherford and Geiger on the gold foil experiment as an undergraduate.
Example: Observed unexpected large-angle scattering of alpha particles.
From Section 2.2: Evolution of Atomic Theory
nucleusAtomic Structure
The small, relatively heavy, positively charged center of an atom containing protons and neutrons.
Example: Contains over 99.9% of the atom
From Section 2.2: Evolution of Atomic Theory
protonSubatomic Particles
A positively charged subatomic particle located within the atomic nucleus.
Example: Number of protons determines the element
From Section 2.2: Evolution of Atomic Theory
SoddyHistorical Figures
English chemist who discovered isotopes, demonstrating that elements can have atoms of different masses.
Example: Frederick Soddy won the 1921 Nobel Prize in Chemistry.
From Section 2.2: Evolution of Atomic Theory
isotopesAtomic Structure
Atoms of the same element that share the same number of protons but differ in mass due to different neutron counts.
Example: Carbon-12 and Carbon-14.
From Section 2.2: Evolution of Atomic Theory
ChadwickHistorical Figures
English physicist who discovered the neutron in 1932.
Example: James Chadwick
From Section 2.2: Evolution of Atomic Theory
neutronsSubatomic Particles
Uncharged, subatomic particles with a mass approximately equal to that of a proton, located in the nucleus.
Example: Neutrons stabilize nuclei and account for isotope mass differences.
From Section 2.3: Atomic Structure and Symbolism
unified atomic mass unit (u)Measurement
A unit of mass defined as exactly 1/12 of the mass of a single carbon-12 atom, equal to 1.6605 × 10⁻²⁴ g.
Example: 1 u ≈ mass of one proton or neutron
From Section 2.3: Atomic Structure and Symbolism
fundamental unit of charge (e)Subatomic Properties
The elementary unit equal to the magnitude of the charge of a single electron, equal to 1.602 × 10⁻¹⁹ C.
Example: Electron charge is 1-, proton charge is 1+
From Section 2.3: Atomic Structure and Symbolism
dalton (Da)Measurement
An alternative mass unit that is numerically equivalent to the unified atomic mass unit (u).
Example: 1 Da = 1 u
From Section 2.3: Atomic Structure and Symbolism
atomic number (Z)Atomic Structure
The number of protons in the nucleus of an atom, which determines the chemical identity of the element.
Example: Carbon always has Z = 6
From Section 2.3: Atomic Structure and Symbolism
mass number (A)Atomic Structure
The total number of protons and neutrons in the nucleus of an atom.
Example: Carbon-12 has a mass number of 12
From Section 2.3: Atomic Structure and Symbolism
ionCharged Species
An electrically charged atom or molecule formed by gaining or losing one or more electrons.
Example: Na⁺ or Cl⁻
From Section 2.3: Atomic Structure and Symbolism
anionCharged Species
A negatively charged ion formed when an atom gains one or more electrons.
Example: O²⁻ or Cl⁻
From Section 2.3: Atomic Structure and Symbolism
cationsCharged Species
Positively charged ions formed when an atom loses one or more electrons.
Example: Na⁺ or Mg²⁺
From Section 2.3: Atomic Structure and Symbolism
chemical symbolNotation
A one-, two-, or three-letter abbreviation used to represent an element or an atom of an element.
Example: C for carbon, Hg for mercury
From Section 2.3: Atomic Structure and Symbolism
atomic massMeasurement
The weighted average mass of all naturally occurring isotopes of an element, expressed in atomic mass units (amu).
Example: Boron has an atomic mass of 10.81 amu
From Section 2.4: Chemical Formulas
molecular formulaFormulas
A representation of a molecule using chemical symbols and subscripts to indicate the exact number of each type of atom present.
Example: C₆H₁₂O₆ (glucose)
From Section 2.4: Chemical Formulas
structural formulaFormulas
A drawing or formula that shows the types and numbers of atoms in a molecule, as well as how those atoms are bonded together.
Example: H-O-H for water
From Section 2.4: Chemical Formulas
empirical formulaFormulas
A formula showing the simplest whole-number ratio of the elements in a compound.
Example: CH₂O for glucose
From Section 2.4: Chemical Formulas
molecular massMass Metrics
The sum of the average atomic masses of all the atoms represented in a molecule
Example: Water (H₂O) molecular mass = 18.02 amu
From Section 2.4: Chemical Formulas
isomersMolecular Structure
Compounds that have identical molecular formulas but different structural arrangements or spatial orientations.
Example: Acetic acid and methyl formate
From Section 2.4: Chemical Formulas
structural isomersMolecular Structure
Isomers in which the molecules differ in how their atoms are connected to one another.
From Section 2.4: Chemical Formulas
spatial isomersMolecular Structure
Isomers in which the atoms are connected in the same order but have different relative orientations in 3D space.
From Section 2.4: Chemical Formulas
moleQuantities
The amount unit representing 6.022 × 10²³ discrete entities (atoms or molecules).
Example: 1 mol of C = 6.022 × 10²³ carbon atoms
From Section 2.4: Chemical Formulas
AvogadroConstants
The experimentally determined number of entities composing 1 mole of substance, equal to 6.022 × 10²³ mol⁻¹.
From Section 2.4: Chemical Formulas
molar massMass Metrics
The mass in grams of 1 mole of a substance, numerically equivalent to its atomic or formula weight in amu.
Example: Carbon molar mass = 12.01 g/mol