Atomic Theory, Periodic Table, and Chemical Nomenclature

Vocabulary flashcards reviewing atomic theory history, atomic models, periodic table classifications, and chemical nomenclature rules.

Atomists

Greek philosophical school represented by Democritus and Leucippus proposing that matter is composed of small, indivisible parts called atoms (atomos) which cannot be created or destroyed, but are always changing.

Empedocles

Elementalist philosopher who believed in the "Four Roots" of matter: water, earth, fire, and air.

Aristotle

Elementalist who believed all matter is continuous and derived from the four elements, positing that change occurs through matter shifting from its potential state to an actual state.

Law of Conservation of Mass

Fundamental law established by Antoine Lavoisier stating that matter cannot be created or destroyed, only changed, shifting chemistry from qualitative alchemy to a quantitative physical science.

Hennig Brand

Scientist who discovered phosphorus and laid the groundwork for atomic theory.

Michael Faraday

Scientist who laid the foundations for the modern electrical age and discovered electromagnetic induction.

Law of Definite Proportions

Law proposed by Joseph Proust stating that chemical compounds always contain elements combined in fixed whole number proportions.

Robert Boyle

Referred to as the "Father of Chemistry" for replacing philosophical speculation with empirical experimentation.

Billiard Ball Model

John Dalton's atomic model depicting atoms as tiny, indivisible, homogeneous, and indestructible solid spheres.

Dalton's Atomic Theory

Theory stating that all matter is made of tiny, indivisible, homogeneous, and indestructible atoms; atoms of a given element are identical; atoms of different elements are different; and compounds form in fixed whole number ratios.

Law of Triads

Concept proposed by Johan Wolfgang Döbereiner where three elements were grouped together based on similar properties, laying groundwork for the Periodic Table.

Dmitri Mendeleev

Chemist who created the first Periodic Table arranged by atomic weights and valency; he left gaps to predict the existence and properties of undiscovered elements.

Antoine-Henry Becquerel

Scientist who discovered spontaneous radioactivity, established a connection between phosphorescence and x-rays, and helped disprove Dalton's idea of indivisible, immutable atoms.

Pierre Curie

Scientist who discovered piezoelectricity, formulated Curie's Law and Curie's Point regarding temperature-altered magnetism, and proved atoms naturally decay and split.

Marie Curie

Scientist who discovered the radioactive elements polonium and radium and theorized that certain materials are radioactive, proving atoms can naturally decay and split.

Plum Pudding Model

Atomic model proposed by J.J. Thomson (likened to a blueberry muffin) asserting that negatively charged bits (electrons) exist distributed within a positively charged space.

Cathode Ray Tube

Apparatus used by J.J. Thomson in experiments similar to Röntgen's to isolate and discover the electron.

Gold Foil Experiment

Experiment conducted by Ernest Rutherford in which ions were fired at a thin sheet of gold foil, revealing the atomic nucleus.

Rutherford's Atomic Model

Nuclear model of the atom featuring a dense positive nucleus surrounded by mostly empty space and negative electrons; its theoretical issue was that electrons should fall into the positive nucleus.

Bohr Model

Planetary model of the atom proposing a dense positive nucleus with electrons moving on fixed energy orbitals (rings), becoming "excited" when absorbing energy and "relaxing" when losing it.

Electron Cloud Model

The current Quantum Mechanical model of the atom, consisting of a central nucleus surrounded by a cloud of probability representing spaces where electrons are likely to be found in energy levels.

Groups (Families)

The $18$ vertical columns on the Periodic Table containing elements with similar chemical properties.

Periods

The $7$ horizontal rows across the Periodic Table.

Alkali Metals

Group 1 elements (excluding hydrogen) characterized by being very reactive with low density. 1 valence electron

Alkaline Earth Metals

Group 2 elements characterized by being somewhat reactive. 2 valence electrons

Transition Metals

Elements occupying Groups 3 through 12 on the Periodic Table whose charges cannot typically be predicted by group number alone.

Halogens

Group 17 elements characterized as very reactive, toxic nonmetals. 7 valence electrons

Noble Gases

Group 18 elements characterized as being very unreactive or inert with a charge of 0. 8 valence electrons (except for helium, which has only 2 total electrons)

Metals

Elements that are shiny/lustrous, reactive, malleable (can break into sheets), ductile (can be pulled into wire), and good conductors of heat and electricity.

Nonmetals

Elements that are dull in appearance, mostly unreactive, brittle, and act as electrical and thermal insulators.

Metalloids

Elements that possess chemical and physical properties intermediate between metals and nonmetals.

Isotopes

Atoms of the same element that have the same number of protons but different numbers of neutrons.

Atomic Number

The total number of protons found in the nucleus of an atom.

Mass Number

The total sum of protons plus neutrons in an atom's nucleus.

Molecule

A chemical particle formed from combinations of non-metal atoms.

Ionic Compound

A neutral chemical compound composed of combinations of metal cations and nonmetal anions.

Ions

Positively or negatively charged particles formed when atoms lose or gain electrons.

Cation

A positively charged ion formed when an atom loses electrons.

Anion

A negatively charged ion formed when an atom gains electrons.

Polyatomic Ions

Charged entities consisting of more than one atom combined that act collectively as a single ion.

Stock System

Nomenclature system for transition metal compounds using Roman numerals in parentheses to denote the charge of the metal cation (e.g., Iron (II) Chloride).

Common System (Nomenclature)

Naming system using suffixes on the metal root: "-ous" indicates the lower charge (e.g., $Fe^{2+}$ is Ferrous) and "-ic" indicates the higher charge (e.g., $Fe^{3+}$ is Ferric).

Hydrate

An ionic compound paired with water molecules in its formula, named by writing the salt followed by a Greek prefix indicating the number of $H_2O$ units (e.g., dihydrate for $2H_2O$).

Acid (Nomenclature)

Ionic compounds containing an $H^+$ cation; named as "hydro-[root]-ic acid" if the anion has no oxygen, "[root]-ous acid" if the anion ends in "-ite", and "[root]-ic acid" if the anion ends in "-ate".

Octet Rule

Chemical rule stating that atoms tend to gain, lose, or share electrons until they have full valence electron shells.

Dot Method

Technique for finding ionic charges by drawing an element's valence electrons to see how many it must lose (positive charge) or gain (negative charge).

Charge Balancing Formula

$(\text{Charge of Cation}) \times (\text{Subscript of Cation}) + (\text{Charge of Anion}) \times (\text{Subscript of Anion}) = 0$; used to calculate an unknown ion charge in neutral compounds.

Cross, Drop, Reduce

Method for writing neutral ionic formulas by crossing numerical charge values across to the opposite subscript, dropping positive/negative signs, and simplifying subscripts to the smallest ratio.

Ammonium

NH4^+

Nitrate

NO3^-

Sulfate

SO4^2-

Hydroxide

OH^-

Carbonate

CO3^2-

Acetate

C2H3O2^-

Phosphate

PO4^3-

HCl

Hydrochloric acid

HBr

Hydrobromic acid

HNO3

Nitric acid

H2SO4

Sulfuric acid

H3PO4

Phosphoric acid

HC2H3O2

Acetic acid