Chemistry: Atoms, Molecules, and Ions

Flashcards covering key concepts from Chapter 2 'Atoms, Molecules, and Ions' including atomic theory, subatomic particles, isotopes, the periodic table, chemical formulas, ions, and chemical nomenclature.

Atomic Theory of Matter

An organized theory developed by John Dalton in the early 1800s, based on laws of constant composition, conservation of mass, and multiple proportions, suggesting that matter is made of indivisible particles called atoms.

Atomos

A term used by some Greek philosophers, like Democritus, meaning 'uncuttable,' referring to the smallest particle that made up all of nature.

Law of Constant Composition

Discovered by Joseph Proust, it states that compounds have a definite composition, meaning the relative number of atoms of each element in the compound is the same in any sample.

Law of Conservation of Mass

Discovered by Antoine Lavoisier, it states that the total mass of substances present at the end of a chemical process is the same as the mass of substances present before the process took place.

Law of Multiple Proportions

Discovered by John Dalton, it states that if two elements, A and B, form more than one compound, the masses of B that combine with a given mass of A are in the ratio of small whole numbers.

Dalton's Atomic Theory Postulate 1

Each element is composed of extremely small particles called atoms.

Dalton's Atomic Theory Postulate 2

All atoms of a given element are identical to one another in mass and other properties, but the atoms of one element are different from the atoms of all other elements.

Dalton's Atomic Theory Postulate 3

Atoms of an element are not changed into atoms of a different element by chemical reactions; atoms are neither created nor destroyed in chemical reactions.

Dalton's Atomic Theory Postulate 4

Atoms of more than one element combine to form compounds; a given compound always has the same relative number and kind of atoms.

Electron

A negatively charged subatomic particle, discovered by J. J. Thomson, streams of which emanate from cathode tubes.

J. J. Thomson

Credited with the discovery of the electron (cathode rays) in 1897 and measured its charge/mass ratio.

Millikan Oil-Drop Experiment

An experiment conducted by Robert Millikan in 1909 that determined the charge on the electron.

Radioactivity

The spontaneous emission of high-energy radiation by an atom, first observed by Henri Becquerel.

Ernest Rutherford (radioactivity)

Discovered three types of radiation: alpha particles (positively charged), beta particles (negatively charged, like electrons), and gamma rays (uncharged).

Plum Pudding Model

A model of the atom proposed by J. J. Thomson around 1900, featuring a positive sphere of matter with negative electrons embedded in it.

Rutherford's Gold Foil Experiment

An experiment where Ernest Rutherford shot alpha particles at a thin sheet of gold foil, revealing that the plum pudding model was incorrect and leading to the discovery of the atomic nucleus.

Nuclear Atom

Rutherford's model of the atom, which postulates a very small, dense positive center (nucleus) with electrons moving around the outside, indicating that most of the atom is empty space.

Proton

A subatomic particle found in the nucleus with a positive charge (+1) and a relative mass of approximately 1 amu.

Neutron

A neutral subatomic particle found in the nucleus with no charge and a relative mass of approximately 1 amu.

Atomic Number

The number of protons in the nucleus of an atom, which determines the identity of the element. In a neutral atom, it also equals the number of electrons.

Mass Number

The total number of protons and neutrons in the nucleus of an atom, written as a superscript before the element symbol.

Isotopes

Atoms of the same element that have the same number of protons but different numbers of neutrons, resulting in different masses.

Atomic Mass Unit (amu)

A base unit used for measuring atomic-level masses, defined as 1/12th the mass of a carbon-12 atom.

Atomic Weight

The average mass of all isotopes of an element, weighted by their relative abundances, used for calculations in realistic scenarios involving large amounts of atoms.

Periodic Table

A systematic organization of the elements, arranged in order of increasing atomic number, where elements with similar chemical properties are grouped together.

Periods

The horizontal rows on the periodic table.

Groups

The vertical columns on the periodic table, where elements in the same group have similar chemical properties.

Periodicity

The repeating pattern of chemical properties and reactivity observed when looking at elements arranged in the periodic table.

Alkali metals

Elements in Group 1A of the periodic table (Li, Na, K, Rb, Cs, Fr).

Alkaline earth metals

Elements in Group 2A of the periodic table (Be, Mg, Ca, Sr, Ba, Ra).

Chalcogens

Elements in Group 6A of the periodic table (O, S, Se, Te, Po).

Halogens

Elements in Group 7A of the periodic table (F, Cl, Br, I, At).

Noble gases

Elements in Group 8A of the periodic table (He, Ne, Ar, Kr, Xe, Rn).

Metals

Elements typically found on the left side of the periodic table, characterized by shiny luster, conductivity of heat and electricity, and generally existing as solids (except mercury).

Nonmetals

Elements typically found on the right side of the periodic table (including H), which can be solid, liquid, or gas at room temperature and generally lack metallic properties.

Metalloids

Elements located on the steplike line of the periodic table (except Al, Po, At), exhibiting properties that are sometimes like metals and sometimes like nonmetals.

Chemical Formula

A representation that uses subscripts to the right of an element's symbol to indicate the number of atoms of that element in one molecule or formula unit of a compound.

Molecular Compounds

Compounds composed of molecules that almost always contain only nonmetals.

Diatomic Molecules

Elements that occur naturally as molecules containing two atoms (H2, N2, O2, F2, Cl2, Br2, I2).

Empirical Formulas

Chemical formulas that give the lowest whole-number ratio of atoms of each element in a compound.

Molecular Formulas

Chemical formulas that give the exact number of atoms of each element in a compound.

Structural Formulas

Formulas that show the order in which atoms are attached in a molecule, without necessarily depicting their three-dimensional shape.

Ions

Atoms or groups of atoms that have lost or gained electrons, resulting in an overall electrical charge.

Cations

Ions formed when an atom or group of atoms loses at least one electron, resulting in a positive charge. Monatomic cations are typically formed by metals.

Anions

Ions formed when an atom or group of atoms gains at least one electron, resulting in a negative charge. Monatomic anions are typically formed by nonmetals (except noble gases).

Polyatomic Ions

Groups of atoms that collectively gain or lose electrons, resulting in an overall electrical charge.

Ionic Compounds

Compounds generally formed between metals and nonmetals where electrons are transferred from the metal to the nonmetal, creating oppositely charged ions that attract each other.

Chemical Nomenclature

The systematic method or system used for naming chemical compounds.

Oxyanion Nomenclature (fewer oxygens)

When comparing two oxyanions of the same element, the one with fewer oxygens ends in -ite (e.g., nitrite NO2-).

Oxyanion Nomenclature (more oxygens)

When comparing two oxyanions of the same element, the one with more oxygens ends in -ate (e.g., nitrate NO3-).

Oxyanion Nomenclature (fewest oxygens)

For a series of oxyanions with varying oxygen atoms, the one with the fewest oxygens has the prefix hypo- and ends in -ite (e.g., hypochlorite ClO-).

Oxyanion Nomenclature (most oxygens)

For a series of oxyanions with varying oxygen atoms, the one with the most oxygens has the prefix per- and ends in -ate (e.g., perchlorate ClO4-).

Acid Nomenclature (-ide anion)

If the anion in the acid ends in -ide, change the ending to -ic acid and add the prefix hydro- (e.g., HCl is hydrochloric acid).

Acid Nomenclature (-ite anion)

If the anion in the acid ends in -ite, change the ending to -ous acid (e.g., HClO2 is chlorous acid).

Acid Nomenclature (-ate anion)

If the anion in the acid ends in -ate, change the ending to -ic acid (e.g., HClO3 is chloric acid).

Nomenclature of Binary Molecular Compounds

A system for naming compounds formed between two nonmetals, using prefixes to denote the number of atoms of each element and changing the ending of the second element to -ide (e.g., CO2 is carbon dioxide).

Organic Chemistry

The branch of chemistry dedicated to the study of carbon compounds.

Hydrocarbons

Organic compounds containing only carbon and hydrogen.

Alkanes

The simplest type of hydrocarbons, named by a prefix indicating the number of carbons followed by the suffix -ane (e.g., methane for 1 carbon).

Functional Group

A specific group of atoms within a molecule that is responsible for the characteristic chemical reactions of that molecule.

Alcohols

Organic compounds formed when a hydrogen in an alkane is replaced with a hydroxyl (-OH) functional group; their names typically end in -ol.

Isomers

Molecules that have the same chemical formula but different structural arrangements of atoms.