General Chemistry: Electronic Structure, Atomic Theory, and Measurement

A comprehensive vocabulary review of 120 key concepts, laws, formulas, and historical figures from Chemistry Chapter 1 (Matter and Energy), Chapter 2 (Atoms, Ions, and Elements), and Chapter 3 (Electronic Structure of the Atom).

Dimensional Analysis

A problem-solving strategy that uses conversion factors and units as a guide to convert one measured quantity to another.

SI Units

An internationally agreed-upon set of standard metric units used in scientific measurements.

Temperature

A measure of the average kinetic energy of the particles that compose a sample of matter.

Celsius Scale

A metric temperature scale based on the freezing point ($0\,^{\circ}\text{C}$) and boiling point ($100\,^{\circ}\text{C}$) of pure water.

Kelvin Scale

The SI base unit scale for temperature, based on gas properties and defined as an absolute scale where $K = {}^{\circ}\text{C} + 273.15$.

Giga (G)

A metric prefix representing $10^9$ times the base unit.

Mega (M)

A metric prefix representing $10^6$ times the base unit.

Kilo (k)

A metric prefix representing $10^3$ times the base unit.

Deci (d)

A metric prefix representing $10^{-1}$ of the base unit.

Centi (c)

A metric prefix representing $10^{-2}$ of the base unit.

Milli (m)

A metric prefix representing $10^{-3}$ of the base unit.

Micro ($\mu$)

A metric prefix representing $10^{-6}$ of the base unit.

Nano (n)

A metric prefix representing $10^{-9}$ of the base unit.

Scientific Notation

A method of writing numbers as a product of a decimal number between 1 and 10 and 10 raised to an integer power.

Density

An intensive physical property unique to a substance, calculated as mass divided by volume ($\text{density} = \frac{m}{V}$).

Intensive Property

A property of matter that is independent of the amount of substance present, such as density, temperature, or color.

Extensive Property

A property of matter that depends directly on the amount of substance present, such as mass, volume, or length.

Exact Numbers

Values that have no uncertainty, originating from precise counting of objects, defined unit conversions, or integral numbers in formulas.

Inexact Numbers

Measured quantities obtained from experimental equipment that contain an inherent degree of uncertainty.

Significant Figures

All certain digits plus one estimated uncertain digit recorded in a measurement, reflecting the precision of the measured value.

Sandwiched Zeros

Zeros situated between two nonzero digits, which are always significant.

Leading Zeros

Zeros located to the left of the first nonzero digit in a number, which function only as place holders and are never significant.

Trailing Zeros

Zeros at the end of a number, which are significant if a decimal point is explicitly present.

Sig Fig Addition and Subtraction Rule

The final calculated answer cannot have more digits to the right of the decimal point than any of the original numbers in the calculation.

Sig Fig Multiplication and Division Rule

The number of significant figures in the result is set by the original number that contains the fewest significant figures.

John Dalton

The English chemist who proposed the first formal atomic theory of matter in 1808 based on existing scientific work.

Dalton's Atomic Theory

A theory proposing that matter consists of indivisible atoms, all atoms of a given element are identical, and atoms combine in small whole-number ratios to form compounds.

Law of Conservation of Matter

A fundamental law stating that matter cannot be created or destroyed in a chemical reaction.

Antoine Lavoisier

French chemist recognized as the father of modern chemistry who established the Law of Conservation of Matter and discovered the role of oxygen in combustion.

Law of Definite Proportions

A law formulated by Joseph Proust stating that all samples of a given chemical compound always contain the same elements in the exact same proportion by mass.

Joseph Proust

The French chemist who established the Law of Definite Proportions (also known as the Law of Constant Composition) in 1800.

Law of Multiple Proportions

A law formulated by John Dalton in 1804 stating that when two elements form two different compounds, the masses of element B that combine with a fixed mass of element A are in a ratio of small whole numbers.

Cathode Ray Tube

A sealed glass tube evacuated of air across which a high voltage is applied, generating a beam of negatively charged cathode rays.

J. J. Thomson

The British physicist who discovered the electron in 1897 by measuring its charge-to-mass ratio using cathode ray tubes.

Electron Charge-to-Mass Ratio

The quantitative ratio determined by J. J. Thomson for an electron, equal to $-1.76 \times 10^8\,\text{C/g}$.

Robert Millikan

The American physicist who measured the exact charge of a single electron in 1909.

Millikan Oil Drop Experiment

An experiment measuring the electric field needed to balance charged oil droplets, revealing the electron's charge to be $-1.60 \times 10^{-19}\,\text{C}$ and its mass to be $9.10 \times 10^{-28}\,\text{g}$.

Plum Pudding Model

An early atomic model proposed by J. J. Thomson featuring negatively charged electrons embedded within a uniform, positively charged sphere.

Alpha ($\alpha$) Particles

Positively charged particles emitted during radioactive decay, identified by Ernest Rutherford.

Beta ($\beta$) Particles

Negatively charged particles emitted during radioactive decay, corresponding to high-speed electrons.

Gamma ($\gamma$) Rays

High-energy, neutral electromagnetic radiation emitted by radioactive substances, unaffected by electric or magnetic fields.

Ernest Rutherford

New Zealand physicist who discovered alpha and beta radiation and proposed the nuclear model of the atom based on his gold foil experiment.

Rutherford's Gold Foil Experiment

A 1911 experiment in which alpha particles were shot at gold foil; the deflection of a few particles at large angles disproved the plum pudding model.

Nucleus

The small, extremely dense central region of an atom containing all of its positive charge and virtually all of its mass.

James Chadwick

The English physicist who discovered the neutron in 1932.

Neutron

A neutral subatomic particle in the nucleus of an atom with a mass of approximately $1\,\text{amu}$ ($1.675 \times 10^{-27}\,\text{kg}$) and no electrical charge.

Proton

A positively charged subatomic particle in the nucleus of an atom with a relative charge of $+1$ and a mass of approximately $1\,\text{amu}$ ($1.673 \times 10^{-27}\,\text{kg}$).

Atomic Mass Unit (amu)

A mass scale unit defined as exactly $\frac{1}{12}$ the mass of a single carbon-12 atom.

Atomic Number ($Z$)

The number of protons in an atom's nucleus, which uniquely defines the identity of an element.

Mass Number ($A$)

The total number of protons and neutrons in the nucleus of an atom ($A = \text{protons} + \text{neutrons}$).

Cation

A positively charged ion formed when a neutral atom loses one or more electrons.

Anion

A negatively charged ion formed when a neutral atom gains one or more electrons.

Isotopes

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

Natural Abundance

The percentage of an element in nature that consists of a specific isotope.

Average Atomic Mass

The weighted average mass of an element's naturally occurring isotopes, calculated based on their relative abundances.

Mass Spectrometry

An analytical technique that measures isotopic masses and relative abundances by separating charged particles according to their mass-to-charge ratio.

Dmitri Mendeleev

The Russian chemist who published the first periodic table in 1869, arranging elements by atomic weight and predicting the properties of undiscovered elements.

Periodicity

The regular recurrence of similar physical and chemical properties among elements when ordered systematically.

Henry Moseley

The English physicist who defined atomic numbers in 1913 using X-ray frequencies, establishing the modern arrangement of the periodic table.

Main-Group Elements

Elements in the A columns of the periodic table (groups 1, 2, and 13–18) whose properties are highly predictable based on their group position.

Transition Elements

Elements in the d-block and f-block (designated B groups) of the periodic table whose properties are less predictable.

Metals

Elements located on the left and center of the periodic table that exhibit shiny luster, malleability, ductility, high conductivity, and tend to form cations.

Nonmetals

Elements on the upper right of the periodic table that are poor conductors, brittle when solid, and tend to gain electrons to form anions.

Metalloids

Elements located along the diagonal stairstep of the periodic table that display properties intermediate between metals and nonmetals.

Alkali Metals

Group 1A metals (excluding hydrogen) that are soft, highly reactive solids that react violently with water to form basic solutions.

Alkaline Earth Metals

Group 2A metals that are harder, denser, and less reactive with water than alkali metals.

Halogens

Group 7A reactive nonmetals that exist as diatomic molecules and are commonly called "salt formers."

Noble Gases

Group 8A nonmetals that exist as extremely unreactive, monatomic gases due to their filled valence shells.

Diatomic Elements

The seven elements that exist naturally as two-atom molecules: $H_2$, $N_2$, $O_2$, $F_2$, $Cl_2$, $Br_2$, and $I_2$.

Quantum-Mechanical Model

A modern atomic model describing electrons in terms of wave-particle duality and statistical probability distributions rather than fixed classical paths.

Wavelength ($\lambda$)

The distance between consecutive corresponding points of a wave, which determines the color of visible light.

Amplitude

The height of a wave from its centerline to its crest, which determines the intensity or brightness of light.

Frequency ($\nu$)

The number of wave cycles that pass a given point per unit of time, measured in hertz ($\text{Hz} = \text{s}^{-1}$).

Speed of Light ($c$)

The constant speed at which all electromagnetic radiation travels in a vacuum, equal to $3.00 \times 10^8\,\text{m/s}$.

Wave Diffraction

The bending of traveling waves as they encounter an obstacle or pass through an opening in a barrier.

Electromagnetic Spectrum

The complete range of electromagnetic radiation arranged by wavelength and frequency, ranging from radio waves to gamma rays.

Johannes Rydberg

The Swedish physicist who analyzed the spectrum of hydrogen and derived an empirical formula describing its emission wavelengths.

Rydberg Constant ($R_H$)

The empirical constant in the Rydberg equation, equal to $1.097 \times 10^7\,\text{m}^{-1}$ or $3.29 \times 10^{15}\,\text{Hz}$.

Blackbody

An idealized physical body that can absorb and emit electromagnetic radiation at all frequencies.

Incandescence

The emission of electromagnetic radiation from a object caused by high thermal energy.

Ultraviolet Catastrophe

The flaw in classical Rayleigh-Jeans theory that incorrectly predicted blackbodies would emit infinite energy at high frequencies.

Max Planck

The German physicist who resolved the ultraviolet catastrophe in 1900 by proposing that energy is quantized in discrete packets called quanta.

Quanta

Small, discrete packets of energy emitted or absorbed by matter.

Planck's Constant ($h$)

The fundamental constant of proportionality relating the energy of a quantum to its frequency, equal to $6.626 \times 10^{-34}\,\text{J}\cdot\text{s}$.

Photoelectric Effect

The ejection of electrons from a metal surface when it absorbs light above a specific threshold frequency.

Heinrich Hertz

The German physicist who discovered the photoelectric effect in 1887.

Threshold Frequency ($\nu_0$)

The minimum frequency of incident radiation required to eject electrons from a specific metal surface.

Albert Einstein

The physicist who explained the photoelectric effect in 1905 by proposing that light consists of particles called photons.

Photon

A quantum or particle of light whose energy is proportional to its frequency ($E_{\text{photon}} = h\nu = \frac{hc}{\lambda}$).

Niels Bohr

The Danish physicist who introduced a quantized model of the hydrogen atom in 1914.

Bohr Model

An atomic model proposing that electrons move in fixed, quantized circular orbits around the nucleus.

Bohr Radius

The radius of the lowest allowed electron orbit ($n = 1$) in the Bohr hydrogen atom.

Ground State

The lowest possible electronic energy state of an atom.

Excited State

Any electronic energy state of an atom higher in energy than its ground state.

Balmer Series

The emission series of hydrogen corresponding to electron transitions ending at $n_f = 2$, producing light in the visible spectrum.

Louis de Broglie

The French physicist who theorized wave-particle duality in 1924, proposing that moving matter exhibits wave properties.

De Broglie Wavelength

The wavelength associated with a particle of mass $m$ moving at velocity $v$, calculated as $\lambda = \frac{h}{mv}$.

Werner Heisenberg

The German theoretical physicist who formulated the Uncertainty Principle in 1927.

Uncertainty Principle

The principle stating that it is impossible to simultaneously determine both the exact position ($\Delta x$) and momentum ($\Delta mv$) of a subatomic particle.

Determinacy

A classical physics concept where the trajectory of a particle into the future can be definitely predicted if position and velocity are known.