Atomic Theory and Structure Review

Flashcards covering key vocabulary, models, experiments, and principles from the Atomic Theory lecture notes, Chapters 4 and 11.

Scientific Models

Represent things difficult to visualize, scaled-down for large objects (solar system) or scaled-up for small ones (atoms).

Democritus's Atomic Model (400 BC)

Proposed 'atomos' (uncuttable or indivisible) particles as the basic structure of matter.

Law of Conservation of Mass (Lavoisier)

Mass is neither created nor destroyed during ordinary chemical reactions; mass of products equals mass of reactants.

Law of Definite Proportions (Proust)

A chemical compound contains the same elements in the same percent by mass regardless of the sample size or source.

Law of Multiple Proportions (Dalton)

When elements combine, they do so in the ratio of small whole numbers (e.g., CO2 and CO).

Dalton's Atomic Theory (1808)

Principles stating that all matter is composed of extremely small, indivisible atoms; atoms of an element are identical; atoms cannot be created, destroyed, or transformed; compounds form from atoms combining in whole-number ratios.

Dalton's Billiard Ball Model

Likens the atom to a solid, indivisible, indestructible sphere, representing a single, complete unit of matter.

Cathode Ray Tube (CRT)

A sealed glass tube with two metal plates (cathode and anode) connected to an electricity source, used to study the flow of charge through gas.

Cathode Rays

Rays observed originating at the negative electrode (cathode) and moving to the positive electrode (anode) in a CRT, later identified as streams of negatively charged particles.

Properties of Cathode Rays

Travel in straight lines, consist of negatively charged particles (bend towards positive plate in an electric field), and have mass (turn a paddle wheel).

J.J. Thomson (1897)

Credited with discovering the electron and determining its charge-to-mass ratio, concluding that electrons are fundamental particles of all matter.

Millikan's Oil Drop Experiment (1909)

Experiment that determined the fundamental charge of an electron by observing the motion of tiny oil drops in an electric field, leading to the calculation of the electron's mass.

Inferences from Electron Properties

Atoms must contain positively charged particles to balance the negatively charged electrons and other particles to account for most of their mass.

Thomson's Plum Pudding Model

Proposed that the atom is a sphere of uniformly distributed positive charge with negatively charged electrons embedded within it, like 'raisins in plum pudding'.

Rutherford's Gold Foil Experiment (1911)

Experiment firing alpha particles at gold foil, leading to the discovery that atoms are mostly empty space with a tiny, dense, positively charged nucleus.

Nucleus

The tiny, dense, positively charged central region of an atom containing most of its mass and all of its positive charge.

Proton (Rutherford's Proposal)

Positively charged particles proposed by Rutherford to reside within the nucleus.

Moseley's Atomic Number (1913)

Discovered using X-rays that each element differs from the next by having one more positive charge in the nucleus, defining the atomic number as the number of protons.

Continuous Spectrum

A spectrum that contains all wavelengths of visible light, produced when white light passes through a prism.

Line Emission Spectrum

A spectrum containing only a few discrete lines (specific wavelengths), unique to each element, produced when excited gas emits light.

Bohr's Theory of the Hydrogen Atom

Proposed that electrons orbit the nucleus in specific circular orbits with quantized energy levels, emitting or absorbing photons when changing orbits.

Quantized Energy Levels

The concept that electrons can only exist in specific, discrete energy states or orbits within an atom, meaning energy is emitted or absorbed in fixed packets.

Electromagnetic Radiation

A form of energy that exhibits wave-like properties and travels at the speed of light, consisting of oscillating electric and magnetic fields.

Wavelength (λ)

The distance between identical points on successive waves.

Frequency (ν)

The number of waves that pass a point per second, measured in Hz or 1/s.

Speed of Light (c)

The constant speed at which all electromagnetic radiation travels in a vacuum, approximately 2.9979 x 10^8 m/s.

Double-Slit Experiment

An experiment demonstrating the wave nature of light (and later electrons) by producing interference patterns when passing through two closely spaced slits.

Electron Diffraction

Experiments demonstrating the wave properties of electrons by producing diffraction patterns.

Electron as a Standing Wave

The concept that only certain circular orbits are allowed for electrons because they must accommodate a whole number of wavelengths of the electron's standing wave.

Quantum Theory

The theory that energy is not continuous but composed of discrete units called quanta, challenging the classical view of matter and energy.

Blackbody Radiation

Light given off by a hot object, where the amount of energy emitted at a certain temperature depends on the wavelength; classical physics failed to explain this.

Quantum (Planck's Concept)

The smallest discrete quantity (bundle) of energy that can be absorbed or emitted as electromagnetic radiation (E = hν).

Planck's Constant (h)

A fundamental physical constant relating the energy of a photon to its frequency (6.626 x 10^-34 J·s).

Photoelectric Effect

The ejection of electrons from the surface of a metal when exposed to light of a certain minimum (threshold) frequency, explained by Einstein.

Photon

A particle of light, as proposed by Einstein, that carries energy given by E = hν.

Threshold Frequency

The minimum frequency of light required to eject electrons from a metal surface in the photoelectric effect.

Dual Nature of Light

The concept that light exhibits both wave properties (like interference) and particulate properties (like photons in the photoelectric effect).

Heisenberg Uncertainty Principle

States that it is impossible to know simultaneously both the exact momentum and the exact position of a particle with certainty.

Modern Atomic Theory

Describes an atom as an electrically neutral, spherical entity with a tiny, dense, positively charged central nucleus surrounded by a 'cloud' of rapidly moving, negatively charged electrons.

Modern Model of the Atom

States that precise electron paths cannot be determined, but the probability of finding electrons in specific locations (orbitals) can be specified by shell, subshell, orbital, and spin number.

Proton

A subatomic particle with a +1 charge, relative mass of 1 amu, and actual mass of 1.7 x 10^-24 g, found in the nucleus.

Neutron

A neutral subatomic particle (0 charge) with a relative mass of 1 amu and actual mass of 1.7 x 10^-24 g, found in the nucleus.

Electron

A subatomic particle with a -1 charge, relative mass of 0 amu, and actual mass of 9.1 x 10^-28 g, found in the electron cloud around the nucleus.

Atomic Mass Unit (amu or u)

A relative mass scale unit used for atoms and subatomic particles, defined as 1/12 the mass of a carbon-12 atom.

Atomic Number

The number of protons in the nucleus of an atom, which uniquely defines an element and determines its identity.

Neutral Atom

An atom in which the number of protons equals the number of electrons, resulting in a net electrical charge of zero.

Mass Number

The total number of protons plus neutrons in an atom's nucleus, always a whole number and not found on the periodic table.

Isotopes

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

Nuclear Symbol / Hyphen Notation

Ways to represent an isotope, e.g., ^12C or Carbon-12, indicating the mass number and atomic number (or element identity).