PPT 3 - Bennett

Matter

Anything that has mass and occupies space.

Atom

Smallest unit of an element that retains its chemical identity.

Element

Pure substance made of one type of atom.

Compound

Substance containing different elements chemically combined in fixed proportions.

Molecule

Two or more atoms chemically bonded together.

Ion

Atom or group of atoms with a net electrical charge.

Cation

Positively charged ion; formed by losing electrons.

Anion

Negatively charged ion; formed by gaining electrons.

Isotope

Atoms of the same element with different numbers of neutrons.

Atomic number (Z)

Number of protons in an atom's nucleus.

Mass number (A)

Total number of protons plus neutrons.

Atomic mass unit (amu)

Unit used to express atomic-scale masses.

Average atomic mass

Weighted average of naturally occurring isotope masses.

Molar mass

Mass of one mole of a substance, in g/mol.

Mole

Amount of substance containing 6.02 × 10^23 entities.

Avogadro's number

6.02 × 10^23 particles per mole.

Molar volume at STP (course convention)

22.4 liters per mole of gas.

STP (course convention)

Standard temperature and pressure; use 22.4 L/mol for gases.

Hydrate

Compound with water associated in a fixed formula ratio.

Anhydrous

Without water; a hydrate after its water is removed.

Heat of fusion

Energy required to melt a substance at its melting point.

Molar heat of fusion

Energy needed to melt one mole, usually kJ/mol.

Stoichiometry

Quantitative relationships among substances in chemical calculations.

Chemical symbol

One- or two-letter shorthand for an element.

Chemical formula

Shows types and numbers of atoms in a substance.

Chemical equation

Uses formulas to represent reactants and products.

Coefficient

Number placed before a formula indicating relative amounts.

Subscript

Small number in a formula showing atoms of an element.

Reactant

Starting substance in a chemical reaction.

Product

Substance formed by a chemical reaction.

Mole ratio

Ratio of coefficients in a balanced chemical equation.

Law of conservation of mass

Mass is not created or destroyed in ordinary chemical reactions.

Law of definite proportions

A given compound has the same mass ratio of elements.

Law of multiple proportions

Masses combine in small whole-number ratios for different compounds.

Continuous theory of matter

Matter can be subdivided without limit; ancient view.

Discontinuous theory of matter

Matter consists of discrete atoms separated by empty space.

Plum pudding model

Electrons embedded in diffuse positive material.

Nuclear model

Tiny dense positive nucleus with electrons outside.

Planetary/Bohr model

Electrons occupy specific allowed energy levels.

Quantum mechanical model

Electrons described by probability clouds/orbitals.

Orbital

Three-dimensional region describing likely electron location.

Quantized energy

Energy comes in discrete allowed amounts.

Ground state

Lowest available energy state.

Excited state

Higher energy state after energy absorption.

Photon

Packet of electromagnetic energy.

Wave-particle duality

Matter and light can show both wave and particle behavior.

Uncertainty principle

Cannot know exact position and momentum simultaneously.

Alpha particle

Positively charged helium nucleus.

Cathode ray

Stream of negatively charged electrons in a discharge tube.

Nucleus

Dense central part of an atom containing protons and neutrons.

Proton

Positively charged nuclear particle; approximately 1 amu.

Neutron

Uncharged nuclear particle; approximately 1 amu.

Electron

Negatively charged particle; very small mass.

Neutral atom

Has equal numbers of protons and electrons.

Nuclear charge

Positive charge determined by the number of protons.

Electron cloud

Probability distribution for electrons around the nucleus.

Atomic mass vs mass number

Atomic mass is weighted average; mass number counts nucleons in one isotope.

Protium

Hydrogen-1: 1 proton and 0 neutrons.

Deuterium

Hydrogen-2: 1 proton and 1 neutron.

Tritium

Hydrogen-3: 1 proton and 2 neutrons.

Historical figure: Empedokles

Proposed four elements: earth, air, fire, water.

Historical figure: Aristotle

Supported continuous matter and the four-element tradition.

Historical figure: Plato

Supported the ancient continuous view of matter.

Historical figure: Democritus

Proposed indivisible atoms and empty space between them.

Historical figure: Roger Bacon

Advocated experimental study and the scientific method.

Historical figure: Alchemists

Developed laboratory equipment and methods while seeking transmutation.

Historical figure: Antoine Lavoisier

Law of conservation of mass; careful reactant/product weighing.

Historical figure: Joseph Proust

Law of definite proportions; constant compound composition.

Historical figure: John Dalton

Modern atomic theory and law of multiple proportions.

Historical figure: William Crookes

Cathode-ray experiments; early evidence of electrons.

Historical figure: J. J. Thomson

Identified electrons and their charge-to-mass ratio.

Historical figure: Max Planck

Proposed that energy is quantized.

Historical figure: Albert Einstein

Explained light's particle-like behavior/photoelectric effect.

Historical figure: Robert Millikan

Oil-drop experiment measured electron charge and mass.

Historical figure: Ernest Rutherford (1911)

Gold-foil experiment revealed a small dense nucleus.

Historical figure: Henry Moseley

Linked nuclear charge to atomic number; reordered periodic table.

Historical figure: Niels Bohr

Proposed electrons in fixed energy levels.

Historical figure: Ernest Rutherford (1919)

Identified the proton through nuclear experiments.

Historical figure: Louis de Broglie

Proposed wave-like behavior of matter.

Historical figure: Erwin Schrödinger

Wave equations and three-dimensional electron orbitals.

Historical figure: Werner Heisenberg

Uncertainty principle.

Historical figure: James Chadwick

Discovered the neutron.

What did ancient four-element thinkers call the four properties?

Hot, cold, wet, and dry.

Why was Democritus's theory rejected for centuries?

It lacked experimental evidence; continuous matter was favored.

What did Dalton's theory say about reactions?

Atoms combine, separate, or rearrange.

Which Dalton claim was disproved by isotopes?

All atoms of one element have identical masses.

Which Dalton claim was disproved by subatomic particles?

Atoms cannot be subdivided.

What did Crookes use?

A cathode-ray (vacuum discharge) tube.

What did Thomson's cathode-ray results show?

Electrons are negative particles present in atoms.

What is Thomson's electron charge-to-mass ratio?

About 1.76 × 10^8 C/g (course slides).

What did Millikan measure for electron charge?

About −1.6 × 10^-19 coulomb.

What mass did Millikan determine for an electron?

About 9.1 × 10^-28 grams.

What was Rutherford's gold-foil observation?

Most alpha particles passed through; a few deflected sharply.

What did the gold-foil experiment imply?

Atoms are mostly empty space with a small dense positive nucleus.

Why did Bohr propose energy levels?

To explain discrete spectral lines emitted by atoms.

What happens when an electron absorbs energy?

It may move to a higher allowed energy level.

What happens when an electron falls to a lower level?

It emits energy as a photon.

Why was Bohr's model limited?

It explained hydrogen better than multi-electron atoms.

What is de Broglie's wavelength formula?

λ = h/(mv), as given in the slides.

What replaced Bohr's fixed electron paths?

Three-dimensional orbitals in the quantum model.