History of the Atom
Greek Philosophers
1. Democritus (460-370): Particle could be divided into smaller particles until a basic particle of
matter that could not be divided further was reached
a. Atoms: Ultimate, indestructible particles
2. Aristotle (384-322 BC): Believed all matter was continuous
Atom: The smallest particle of an element that retains the properties of that element that can exist alone or in
combination with other atoms
Lavoisier (1743-1794)
1. Law of the conservation of matter: Matter, like energy, is neither created nor destroyed in any
process
John Dalton (1803)
1. Atomic Theory
a. All matter is composed of extremely small particles called atoms
b. Atoms of a given element are identical in size, mass, and other properties, atoms of
different elements differ in size, mass and other properties
c. Atoms cannot be subdivided, created, or destroyed
d. Atoms of different elements can combine in simple whole # ratios to form chemical
compounds
e. In chemical reactions, atoms are combined, separated or rearranged
1. Atomic theory has been expanded to explain new observations
a. Successfully explained laws of conservation and definite composition (a
and d)
b. Exceptions:
1. Atoms are divisible
2. Elements can have atoms of
different masses
a. Isotopes
2. Conservation of mass: Mass is conserved in chemical reactions
a. A + B  AB
b. Mass ratio:
Proust (1799):
1. Law of Definite Proportions: If 2 or more different compounds are composed of the same 2
elements the masses of the second element combined with a certain mass of the first element can be
expressed as ratios of small whole #s
a. Compounds are always in certain proportion by mass
1. Pyrite (Fool’s gold)
Joseph John (JJ) Thomson
1. Discovery of the negative electron
a. Electron: Negatively charged subatomic particle
2.Cathode Ray tube experiments
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a. Cathode: Electrode connected to neg terminal
b. Anode: Electrode connected to pos terminal
3. Found the charge to mass ratio
a. (1.76 x 108 coulombs/g)
b. Concluded if charge could be determined, mass could be found
Robert A. Millikan 1868-1953
1. Discovery of charge (1.6 x 10-19 C) and thus the mass (9.11 x 10-28g) of an electron
2. Oil drop experiment
Rutherford: Discovery of positive nucleus
1. Bombarded thin metal foils with positively charged fast moving particles
2. Some of the particles deflected
a. Caused by a small densely packed bundle of matter with a positive charge
1. Discovered the proton
b. Concluded that electrons surrounded the positive nucleus
3. 1908: Nobel Prize for work in radioactivity
James Chadwick
1. Discovered the neutron
Isotopes: Atoms of the same element that have different masses (different #s of neutrons)
1. Number of protons is what determines the element
2. 3 Isotopes of Hydrogen
a. Protium: 1 proton
b. Deuterium: 1 proton and 1 neutron
c. Tritium: 1 proton and 2 neutrons
Nuclear (Symbolic) Notation
Atomic number: Represents number of protons in an element
Mass number: Represents total number of protons and neutrons in the nucleus of an isotope
Hyphen Notation
Lithium-7
Li-7
Relative Atomic Mass
1. Carbon 12 is the standard
a. Chosen by the International Organization of Scientists that govern units of measurement
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b. 12 atomic mass units
Atomic mass unit: U, amu
1. One atomic mass unit is exactly 1/12 the mass of the Carbon-12 atom
MASS NUMBER AND ATOMIC MASS ARE NOT EQUAL
Average Atomic Mass: Weighted average of the atomic masses of naturally occurring isotopes of an element
Ex: Copper-63 (atomic mass= 62.939598) 69.17%
Copper-65 (atomic mass = 64.927793) 30.83%
Atomic Mass: The relative atomic mass of atoms as written on the periodic table
The Mole
Avagodro’s Number: 6.022 137 x 1023 is the number of particles in exactly one mole of pure substance
Mole: Counting unit used in chemistry to count atoms, particles and molecules
1 mole = 6.02 x 1023 atoms = atomic mass of element
Molar Mass: The mass in grams of one mole of an element
-Numerically equal to the atomic mass of an element
-Atomic mass: amu’s or u
-Molar mass: grams
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