Atoms: The Building Blocks of
Matter
MODERN CHEMISTRY CH 3
3-1 The Atom: From Philosophical Idea to
Scientific Theory
 Democritus (400B.C.) – Greek philosopher who first
used the term atom to describe a basic indivisible
particle of matter.
 Chemical Reaction – transformation of substances into
one or more new substances
 Law of Conservation of Mass – mass is neither created or
destroyed during ordinary chemical reactions or physical
changes
 Law of Definite Proportions – chemical compounds
contain the same elements in exactly the same
proportions my mass (water is always 2-H and 1-O)
EX: Every sample of pure salt (NaCl) contains 39.34%
Na and 60.66% Cl by mass.
3-1 The Atom: From Philosophical Idea to
Scientific Theory
 Law of Multiple Proportions – If two or more different
compounds are composed of the same two elements, then
the ratio of the masses of the second element combined
with a certain mass of the first element is always a ratio of
small whole numbers.
- carbon monoxide (CO) contains 1.33 g O and 1.0 g C
- carbon dioxide (CO2) contains 2.66 g O and 1.0g C
- the ratio of masses of oxygen in the two compounds is
2.66 = 2
1.33
1
3-1 The Atom: From Philosophical Idea to
Scientific Theory
 John Dalton’s Atomic Theory p. 66
1) All matter is composed of extremely small atoms.
2) Atoms of a given element are identical in size mass, and other properties.
3) Atoms cannot be subdivided, created, or destroyed.
4) Atoms of different elements combine in simple whole number ratios to form
compounds.
5 ) In chemical reactions, atoms are combined, separated, or rearranged, but
they are NOT
created or destroyed.
 Modern Atomic Theory –
-
Atoms can be subdivided into smaller particles.
All atoms of the same element are NOT identical.
3-2 The Structure of the Atom
 Atom – smallest particle of an element that retains the chemical properties of
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



that element
Nucleus – small region near the center of the atom (positively charged)
Subatomic Particles – positively charged protons, neutral neutrons, negatively
charged electrons
Cathode Ray Tubes – glass tubes containing gases at low pressure exposed to
electrical currents
J.J. Thomson – (1897, English Physicist) measured the charge to mass ratio of
the cathode rays. The rays were composed of negatively charged particles later
named electrons. He developed the Plum Pudding Model of the atom.
Robert Millikan – (1909, American Physicist) he used the Oil Drop Experiment
to determine the exact charge and mass of the electron. Electron has a mass of
1/1837th of a single hydrogen atom (basically 1/1837th the mass of a proton) [p.
71]
3-2 The Structure of the Atom
 From these two experiments we now know
- atoms are electrically neutral – they have as many positive charges as they do
negative charges
- electron has so much less mass than atoms the atoms must contain other
particles to account for their mass.
 Ernest Rutherford – (1911, New Zealand) In the Gold Foil Experiment he used
a radioactive source to fire positively charged alpha particles toward a sheet of
gold foil. Protons must be contained in a dense nucleus. The atom is mostly
empty space. (p. 72)
 Niels Bohr – (1913, Denmark) his work showed that electrons follow specific
paths or orbits around the nucleus of the atom in an electron cloud. (p. 96-97)
2-3 The Structure of the Atom
Subatomic Particles
Particle Symbol Charge Mass # Relative Mass (a.m.u.)
Actual Mass (kg)
Electron e-, 0-1e
-1
0
0.0005486
9.1x10-31
Proton p+, 11H
+1
1
1.007276
1.64x10-27
Neutron n0, 10n
0
1
1.008665
1.68x10-27
3-2 The Structure of the Atom
 Nuclear Forces – short range proton-neutron, proton-proton, and neutron-
neutron forces holding the nuclear particles together (overcomes the repulsive
forces of like charges)
 Electron Cloud – region containing negative charge (electrons)
 Atomic Radii – range from 40-270 pm (1 pm = 1x10-12)
3-3 Counting Atoms
 Atomic Number (Z) – identifies the element and gives the number of protons in
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the nucleus of each atom of that element (whole number on the periodic table)
Neutral Atoms - # protons = #electrons
Isotopes – atoms of the same element have the same #p+, but different #n0
and, therefore different masses (p. 75-77)
EX: Hydrogen isotopes:
Protium (Hydrogen -1)
1 proton, 1 electron
Deuterium (Hydrogen -2) 1 proton, 1 electron, 1 neutron
Tritium (Hydrogen – 3) 1 proton, 1 electron, 2 neutrons
Mass Number – total number of protons and neutrons (#p+ + #n0 = mass#)
Nuclear Symbol – shows composition of isotope’s nucleus (superscript is the
mass number and the subscript is the atomic number)
EX: 23592U (uranium – 235) and 21H (hydrogen – 2)
3-3 Counting Atoms
 Nuclide – general term for any isotope of any element
 EX #1:
How many protons, electrons, and neutrons are there is an
atom of chlorine – 37?
atomic number = 17
#p+ = 17
#e- = 17
mass number = 37
#n0 = mass number – atomic number
= 37 -17 = 20
Pause for LOTS of practice with this! (Atomic Structure WS I and II)
3-3 Counting Atoms
 Relative Atomic Mass – the mass of an atom compared to the carbon-12 isotope
 Atomic Mass Unit – exactly 1/12 the mass of a carbon-12 atom
 Average Atomic Mass – weighted average of the atomic masses of the naturally
occurring isotopes of an element
EX: What is the average atomic mass of naturally occurring copper which
consists of 69.17% copper-63 (mass of 62.929598 amu) and 30.83 %
copper-65 (mass of 64.927793 amu)?
K:
69.17% Cu-63
30.83% Cu-65
unk:
average atomic mass
(0.6917)(62.929598amu) + (0.3083)(64.927793) = 63.55 amu
3-3 Counting Atoms
 Composition Stoichiometry - this is basically doing dimensional analysis
with atoms and mass
 Mole (mol) – SI unit for the amount of substance
1 mole contains as many particles as there are in exactly 12 g of carbon-12
Avogadro’s Number = 6.022x1023 the number of particles in exactly one
mole of a pure substance
Molar Mass – the mass of one mole of a pure substance (g/mol);
numerically equal to the atomic mass of the element in atomic mass
units
(periodic table)
EX:
1 mol Xe = 131.29 g Xe
1 mol Na = 22.98977 g Na
3-3 Counting Atoms
 1 mole = 6.022x1023 atoms (or particles)
 1 mole = weight in grams from the P.T.
 We will now work SEVERAL compositional
stoichiometry practice problems from your book p.
83-85.
3-3 Counting Atoms
 Compositional Stoichiometry Examples