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Chapter 3
Atoms: The Building Blocks of
Matter
Democritus – (460-370 BC)
Atom – “indivisible” part of matter
Matter can only be divided into
smaller parts a limited number of
times until you reach the essence
of what that matter is.
Antoine Lavoisier (1743-1794) –
“Father of Modern Chemistry”
• Law of the Conservation of Mass – Matter
cannot be created or destroyed in a
chemical reaction. (1789)
Joseph Proust (1754-1826)
• The Law of Definite Proportions – a
chemical compound contains the same
elements in exactly the same proportions
regardless of the sample size.
• NaCl is always 39.34%Na and 60.66%Cl
• H2O is always 11.11%H and 88.89%O
John Dalton (1776-1844)
• The Law of Multiple Proportions – If two or
more 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.
Dalton proposed but couldn’t prove:
• Reacting gases and products
combine in ratios of small
whole numbers. J.L. GayLussac.
• Equal volumes of gases, under
the same conditions of
temperature and pressure, have
the same number of molecules.
Amedeo Avogadro.
Dalton’s Atomic Theory
(1808)
1. All matter is composed of small particles called
atoms.
2. Atoms of the same element are exactly alike,
atoms of different elements are different.
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 simply
combined, separated or rearranged.
Atomic Structure
• Atom – smallest particle
of an element that still
has the properties of that
element.
• Nucleus – very small,
very dense central portion
of the atom where
protons and neutrons are
located.
• Electron Cloud – area
surrounding the nucleus
where electrons are
located.
“Discovery” of the Electron
• JJ Thomson – (1897) – “found” small
negatively charged particles in the beam
of a cathode ray tube and called the
“corpuscles”. (We call them electrons)
He also modified a CRT
and found small positive
particles (protons)
Robert Millikan – (1909) – Oil Drop
Experiment
• Found the charge to
mass ratio of an
electron.
• From that calculated
the mass of an
electron.
• Me- = 9.109x10-31Kg
Ernest Rutherford – (1911) – Gold Foil
Experiment – “discovery” of the Nucleus
Rutherford led the experiment with the
assistance of Hans Geiger and Ernest Marsden
Walter Bothe (1930) & James
Chadwick (1932) – “discover” Neutron
In the 1930s, Bothe found that the
radiation emitted by beryllium when it is
bombarded with alpha particles was a
new form of penetrating high energy
radiation, which was later shown by
Chadwick to be neutrons.
Atomic Model Development
Dalton’s Model
Thomson’s Plum
Pudding Model
Quantum
Model
Bohr’s Model (next chapter)
Rutherford’s Model
Subatomic Particle Size
Charge
Electron -1
Proton
+1
Neutron 0
Mass #
0
1
1
Relative
Actual
Mass
Mass (Kg)
(amu)
0.0005486 9.109x10-31
1.007276
1.673x10-27
1.008665
1.675x10-27
Subatomic Particle Size
Charge
Electron -1
Proton
+1
Neutron 0
Mass #
0
1
1
Relative
Actual
Mass
Mass (Kg)
(amu)
-31
1836
electrons 9.109x10
0.0005486
equal
1.007276
1 proton
1.673x10-27
1.008665
1.675x10-27
Counting Atomic Particles
• Atomic Number (Z) –
the number of protons
• Mass Number (A) –
the total number of
protons and neutrons
• Number of Neutrons
=A-Z
Isotopes – atoms of the same element that
differ in mass due to a different number of
neutrons
• Nuclide – general
term for a specific
isotope of an element
• Carbon – 12
• Carbon – 14
Nuclear Symbol
?
?
?
Atomic Mass Units – amu- 1/12 the
mass of the carbon-12 atom.
• Used because atomic masses are so
small in Kg and g.
• 1 atom of Oxygen-16 = 2.657x10-23g
or
15.994915 amu
Mole – amount of a substance that
contains the same number of
particles as 12 g of carbon-12
Avogadro’s Number –
the number of particles
in one mole of a
substance – 6.022x1023
Units can be – atoms/mol or
molecules/mol or particles/mol
A mole is like a dozen (12) or a gross (144)
only it is a mole (6.022x1023)
Size of Avogadro ?
Molar Mass – mass in grams of
one mole of a pure substance.
• Numerically
equal to the
average atomic
mass
Carbon
Iron
Atomic
Mass
(amu)
12.0107
55.845
Molar
Mass (g)
12.0107
55.845
Conversions
grams – moles - atoms
6.022 x 1023 Atoms of each
Example: How many moles are in 22 grams of copper metal?
Example: How many atoms are in 22 grams of copper metal?
Homework
• Pages 89-90
• Numbers 2,4,6,8,11,15,17,18,19,21,22,23