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Chapter 3
Atoms: The Building Blocks of
Matter
Early History of Chemistry
400 BC
– The Greeks characterized matter into four
fundamental substances – Earth, Water,
Wind, and Fire. (What do these correspond to?)
– Greek philosopher, Democritus, proposed that
matter is composed of small indivisible
particles, called “atomos” (means
“uncuttable”)
Alchemy
For the next 2000 years, chemical history
was dominated by the pseudoscience
“alchemy”.
Some alchemists were mystics who
obsessed with turning cheap metals into
gold.
Other alchemists were serious scientists
and they discovered several elements and
developed many laboratory techniques.
Foundations of Quantitative
Chemistry
By the late 1700s, scientists had
discovered several laws related to the
composition of matter and the changes
that matter undergoes.
The Law of Conservation of Mass
Mass is neither created nor destroyed in
an ordinary chemical reaction or during a
physical change.
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Law of Definite Proportions
A given compound always contains
exactly the same proportion of elements
by mass regardless of the source of the
compound or the size of the sample.
For pure water:
2 atoms H = 2.0 amu
1 atom O = 16.0 amu
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Law of Multiple Proportions
For series of compounds containing the same
two elements, the ratios of the masses of the
first element that combine with a constant mass
of the second element can be reduced to small
whole numbers
H2O
H2O2
H2O
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H2O2
Mass H 2.0 g
2.0 g
Mass O 16.0 g
32.0 g
Dalton’s Atomic Theory
1808 – John Dalton published a theory to
explain these laws. There are 5 parts of
his theory.
1. Each element is made up of tiny particles
called atoms.
Atoms can be treated
like solid spheres
2. The atoms of a given element are identical.
Atoms of different elements are different in
some fundamental way.
Dalton’s Atomic Theory (cont)
3. Atoms cannot be created or destroyed
4. Atoms of different elements combine in
simple ratios to form chemical compounds
CO2
CCl4
5. Chemical reactions involve the rearrangement
of atoms, changing the way they are bound
together.
Subatomic particles
The atom is the smallest unit of an
element that retains the properties of that
element.
We now know that atoms contain
subatomic particles (protons, neutrons,
and electrons).
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Discovery of the Electron
In 1897, JJ Thomson discovered electrons
using cathode ray tubes.
Electric currents were passed through low
pressure gases. A ray of particles
travelled from the cathode to the anode.
This ray was shown to be deflected by a
negatively charged magnet.
He found the mass to charge ratio of the
electron
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More about the Electron
In 1909, Millikan determined the charge of
the electron.
He did this by suspending tiny oil drops
covered with electrons within a magnetic
field
He then used Thomson’s mass/charge
ratio to calculate the mass of an electron
Mass of electron = 9.11 x 10-31 kg
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Discovery of the Nucleus
In 1911, Ernest Rutherford conducted his
famous Gold Foil Experiment.
He directed a stream of alpha particles
(positively charged particles) towards a
very thin sheet of gold foil. Most of the
alpha particles passed through the sheet,
but some were deflected.
Proved the existence of a small, dense,
positively charged center of the atom (the
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nucleus)
Subatomic Particles Summary
name
Relative Relative
charge mass (amu)
proton
+1
1.007276
neutron
0
1.008665
-1
0.0005486
(1/1823)
electron
Estimated location
mass
(amu)
In
1
nucleus
In
1
nucleus
0
Outside
nucleus
Atomic Number
The atomic number of any atom is equal to
the number of protons in that atom.
The elements are arranged on the periodic
table in order of their atomic numbers.
How many protons in an atom of:
Helium? 2
26
Iron?
54
Xenon?
Mass Number
The mass number of an atom is equal to
the number of protons plus the number of
neutrons.
A nucleon is a term used to describe the
particles found in the nucleus. (The
number of nucleons in an atom is always
equal to the mass number of that atom).
The atomic mass shown on the periodic
table is NOT the mass number.
Atomic Notation
Mass number
A
Atomic Number
Z
X
Element
Symbol
Isotopes
Isotopes are atoms with the same number
of protons, but a different number of
neutrons
Most elements have several isotopes.
Ex: Helium
helium-3
3He
helium-4
4He
2
2
Average Atomic Mass
The average atomic mass of an element
takes into account the relative abundance
of each isotope of the element.
Ave mass = (mass1)(abund1) +
(mass2)(abund2) + etc
Example: A sample of lead contains the
following isotopes. 1.40% Pb-204,
24.10% Pb-206, 22.10% Pb-207, and
52.40% Pb-208. Find the average atomic
mass.
(204amu)(0.0140)+(206amu)(0.2410)+
(207amu)(0.2210)+(208amu)(0.5240) =
207.2 amu
The average atomic mass is the
mass listed on the periodic table.
Finding % abundance of isotopes
You will only be asked to do this if there
are only two isotopes of the element.
Since the percent abundances must add
up to 100% (1), you can set it up as
shown.
ave mass = (mass1)(x) + (mass2)(1-x)
Copper consists of two isotopes, Cu-63 and
Cu-65. Find the percent abundance of
each isotope.
63.55 amu = (63amu)(x) + (65amu)(1-x)
Solving for x:
x = 0.725
Therefore:
Cu-63 is 72.5% and
Cu-65 is 27.5%
The mole
A mole is a unit which measures the
number of particles in a sample.
1 mole = 6.02 x 1023 particles
This number is also called
Avogadro’s number
Molar mass
The mass of one mole of a substance is
called the molar mass.
Molar masses for atoms can be found on
the periodic table.
– The molar masses of some elements
Helium = 4.0 g/mol
Chromium = 52.0 g/mol
Iron = 55.8 g/mol
Molar mass of molecules
The molar mass of molecules can be
determined by adding up the masses of
the atoms composing the molecule.
– Ex: Find the molar mass of C6H12O6
6(12.0g)+12(1.0g)+6(16.0g) = 180 grams/mol
Mole conversions
w/Dimensional Analysis
You can use Avogadro’s number to
convert between moles and number of
particles.
6.02 x 1023 particles
2.50 moles x
1 mole
= 1.50 x 1024 particles
You can use molar mass to convert
between grams and moles.
35.0 grams H2O
1 mole H2O
18.0 grams H2O
= 1.94 moles H2O