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Unit 3
Atoms and their structure
How we started to think about
atoms
• Original idea came from Ancient Greece (400
B.C.)
– Democritus and Leucippus were Greek philosophers
How we started to think about
atoms
• Democritus looked
at the beach
– Made of sand
– If you cut sand
particles, you get
smaller sand
particles
How we started to think about
atoms
• There must be a smallest possible piece
– Called those pieces “Atomos” – not able to be
cut
Another Greek
• Aristotle - Famous philosopher
– All substances are made of 4 elements
• Fire - Hot
• Air - light
• Earth - cool, heavy
• Water - wet
• Blend these in different proportions to
get all substances
Who Was Right?
• None of the philosophers experimented to
determine who was right
– Greeks settled disagreements by argument
– Aristotle was a better debater - He won
– His ideas carried through middle ages
• Later on, alchemists tried to change
lead to gold (they did not understand
atoms)
Who’s Next?
• England in the late 1700’s - John Dalton
– Teacher who summarized results of his
experiments and those of others
– Elements are substances that can’t be broken
down
– In Dalton’s Atomic Theory, he combined the idea
of elements with that of atoms
Dalton’s Atomic Theory
 All matter is made of tiny indivisible
particles called atoms.
 Atoms of the same element are identical,
those of different atoms are different.
 Atoms of different elements combine in
whole number ratios to form compounds.
 Chemical reactions involve the
rearrangement of atoms. No new atoms are
created or destroyed.
Law of Definite Proportions
(part 3 in Dalton’s Theory)
• Each compound has a specific ratio of
elements
– It is a ratio by mass
– Water is always 8 grams of oxygen for each
gram of hydrogen
Law of Multiple Proportions
• If two elements form more than one
compound, the ratio elements in each
compound, is a simple whole number
– The ratio of the ratios is also a whole number
What?
• Water is 8 grams of oxygen per gram of
hydrogen
• Hydrogen peroxide is 16 grams of oxygen
per gram of hydrogen
• 16 to 8 is a 2 to 1 ratio
• This happens because you have to add a
whole atom, you can’t add a piece of an
atom
Parts of Atoms
• J. J. Thomson - English physicist, 1897
– Made a piece of equipment called a cathode ray
tube
• It is a vacuum tube - all the air has been
pumped out
• A limited amount of other gases are put
in and an electric current is applied to the
tube
Thomson’s Cathode Ray Tube
Voltage source
-
+
Metal Disks
Thomson’s Experiment
Voltage source

+
Passing an electric current makes a beam
appear to move from the negative to the
positive end
Thomson’s Experiment
Voltage source
+
 By adding an electric field he found that the
moving pieces were negatively charged
Thomson & his atomic model
• Discovered the electron
• He did not know where
positive charges were
• Said the atom was like
plum pudding
– A bunch of positive stuff,
with the electrons able to
be removed
Rutherford’s Experiment
• Ernest Rutherford - English physicist, 1910
– Believed the plum pudding model of the atom
was correct
– Wanted to see how big atoms are
– Used radioactivity
• Alpha particles - positively charged
pieces given off by uranium
• Shot them at gold foil which can be
made a few atoms thick
Rutherford’s experiment
• When the alpha particles hit a fluorescent
screen, it glows
• Here’s what it looked like
Lead
block
Fluorescent
Screen
Uranium
Gold Foil
Lead
block
Uranium
Fluorescent
Screen
Gold Foil
He expected that…
• The alpha particles would pass through
without changing direction very much
• Because…
– The positive charges were spread out evenly alone they would not be enough to stop the alpha
particles
What he expected
Because, he thought
the mass was evenly
distributed in the atom
What he got
How he explained it
• Atom is mostly empty
space
– There is a small dense,
positive piece at the
center
– Alpha particles are
deflected by it, if they
get close enough
+
+
Modern View
• The atom is mostly
empty space
– Two regions
• Nucleus - protons
and neutrons
• Electron cloud region where you
might find an
electron
Density and the Atom
• Since most of the particles went straight
through the gold foil, the atom was mostly
empty space
• Because the alpha particles turned so much,
the positive particles must have been heavy
• Small volume and big mass = big density.
• This small dense positive area is the nucleus
Subatomic particles
• Electron – located outside of nucleus, has
negative charge
• Proton - positively charged particles inside of
nucleus that are many times heavier than the
electron
• Neutron - no charge but about the same mass
as a proton, located inside of nucleus
Subatomic particles
Relative
Particle
mass Actual
Name Symbol Charge (amu) mass (g)
Electron e-1 1/1840 9.11 x 10-28
Proton
p+
+1
1
1.67 x 10-24
Neutron
n0
0
1
1.67 x 10-24
Structure of the Atom
• There are two regions
– The nucleus
• With protons and neutrons
• Positive charge
• Almost all the mass
– Electron cloud - most of the volume of an atom
• The region where the electron can be
found
Size of an atom
• Atoms are small
– Measured in picometers, 10-12 meters
– Hydrogen atom, 32 pm radius
– Nucleus is very tiny compared to atom
• If the atom was the size of a stadium, the
nucleus would be the size of a marble
• Radius of the nucleus is near 10-15 m
• Density near 1014 g/cm3
Counting the Pieces
• Atomic Number = number of protons
– # of protons determines kind of atom
– the same as the number of electrons in the
neutral atom
• Mass Number = the number of protons +
neutrons
– Includes all the particles with mass
– NOT found on the periodic table
What about when Electrons ≠ Protons
• Electrons may be gained or lost
– Gaining electrons gives a negatively charged ion
called an anion
– Losing electrons gives a positively charged ion
called a cation
What about when Electrons ≠ Protons
• Practice with ions…
– Magnesium makes ions with a 2+ charge. Are
electrons lost or gained? How many electrons are
moved?
– Fluorine makes ions with a 1- charge. Are electrons
lost or gained? How many electrons are moved?
– An ion has 13 p+ and 10 e-. Give the symbol and
charge for the ion.
– An ion has 34 p+ and 36 e-. Give the symbol and
charge for the ion.
Isotopes
• Dalton was wrong
–Atoms of the same element can have
different numbers of neutrons
• different mass numbers (will have
the same atomic number
• called isotopes
Symbols for Isotopes
• Nuclear Notation
– Contains the symbol of the element, the mass
number and the atomic number
Mass
number
Atomic
number
X
Symbols for Isotopes
• Hyphen Notation
– Contains the symbol (or name) of the element and
the mass number.
• carbon- 12
• carbon -14
• uranium-235
Symbols for Isotopes
• Find the
–number of protons
–number of neutrons
–number of electrons
–Atomic number
–Mass Number
–Name
24
11
Na
Symbols for Isotopes
• Find the
–number of protons
–number of neutrons
–number of electrons
–Atomic number
–Mass Number
–Name
80
35
Br
Symbols for Isotopes
–if an element has an atomic
number of 34 and a mass number
of 78 what is the
–number of protons
–number of neutrons
–number of electrons
–Symbol – Nuclear & Hyphen
notation
–Name
Symbols for Isotopes
 if
an element has 91 protons and
140 neutrons what is the
–Atomic number
–Mass number
–number of electrons
–Symbol – Nuclear & Hyphen
notation
–Name
Symbols for Isotopes
 if
an element has 78 electrons and
117 neutrons what is the
–Atomic number
–Mass number
–number of electrons
–Symbol – Nuclear & Hyphen
notation
–Name
Atomic Mass
• How heavy is an atom of oxygen?
– There are different kinds of oxygen atoms
– More concerned with average atomic mass
• Based on abundance of each element in
nature
• Don’t use grams because the numbers
would be too small
• Is not a whole number because it is an average
• are the decimal numbers on the periodic table
Measuring Atomic Mass
• Unit is the Atomic Mass Unit (amu)
–One twelfth the mass of a carbon-12
atom
• 6 p+ and 6 n0
–Each isotope of an element has its own
atomic mass
• we get the average atomic mass of an
element using weighted averages (need
mass & percent abundance)
Calculating averages
• You have five rocks, four with a mass of 50 g,
and one with a mass of 60 g. What is the
average mass of the rocks?
• Total mass = (4 x 50) + (1 x 60) = 260 g
• Average mass = (4 x 50) + (1 x 60) = 260 g
5
5
Calculating averages
• If 80% of the rocks were 50 grams and 20% of
the rocks were 60 grams what is the weighted
average mass of the rocks?
• Weighted Average =
(% as decimal x mass) +
(% as decimal x mass) + …
• Weighted Average =
(0.8 x 50 g) + (0.2 x 60 g) = 52 g