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Chapter 4
Atoms and their structure
History of the atom
Not the history of atom, but the idea of the
atom.
 Original idea Ancient Greece (400 B.C.)
 Democritus and Leucippus- Greek
philosophers.

History of Atom
Looked at beach Smallest possible
piece?
 Made of sand
 Cut sand - smaller Atomos - not to be cut
sand

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?
Did not experiment.
 Greeks settled disagreements by
argument.
 Aristotle was a better debater - He won.
 His ideas carried through middle ages.
 Alchemists tried to change lead to gold.

Who’s Next?
Late 1700’s - John Dalton- England.
 Teacher- summarized results of his
experiments and those of others.
 Elements substances that can’t be
broken down
 In Dalton’s Atomic Theory
 Combined 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.
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

Thomson’s Experiment
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
Thomson’s Experiment
Voltage source
+
 By adding an electric field he found that the
moving pieces were negative
Thomson’s Experiment
Used many different metals and gases
 Beam was always the same
 By the amount it bent he could find the
ratio of charge to mass
 Was the same with every material
 Same type of piece in every kind of
atom

Thomsom’s Model
Found the electron.
 Couldn’t find
positive (for a while).
 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 they 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 florescent
screen, it glows.
 Here’s what it looked like (pg 72)

Lead
block
Uranium
Flourescent
Screen
Gold Foil
He Expected
The alpha particles to pass through
without changing direction very much.
 Because…
 The positive charges were spread out
evenly. Alone they were not enough to
stop the alpha particles.

What he expected
Because
Because, he thought the mass was
evenly distributed in the atom
Because, he thought
the mass was evenly
distributed in the atom
What he got
How he explained it
Atom is mostly empty.
 Small dense,
positive piece
at 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 cloudregion where you
might find an
electron.

Density and the Atom
Since most of the particles went
through, it was mostly empty.
 Because the pieces turned so much, the
positive pieces were heavy.
 Small volume, big mass, big density.
 This small dense positive area is the
nucleus.

Other pieces
Proton - positively charged pieces 1840
times heavier than the electron.
 Neutron - no charge but the same mass
as a proton.
 Where are the pieces?

Subatomic particles
Relative Actual
mass (g)
Name Symbol Charge mass
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.
-12 meters.
 Measured in picometers, 10
 Hydrogen atom, 32 pm radius.
 Nucleus 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 10-15m.
 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.
 All the things with mass.
 NOT on the periodic table

Symbols

Contain the symbol of the element, the
mass number and the atomic number.
Mass
number
Atomic
number
X
Symbols

Find the
– number of protons
– number of neutrons
– number of electrons
– Atomic number
– Mass Number
– Name
24
11
Na
Symbols
 Find
the
–number of protons
–number of neutrons
–number of electrons
–Atomic number
–Mass Number
– Name
80
35
Br
Symbols
 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
–Complete symbol
– Name
Symbols
 if
an element has 91 protons and
140 neutrons what is the
–Atomic number
–Mass number
–number of electrons
–Complete symbol
– Name
Symbols
 if
an element has 78 electrons and
117 neutrons what is the
–Atomic number
–Mass number
–number of protons
–Complete symbol
– Name
Unit 3 notes




Ions – charged atoms
Most atoms in their natural state are not stable.
In order to understand stability, we have to look
at the last energy level the electrons fill. This
shell is called the valence shell. If that shell is
full the atom is happy, if not, then the atom will
go out and react to make itself full.
If the atom needs to steal electrons to become
stable it will, if it needs to give up electrons to
become stable it will, and some have the ability
to do both.
Rule is this, the atom will do whatever is
easiest.
Unit 3 notes
Ex. Flourine :2 in the first level 7 in the
second. That second level can hold 8,
so it needs one more, so it will go out
and steal one from something to make
itself happy.
 Now the PNE will change from 9, 10, 9
to 9, 10, 10, now the protons and
electrons are not equal to each other
and so the overall charge is not zero but
now -1. So the flouride ion has an F-1
overall charge.

Unit 3 notes


Where does it get this extra electron from?
If you look at Lithium, the PNE for lithium is 3, 4, 3 2
electrons in the first level and 1 in the second. If
lithium wants to fill its second level it would need to
steal 7 electrons, very difficult. But if it could dump it
to flourine, then the first level (which is now the only
level) would be full and everything could be stable.

Lithium ion would have a PNE of 3, 4, 2 and because
it gave up an negative particle would have a charge of
Li +1

These go hand in hand, you have to have one thing
giving it up if you want to have something take it in. It
is the backbone for every chemical reaction.
Unit 3 notes






If the atom gives up an electron we say that it
has been oxidized (overall charge increases)
If the atoms takes in an electron we say that it
has been reduced (overall charge decreases)
These come from Benjamin Franklin’s names
of oxidation and reduction during a chemical
reaction.
Cation- + ions
Anion - - ions
Something is unique about C though, it can do
either gain 4 or lose 4, so it can have a charge
of + or – 4 depending on whatever it needs to
do.
Isotopes
Dalton was wrong.
 Atoms of the same element can have
different numbers of neutrons.
 different mass numbers.
 called isotopes.

Naming Isotopes
Put the mass number after the name of
the element.
 carbon- 12
 carbon -14
 uranium-235

Relative Abundance
The atomic mass given to you is the
overall average of all of the isotopes
found.
 It is based on the relative abundance of
each of the isotopes found in nature.
 For example, Cu has two isotopes:
Cu-63 and Cu-65
 Because the average is closer to 63,
then Cu-63 is more abundant in nature.

Relative Abundance

To calculate the average of all the
isotopes, use the atomic mass and the
decimal form of the percent abundance
to find the mass contributed by each
isotope and then add them together.
Relative Abundance
What??????
Relative Abundance
Take the mass of one isotope and
multiply it by its percentage that exists
in nature (just use the decimal form)
 Ex. You have element X-10 that is
19.91% abundant in nature and element
X-11 that is 80.80% abundant in nature.
10 (.1991) = 1.991
11 (.8080) = 8.888
Now add them up to get the average:
1.991 + 8.888 = 10.879 amu

Relative Abundance
Now you try it!
 Calculate the average atomic mass of
strontium. Here are the relative
abundances:
 Sr-84 .960%
 Sr-86 9.86%
 Sr-87 7.10%
 Sr-88 82.08%

Relative Abundance

Bromine has two isotopes Br- 79 and
Br- 81. If the average mass of bromine
is 79.9, calculate the relative
abundances?
Relative Abundance

Lead has 4 isotopes. Pb-204 Pb-206
Pb-207 and Pb-208. If the relative
abundance of Pb-204 is 1.4% and the
relative abundance of Pb-206 is 24.1%
Calculate the relative abundance of the
other two isotopes.
Relative Abundance

Element X has three isotopes X-100, X102 and X-104. If the average mass of
element X is 103.2 amu and the relative
abundance of X-104 is 68.45%, what
are the relative abundances of the other
two isotopes?
Symbols

Contain the symbol of the element, the
mass number and the atomic number.
• A Z X  0 -1 e + 214 83 Bi
• 140 56 Ba  140 57 La + A Z X
• 222 86 Rn  4 2 He + A Z X
•A Z Po  A Z X + 206 Z Pb
1
2
3
4
5. Potassium- 42 has a half-life of 12.4 hrs. How
much of a 560 g sample remains after 74.4 hrs?
6. If the half-life of iodine-131 is 8.1 days, how long
will it take a 50 g sample to decay to 12.5 g?
7.What is the half-life of a 100 g sample of nitrogen16 that decays to 6.25 g in 28.8 s?
8. C-14 has a half-life of 5730 yrs, if after 5730 yrs 35
g of C-14 remain, what was the original amount?
Bean Lab
Solve for the average mass of the element “Bean”
- In nature to find average mass, we take the % of each isotope x the
mass of one isotope.
- Directions: Split beans up into their 12 isotopes (varieties, call them
whatever you want)
- Count the number of each bean and record. (#1) for each one (12
lima, 10 split peas etc.)
- Count total number of beans record that. (#2)
- Record mass of the total of each isotope (ex. Put all lima beans on
scale and get the total mass.) (#3) for each one
- Find the % of each bean (#1) for each/ (#2) = (4) for each (take # of
lima beans/total then repeat for each type)
- Find average mass of each bean (#3)/(#1) for each = (#5) for each
(take total mass of lima/ # of lima, then repeat for each one)
- Find the average mass of all the isotopes:
Millikan’s Experiment
Atomizer
+
-
Oil
Microscope
Metal Plates
Millikan’s Experiment
Atomizer
Oil droplets
+
-
Oil
Microscope
Millikan’s Experiment
X-rays
X-rays give some drops a charge by knocking off
electrons
Millikan’s Experiment
+-
Millikan’s Experiment
-
-
+
+
They put an electric charge on the plates
Millikan’s Experiment
-
-
+
+
Some drops would hover
Millikan’s Experiment
-
-
-
-
-
-
-
+
Some drops would hover
+
+
+
+
+ +
+
Millikan’s Experiment
-
-
+
+
From the mass of the drop and the charge on
the plates, he calculated the charge on an electron
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.

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 has its own atomic mass
 we get the average using 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
 Average mass = 4 x 50 + 1 x 60 = 260 g
5
5
5

Calculating averages
Average mass = 4 x 50 + 1 x 60 = 260 g
5
5
5
 Average mass = .8 x 50 + .2 x 60
 80% of the rocks were 50 grams
 20% of the rocks were 60 grams
 Average = % as decimal x mass +
% as decimal x mass +
% as decimal x mass +

Atomic Mass

Calculate the atomic mass of copper if
copper has two isotopes. 69.1% has a mass
of 62.93 amu and the rest has a mass of
64.93 amu.
Atomic Mass
Magnesium has three isotopes. 78.99%
magnesium 24 with a mass of 23.9850
amu, 10.00% magnesium 25 with a mass of
24.9858 amu, and the rest magnesium 25
with a mass of 25.9826 amu. What is the
atomic mass of magnesium?
 If not told otherwise, the mass of the
isotope is the mass number in amu

Atomic Mass
Is not a whole number because it is an
average.
 are the decimal numbers on the periodic
table.

Law of Definite Proportions (#3)
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.
