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Transcript
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
The Hellenic Market
Fire
~
Water
Earth
Air
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
+
+
Homework Questions
1.
2.
Make a table that lists the different
philosophers and scientists (mentioned
so far), identify the experiment they
conducted AND describe what they
discovered.
What is Dalton’s atomic theory? Which
parts are still hold true today?
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 (IONS)
– 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
Atomic Mass
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
Homework Questions
1.
2.
3.
4.
Describe the structure of an atom.
Compare the size of the nucleus to the
atom.
What is the difference between an ion
and an isotope?
Write nuclear notation and hyphen
notation for an element that has 24
protons, 24 electrons, and 26 neutrons.