Download Hein and Arena - faculty at Chemeketa

1
Early Thoughts
2
• The earliest models of the atom
were developed by the ancient
Greek philosophers.
Leucippus of Miletus (490-??? B.C.).
First to introduce the idea of the atom, an
indivisible unit of matter. This idea was later
extended by his student, Democritus.
Democritus (about 470-370 B.C.)
thought that all forms of matter were
made of tiny particles called “atoms”
from the Greek “atomos” indivisible.
3
According to Democritus atoms are:
• Unchangeable and
indivisible.
• Identical except
for their size and
shape.
• Always in motion.
4
Democritus imagined that atoms of iron were shaped like
coils, making iron rigid, strong, and malleable. Atoms of
fire were sharp, lightweight, and yellow.
5
• Aristotle (384-322 B.C.) rejected the
theory of Democritus and endorsed
that of Empedocles that stated that
matter was made of 4 elements: air,
earth, fire , and water.
6
• Empedocles (492-432 B.C.)
believed that these elements
have always existed in fixed
amounts, and that there two
major forces which act upon
these elements to both
create and destroy: Love
and Strife. According to
legend, he died by falling
into a volcano's crater after
failing to become a god as
he predicted.
7
– Aristotle’s influence dominated the
thinking of scientists and
philosophers until the beginning of
the 17th century
8
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Alchemical Symbols
antimony
copper
arsenic
gold
bismuth
iron
10
Alchemical Symbols
magnesium
mercury
platinum
potassium
phosphorus
silver
11
Alchemical Symbols
sulfur
tin
zinc
lead
12
Joseph Priestley
Antoine and Marie Lavoisier
13
Dalton’s Model
of the Atom
14
2000 years after Aristotle, John Dalton,
an English schoolmaster, proposed his
model of the atom–which was based on
experimentation.
15
16
Dalton’s Atomic Theory
1. Elements are composed of minute
indivisible particles called atoms.
2. Atoms of the same element are alike in
mass and size.
3. Atoms of different elements have
different masses and sizes.
4. Chemical compounds are formed by
the union of two or more atoms of
different elements.
Atoms under special circumstances can be decomposed.
Modern research has demonstrated that atoms are composed of
subatomic particles
17
Dalton’s Atomic Theory
5. Atoms combine to form compounds in
simple numerical ratios, such as one to
one, two to two, two to three, and so on.
6. Atoms of two elements may combine in
different ratios to form more than one
compound.
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Dalton’s atoms were individual particles.
Atoms of each element are alike in
mass and size.
19
Dalton’s atoms were individual particles.
Atoms of different elements are not alike
in mass and size.
20
H 2
=
O 1
H 1
=
O 1
Daltons atoms combine in specific ratios
to form compounds.
21
Composition of
Compounds
22
The Law of Definite
Composition
A compound always contains two
or more elements combined in a
definite proportion by mass.
23
Composition of Water
• Water always contains the same two
elements: hydrogen and oxygen.
• The percent by mass of hydrogen in
water is 11.2%.
• The percent by mass of oxygen in
water is 88.8%.
• Water always has these percentages.
If the percentages were different the
compound would not be water.
24
Composition of Hydrogen Peroxide
• Hydrogen peroxide always contains the same
two elements: hydrogen and oxygen.
• The percent by mass of hydrogen in hydrogen
peroxide is 5.9%.
• The percent by mass of oxygen in hydrogen
peroxide is 94.1%.
• Hydrogen peroxide always has these
percentages. If the percentages were different
the compound would not be hydrogen
peroxide.
25
The Law of Multiple Proportions
Atoms of two or more elements may
combine in different ratios to produce
more than one compound.
26
Combining Masses of Hydrogen and Oxygen
Mass
Hydrogen(g)
Mass
Oxygen(g)
Water
1.0
8.0
Hydrogen
Peroxide
1.0
16.0
Hydrogen
peroxide
has peroxide
twice as 16g
much2
mass of oxygen
in hydrogen
=
=
oxygenmass
(by mass)
as
does
water.
of oxygen in water
8g 1
27
Combining Ratios of Hydrogen and Oxygen
• Hydrogen peroxide has twice as many
oxygens per hydrogen atom as does
water.
• The formula for water is H2O.
• The formula for hydrogen peroxide is
H2O2.
28
Learning Check
Which pair of formulas illustrates the law of
multiple proportions?
a.
b.
c.
d.
CH3Cl and CH3OH
H2O and HOH
CuCl2 and CuBr
Na2O and Na2O2
29
Learning Check
Which pair of formulas illustrates the law of
multiple proportions?
a.
b.
c.
d.
CH3Cl and CH3OH
H2O and HOH
CuCl2 and CuBr
Na2O and Na2O2
30
The Nature of
Electric Charge
31
Surrounding the atomic nucleus are electrons. The name electron comes
from the Greek word for amber, a brownish-yellow fossil resin studied
by the early Greeks. They found that when amber was rubbed by a piece
of cloth, it attracted such things as bits of straw. This phenomenon,
known as the amber effect, remained a mystery for almost 2000 years.
In the late 1500s other materials that behaved like amber were called
“electrics”. The concept of electric charge awaited experiments by
Benjamin Franklin nearly two centuries later. Franklin experimented
with electricity and postulated the existence of an electric fluid that
could flow from place to place. An object with an excess of this fluid he
called electrically positive, and one with a deficiency of the fluid he
called electrically negative. The fluid was thought to attract ordinary
matter but to repel itself. We still follow Franklin’s lead in how we
define positive and negative electricity. Franklin’s 1752 experiment with
the kite in the lightning storm showed that lightning is an electrical
discharge between clouds and the ground. This discovery told him that
electricity is not restricted to solid or liquid objects and that it can travel
32
through a gas.
33
Ben Franklin (1746-52 ) flew kites to demonstrate
that lightning is a form of static electricity (ESD).
He would run a wire to the kite and produce sparks
at the ground, or charge a Leyden jar. This led
Franklin to invent the lightning rod. Franklin also
made several electrostatic generators with rotating
glass balls to experiment with. These experiments
led him to formulate the single fluid (imponderable
fluid) theory of electricity. Previous theories had
held there were two electrical fluids and two
magnetic fluids. Franklin theorized just one
imponderable electrical fluid (a fluid under
conservation) in the universe. The difference in
electrical charges was explained by an excess ( + )
or defect ( - ) of the single electrical fluid. This is
where the positive ( + ) and negative ( - ) symbols
come from in electrical science.
http://www.bing.com/videos/search?q=Leyden+Jar&Form=R5FD5#vie
w=detail&mid=FC14394778B561662C72FC14394778B561662C72
Leyden jar
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Discovery of Ions
35
• Michael Faraday discovered that
certain substances, when dissolved in
water, conducted an electric current.
• He found that atoms of some elements
moved to the cathode (negative
electrode) and some moved to the
anode (positive electrode).
• He concluded they were electrically
charged and called them ions (Greek
wanderer).
36
Michael Faraday
37
• Svante Arrhenius reasoned that an ion
is an atom (or a group of atoms)
carrying a positive or negative electric
charge.
• Arrhenius accounted for the electrical
conduction of molten sodium chloride
(NaCl) by proposing that melted NaCl
dissociated into the charged ions Na+
and Cl-.
Δ
NaCl → Na+ + Cl38
NaCl → Na+ + Cl• In the melt the positive Na+ ions moved
to the cathode (negative electrode).
Thus positive ions are called cations.
• In the melt the negative Cl- ions moved
to the anode (positive electrode). Thus
negative ions are called anions.
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Svante Arrhenius
40
Subatomic Parts
of the Atom
41
An atom is very Small
42
This
The diameter
is 1 to 5often
anbillionths
atom is 0.1oftoa
meter.
0.5 nm.
If
Even
the diameter
smaller particles
of this dot
thanis atoms
1 mm
exist.
then 10
These
million
are hydrogen
called subatomic
atoms
particles.
would form a line across the dot.
43
Subatomic Particles
44
Electron
45
In 1875 Sir William Crookes
invented the Crookes tube.
46
• Crookes tubes experiments led the way
to an understanding of the subatomic
structure of the atom.
http://www.youtube.com/watch?v=Sikzu09q6cc
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48
• Crookes tube emissions are called cathode
rays.
• Below are Crookes cathode-ray tubes. The
cathode-rays (streams of electrons) can be
clearly seen.
49
"Maltese Cross" Crookes Tube
Demonstrates that radiant matter is
blocked by metal objects
50
In 1897 Sir Joseph Thomson demonstrated
that cathode rays:
• travel in straight lines.
• are negative in charge.
• are deflected by electric
and magnetic fields.
• produce sharp shadows
• are capable of moving a
small paddle wheel.
51
Paddle Wheel
http://www.youtube.com/watch?v=eDFKyLZ8L8s
http://www.youtube.com/watch?v=Z61zCaAFky4
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Thomson’s Apparatus
batteries
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Thomson’s Lab
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J.J. Thomson determined and is given credit
for finding:
• The charge to mass
(e/m) ratio of the
cathode ray.
• The cathode ray was renamed the “electron”.
• Thomson “discovered”
the electron.
http://www.aip.org/history/mod/fission/fission1/01.html
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Can atoms be split apart? Does each atom have inner workings? Parts which
can be separated? Parts which can perhaps be put to some use? These
questions had already come to mind in 1898, when J. J. Thomson isolated the
electron. That was the first solid proof that atoms are indeed built of much tinier
pieces. Thomson speaks of the electron in this recorded passage...
Could anything at first sight seem more impractical than a body which
is so small that its mass is an insignificant fraction of the mass of an
atom of hydrogen, which itself is so small that a crowd of these atoms
equal in number to the population of the whole world would be too
small to have been detected by any means then known to science.
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Robert Millikan
• Determined the
charge of the
electron.
• Experiment called
the Oil Drop
Experiment.
57
58
Apparatus Used by Millikan
59
Modern Apparatus
60
Anti-Electron
Positron?
61
• Eugen Goldstein, a German physicist,
first observed protons in 1886:
• Thomson determined the
protons’ characteristics.
• Thomson showed that atoms
contained both positive and
negative charges.
• This disproved the Dalton
model of the atom which held
that atoms were indivisible.
62
63
Proton
64
Thomson’s Plum-Pudding Model of
the Atom
65
Ions
66
• Positive ions were explained by
assuming that a neutral atom loses
electrons.
• Negative ions were explained by
assuming that extra electrons can be
added to atoms.
67
Metals seem
to do this in
forming
chemical
bonds
When one or more electrons are lost
from an atom, a cation is formed.
68
Nonmetals
seem to do
this in
forming
chemical
bonds
When one or more electrons are added
to a neutral atom, an anion is formed.
69
Neutron
70
• James Chadwick discovered the neutron
in 1932.
• Its actual mass is
slightly greater than
the mass of a
proton.
71
Electric charge and mass of electrons, protons, & neutrons
Particle
symbol Actual charge
(coulombs)
Relative
charge
Actual
mass (g)
Relative
mass
(amu)
Electron
e–
–1.602  10–19
–1
9.11  10–28
0
Proton
p
+1.602  10–19
+1
1.67  10–24
1
Neutron
n
0
1.68  10–24
1
0
72
The Nuclear Atom
73
X-rays were discovered by Wilhelm
Röentgen in 1895
74
• Röentgen observed that a vacuum
discharge tube enclosed in a thin,
black cardboard box had caused a
nearby piece of paper coated with
the salt barium platinocyanide to
phosphorescence.
• From this and other experiments he
concluded that certain rays, which he
called X-rays, were emitted from the
discharge tube, penetrated the box, and
caused the salt to glow.
75
Calcite
standard Illumination
UV Illumination
76
77
78
• Radioactivity was discovered by Henri
Becquerel in 1896.
79
• Shortly after Röentgen’s discovery,
Antoine Henri Becquerel attempted to
show a relationship between X-rays
and the phosphorescence of uranium
salts.
• Becquerel wrapped a photographic
plate in black paper, sprinkled a sample
of a uranium salt on it, and exposed it
to sunlight.
80
• When Becquerel attempted to repeat
the experiment the sunlight was
intermittent.
• He took the photographic plate
wrapped in black paper with the
uranium sample on it, and placed the
whole setup in a drawer.
81
• Several days later he developed the
film and was amazed to find an intense
image of the uranium salt on the plate.
• He repeated the experiment in total
darkness with the same result.
• This proved that the uranium salt
emitted rays that affected the
photographic plate, and that these rays
were not a result of phosphorescence
due to exposure to sunlight.
82
Here is a photo of the discovery plate
83
• Two years later, in 1898, Marie Curie
coined the name radioactivity.
Radioactivity is the spontaneous emission of
particles and/or rays from the nucleus of an
atom.
84
85
Marie Curie, in a classic experiment, proved
that alpha and beta particles are oppositely
charged.
radiation passes between the
poles of an electromagnet
a radioactive source
was placed inside a
lead block
86
strongly deflected
to the positive pole.
Gamma rays are
not deflected by
the magnet.
Alpha rays are less
strongly deflected to the
negative pole.
three types of radiation are
detected by a photographic
plate
Marie Curie, in a classic experiment, proved
that alpha and beta particles are oppositely
charged.
Beta rays are
87
The Rutherford Experiment
88
Ernest Rutherford
89
• In 1899 Rutherford began to investigate
the nature of the rays emitted by uranium.
• He found two particles in the rays. He
called them alpha and beta particles.
90
• Rutherford in 1911 performed experiments
that shot a stream of alpha particles at a
gold foil.
• Most of the alpha particles passed through
the foil with little or no deflection.
• He found that a few were deflected at large
angles and some alpha particles even
bounced back.
91
Rutherford’s alpha particle scattering experiment.
92
93
94
95
• An electron with a mass of 1/1837 amu
could not have deflected an alpha
particle with a mass of 4 amu.
• Rutherford knew that like charges
repel.
• Rutherford concluded that each gold
atom contained a positively charged
mass that occupied a tiny volume. He
called this mass the nucleus.
96
• If a positive alpha particle approached
close enough to the positive mass it
was deflected.
• Most of the alpha particles passed
through the gold foil.
This led
Rutherford to conclude that a gold
atom was mostly empty space.
97
• Because alpha particles have relatively
high masses, the extent of the
reflections led Rutherford to conclude
that the nucleus was very heavy and
dense.
98
Deflection
Scattering
Deflection and scattering of alpha particles by positive gold nuclei.
99
Ideas about the atom were refined by one of Thomson's students,
Ernest Rutherford. He showed that the mass in an atom is not
smeared out uniformly throughout the atom, but is concentrated in a
tiny, inner kernel: the nucleus. Rutherford wanted to understand the
nucleus, not for any practical purpose, but because he was attracted
to the beauty of its simplicity. Fundamental things should be simple
not complex. Here is how he explains himself in 1931...
The bother is that a nucleus, as you know, is a very small thing,
and we know very little about it. Now, I had the opinion for a long
time, that's a personal conviction, that if we knew more about the
nucleus, we'd find it was a much simpler thing than we suppose,
that these fundamental things I think have got to be fairly simple.
But it's the non-fundamental things that are very complex usually. I
am always a believer in simplicity being a simple person myself.
100
• The gamma ray, a third type of emission from
radioactive material, was discovered by Paul
Villard in 1900.
101
Alpha, Beta, and Gamma Radiation
Name
Alpha
Particle
Symbol
Mass
(amu)
Charge

4
+2
e

1
1837
–1


0
0
Nuclide
Symbol
4
2
He
Beta
0
-1
Gamma Ray
0
0
102
General
Arrangement of
Subatomic Particles
103
• Rutherford’s experiment showed that an
atom had a dense, positively charged
nucleus.
• Chadwick’s work in 1932 demonstrated
the atom contains neutrons.
• Rutherford also noted that light,
negatively charged electrons were
present in an atom and offset the positive
nuclear charge.
104
• Rutherford put forward a model of the
atom in which a dense, positively
charged nucleus is located at the
atom’s center.
• The negative electrons surround the
nucleus.
• The nucleus contains protons and
neutrons
105
106
Atomic Numbers of
the Elements
107
• The atomic number of an element is
equal to the number of protons in the
nucleus of that element.
• The atomic number of an atom
determines which element the atom is.
108
Every atom with an atomic
number of 1 is a hydrogen atom.
Every hydrogen atom contains 1
proton in its nucleus.
109
Every atom with an atomic
number of 6 is a carbon atom.
Every carbon atom contains 6
protons in its nucleus.
110
atomic
number
Every atom with an
atomic number of 1
is a hydrogen atom.
1 proton in the
nucleus
111
atomic
number
Every atom with an
atomic number of 6
is a carbon atom.
6 protons in the
nucleus
112
atomic
number
Every atom with an
atomic number of
92 is a uranium
atom.
92 protons
in the
nucleus
113
Isotopes of the
Elements
114
• Atoms of the same element can have
different masses.
• They always have the same number of
protons, but they can have different
numbers of neutrons in their nuclei.
• The difference in the number of neutrons
accounts for the difference in mass.
• These are isotopes of the same element.
115
Greek roots isos (equal) and
topos (place). Hence: "the
same place," meaning that
different isotopes of a single
element occupy the same
position on the periodic table.
116
Isotopes of the Same
Element Have
Equal numbers of protons
Different numbers of
neutrons
117
Isotopic Notation
Mass number is also the number
of nucleons in the nucleus.
Nucleons = protons and/or neutrons
118
Relationship Between Mass
Number and Atomic Number
119
The mass number minus the atomic
number equals the number of neutrons in
the nucleus.
mass
number
atomic
number
109
47
Ag
atomic
mass number number
109
47
=
=
number of
neutrons
62
120
Isotopic Notation
8 protons + 8 neutrons
16
O
8
8 protons
121
Isotopic Notation
8 protons + 9 neutrons
17
O
8
8 protons
122
Isotopic Notation
8 protons + 10 neutrons
18
O
8
8 protons
123
Hydrogen has three isotopes
1 proton
1 proton
1 proton
0 neutrons
1 neutron
2 neutrons
124
Examples of Isotopes
Element Protons
Electrons Neutrons Symbol
Hydrogen
Hydrogen
Hydrogen
1
1
1
1
1
1
0
1
2
1
1
2
1
3
1
Uranium
Uranium
92
92
92
92
143
146
235
92
U
238
92
U
Chlorine
Chlorine
17
17
17
17
18
20
H
H
H
35
17
37
125
17
Cl
Cl
Atomic Weight
126
• The mass of a single atom is too small to
measure on a balance.
• Using a mass spectrometer, the mass of
the hydrogen atom was determined.
127
A Modern Mass Spectrometer
Positive ions
formed from
sample.
Electrical field at
slits accelerates
positive ions.
From the intensity and positions
of the lines on the mass
spectrogram, the different
isotopes and their relative
amounts can be determined.
Deflection of
positive ions
occurs at magnetic
field.
A mass spectrogram
is recorded.
128
A typical reading from a mass spectrometer. The two
principal isotopes of copper are shown with the
129
abundance (%) given.
Using a mass spectrometer, the mass of one
hydrogen atom was determined to be 1.673
x 10-24 g
130
This number is very small.
small
small
small
small
small
small
small
small
small
small
small
small
small
small
small
small
small
small
small
131
The mass of a hydrogen atom is
very small.
Numbers
of this
sizeproblem
are too small
for of
To overcome
this
a system
practical
use.
relative atomic
weights using “atomic
mass units” was devised to express the
masses of elements using simple
numbers.
1.673 x 10-24 g
132
The standard to which the masses of all
other atoms are compared to was
originally chosen to be the most
abundant isotope of hydrogen.
1
1
H
133
A mass of exactly 1 atomic mass units
(amu) was assigned to
1
1
H
134
The standard to which the masses of all
other atoms are compared to was later
chosen to be the most abundant isotope
of oxygen.
16
8
O
135
A mass of exactly 16 atomic mass units
(amu) was assigned to
16
8
O
136
The standard to which the masses of all
other atoms are compared to was chosen
to be the most abundant isotope of
carbon.
12
6
C
137
A mass of exactly 12 atomic mass units
(amu) was assigned to
12
6
C
138
1
1 amu is defined as exactly equal to
12
the mass of a carbon-12 atom
1 amu = 1.6606 x 10-24 g
12
6
C
139
Average atomic weight 1.00797 amu.
H
140
Average atomic weight 39.098 amu.
K
141
Average atomic weight 248.029 amu.
U
142
Average Relative
Atomic Weight
143
• Most elements occur as mixtures of
isotopes.
• Isotopes of the same element have
different masses.
• The listed atomic mass of an element is
the average relative mass of the isotopes
of that element compared to the mass of
carbon-12 (exactly 12.0000…amu)
144
To calculate the atomic mass multiply the
atomic mass of each isotope by its percent
abundance and add the results.
Isotope
63
29
Cu
65
29
Cu
Isotopic mass
(amu)
Abundance
(%)
62.9298
69.09
64.9278
30.91
Average
atomic mass
(amu)
63.55
(62.9298 amu) 0.6909 = 43.48 amu
(64.9278 amu) 0.3091 = 20.07 amu
63.55 amu
145
Isotope Practice
(Fill-in the Blanks)
symbol
atomic no
mass no
#e
8
16
10
Pt
Pt
#n
#p
117
PP
30
3
30 
–3
53
74
48
34
36
45
Ca+22
Ca
40
146
Isotope Practice
(Fill-in the Blanks)
symbol
atomic no
mass no
#e
#n
O 2
8
16
10
8
Pt
78
195
78
30
18
53
16
8
195
78Pt
PP
30
3
30 –3
15
15
117
15
#p
8
78
15
I
53
127
Kr
36
84
36
48
36
Se
34
79
34
45
34
20
20
127
53
84
36
79
34
Ca
40
2
40Ca+2
20
20
40
18
74
53
147
148