Download Atoms - RCSD

1
2
3
• http://www.youtube.com/watch?v=WjjmzM
-YV1s
• Molar eclipse of the heart
• http://www.youtube.com/watch?v=oIkC7S
RqXP0
4
5
1. The process which a heavy nucleus is split into
lighter nuclei is called __________.
1. The process of combining nuclei to form a
bigger and heavier nucleus: ___________.
1. The bonus result of both reactions is ___, which
is why we are VERY interested in them.
1. One problem with nuclear fission is ____.
6
Atoms:
The Building Blocks of Nature
ἄτομος
As early as 400 BC Democritus
had an idea of matter… But it
couldn’t be tested.
8
18th Century… that’s when it all
started to come together…
9
th
18
Century thinking…
• Chemists accepted that an element couldn’t
be broken down by ordinary chemical
means.
• They assumed that these elements combined
to form compounds…
• They just couldn’t figure out exactly how
the different substances could combine with
one another to form new ones…
10
• The foundation of modern chemistry dates
to this time when scientists finally began to
give rules to how matter interacts.
11
It all led to…
• An emphasis on
analysis of
chemical reactions…
• Improved equipment like balances so mass
was more accurately
…
• The discovery of several basic laws.
12
Three basic laws…
• Law of Conservation of Mass
• Law of Definite Proportions
• Law of Multiple Proportions
13
Law of Conservation of Mass
Mass isn’t created or destroyed during
ordinary chemical reactions or physical
changes.
Carbon + Oxygen = CO
Mass x
Mass y
Mass x + Mass y
14
And the reverse holds true…
CO =
(Mass x + Mass y)
Carbon + Oxygen
Mass x
Mass y
15
Law of Definite Proportions
• A compound contains the same elements in
exactly the same proportions by mass - no
matter how much or how little there is of it.
16
And in simpler language:
Every chemical compound has one and only
one chemical formula.
No matter what process you use to make
water, the formula will always be H2O.
17
Basically
The Law of Definite Proportions states that no
matter how you make a chemical compound,
it's got the same ratio of elements.
18
• An example: Whether you make H2O by
combining H + O or by decomposing H2O2
(hydrogen peroxide),
• the resulting H2O will still be 1 part
by mass of H to eight parts by mass of O.
19
Sodium chloride - NaCl
20
• NaCl has a fixed proportion of elements.
• It is always 39.43% by mass Sodium and
60.66% by mass Chlorine.
The exact same proportions of Na and Cl must
always be combined in order for salt to be
created.
21
H2SO4 - Sulfuric Acid
• Sulfuric acid is made up of the individual
elements of H, S, and O.
• The chemical compound is written H2SO4.
• The same proportions of H, S, and O must
be combined to create sulfuric acid.
22
Law of Multiple Proportions
If two or more different compounds are
composed of the same two elements, then the
ratio of the masses of the second element
combined with a certain mass of the first
element is always a ratio of small whole
numbers.
23
Once again, in a simpler way…
If two elements can combine to form more
than one chemical compound, the ratio of the
mass of one element that combines with a
fixed mass of the other element will be a
whole number ratio for the compounds.
And sense that doesn’t make any sense…
24
Basically…
• The same two elements can combine in
multiple ways to create different
compounds.
• Any time two elements can form more than
one compound with each other, a particular
ratio between the masses takes place.
25
For example… yes, please!
C + O = CO
Carbon monoxide molecules are always
composed of 1 C and 1 O atom.
C + O + O = CO2
Carbon dioxide molecules are always
composed of 1 C and 2 O atoms.
26
Think about it like this:
Definite Proportions means no matter how
you make a compound – it’s always the same
ratio of elements.
Multiple Proportions has to do with the
different compounds you get when you
combine the
.
27
And back to our story…
28
• So John Dalton, an English schoolteacher,
in 1808, proposed an explanation that pulled
together all these laws.
• He reasoned that…
29
Dalton’s Atomic Theory (1803)
1. Elements are composed of extremely small particles
called atoms.
2. All atoms of a given element are identical, having the
same size, mass and chemical properties. The atoms
of one element are different from the atoms of all
other elements.
3. Atoms cannot be subdivided, created, or destroyed.
4. Compounds are composed of atoms of more than
one element. In any compound, the ratio of the
numbers of atoms of any two of the elements present
is either an integer or a simple fraction.
5. A chemical reaction involves only the separation,
combination, or rearrangement of atoms; it does not
result in their creation or destruction.
Or in other words…
1) Everything is made of atoms. (TRUE)
2) All atoms of an element are identical in
every way (Not true due to isotopes)
3) Atoms of different elements are different
(true)
4) Atoms can’t be broken (true for chemical
reactions, but not for nuclear ones)
31
Atoms combine in whole number ratios to
form compounds (i.e., you can’t have half an
atom in a compound) This is true…
5) In chemical reactions, atoms are
rearranged. (true)
32
• Not all aspects of Dalton’s atomic theory
have proven to be correct.
1)Atoms ARE divisible into smaller particles
called subatomic particle: protons, electrons
& neutrons.
2)A given element can have atoms with
different masses (isotopes).
33
Dalton’s Atomic Theory (1803)
1. Elements are composed of extremely small particles
called atoms.
2. All atoms of a given element are identical, having the
same size, mass and chemical properties. The atoms
of one element are different from the atoms of all
other elements.
3. Atoms cannot be subdivided, created, or destroyed.
4. Compounds are composed of atoms of more than
one element. In any compound, the ratio of the
numbers of atoms of any two of the elements present
is either an integer or a simple fraction.
5. A chemical reaction involves only the separation,
combination, or rearrangement of atoms; it does not
result in their creation or destruction.
Dalton turned Democritus’ idea into a
scientific theory that could be tested by
experiment.
•But today we know that atoms are divisible
into even smaller particles (protons, neutrons,
electrons… quarks…leptons…)
•We also know that a given element can have
atoms with different masses. (isotopes)
35
The Law of Conservation of
Mass still holds true!
• Atomic theory hasn’t been discarded – only
modified!
36
The important concepts of it:
1) All matter (elements) is composed of
atoms
2) Atoms of any one element differ in
properties from atoms of another element.
3) A chemical reaction involves separation,
combination or rearrangement of atoms
that are not created or destroyed.
37
Dalton’s Atomic Theory (1808)
Law of Multiple Proportions
38
16 X
+
8Y
Law of Conservation of Mass
8 X2Y
Can we see atoms?!
. . . kind of . . .
42
Atomic force microscopy (AFM)
43
X-ray diffraction
44
In a nutshell
• Law of Conservation of Mass: The amount
of stuff you form in a reaction is equal to
the amount of stuff you started with.
45
The Law of Definite Proportions
Every chemical compound has one and only
one chemical formula. For example, no
matter what process you use to make water,
the formula will always be H2O.
46
Law of Multiple Proportions
• The same two elements can combine in
multiple ways to create different
compounds:
• H2O and H2O2
• NO and NO2
• CO and CO2
47
H2O and H2O2
• In the first compound, the amount of O
needed to combine with 2 g of H is 16 g.
• In the second compound, the amount of O
needed to combine with 2 g of H is 32
grams.
• The ratio of 32:16 works out to a 2:1 ratio,
it follows this law.
48
Bellwork #4 – 10/15
1. A chemical compound contains the same
elements in exactly the same proportions by
mass. Ibuprofen = C13H18O2.
2. H and O combine to form H2O. They also
combine to form H2O2.
3. The amount of stuff you form in a reaction
is equal to the stuff you started with.
49
Bellwork #5: 10/17 and check
1) Give 3 main concepts in Dalton’s
atomic theory.
2) 180.18 g of glucose, C6H12O6,
always contains 72.06 g C, 12.12 g H
and 96.00 g O. What % of each of these
elements is present in glucose?
50
• 1.
CaO + CO2
CaCO3
where the mass of the reactants = mass of
the products
•
• 2. N + O can form NO.
•
They can also form NO2.
•
•
H + O can form H2O.
•
They can also form H2O2.
•
51
3. To make table salt, 1 atom of Na will combine
with 1 atom of Cl. Always.
To make water, 2 atoms of H will combine with 1
atom of O. Always.
52
Bellwork #1
1. Law of Definite Proportions, any 2
samples of KCl have the same ratio of
elements.
2. Law of Conservation of Mass: When Na,
H and O form a compound, the mass of the
compound is equal to the sum of the
masses of the individual elements.
3. The ratio of O to C when 32 g of O
53
combine with 12 g of O = 2:1
Section 2:
The Structure of the Atom
54
Main Ideas of S.2:
1) Atoms contain positive and negative
particles.
2) Atoms have small, dense positivelycharged nuclei.
3) A nucleus contains protons and neutrons.
4) The radii of atoms are expressed in
picometers.
55
Even when the pieces look
different...
56
What I mean is…
• Dalton thought atoms were indivisible,
investigators in the late 1880s proved
otherwise.
• Atoms are actually composed of smaller
particles and the number and arrangement
of these particles determine the atom’s
chemical properties.
57
We’ll get to this in a bit
**Spoiler alert**:
All atoms consist of 2 regions. The nucleus is
a very small region located at the center of an
atom.
The nucleus contains at least one positively
charged particle (the proton) and usually one
or more neutral particles called neutrons.
Surrounding the nucleus is a region occupied
by negatively charged particles (electrons).
58
• So overall, people seemed pretty happy with
Dalton’s laws. That is until his idea of what
atoms are like was disproved by…
59
Mr. J.J. Thomson and his cathode
ray tube!
60
This is what happened
So… Thomson was experimenting with these
cathode ray tubes. He found that when he
connected these long hollow tubes to
batteries, a beam of light would go from one
end to another.
61
Since he had a lot of time on his hands…
He decided to figure out what the deal was
with the light. After all, if there was nothing
in the tube to start with, where’d the light
come from??
He figured it must come from the electrodes –
since the electrodes were made of atoms…
The atoms must be coming apart!
62
Then he put a magnet on it…
When he did this, he found that the beam
would bend toward the positive side of the
magnet and away from the negative side.
From this he figured that the beam must
contain very small particles from the atom and
that they must have negative charge.
63
• These experiments were carried out in glass
tubes hooked up to a vacuum pump – and it
is called a
.
64
65
• Did you notice the difference between the
flow of electrons (the cathode ray) and the
conventional direction of electric current?
66
• This figure depicts a cathode ray, created by
electrons, flowing from the negatively
charge cathode to the positively charged
anode.
• The electric current flows from the anode to
the cathode.
67
This is what they found…
• When a current was passed through the
tube, the surface of the tube directly
opposite the cathode glowed.
• The ray traveled from the cathode to the
anode.
68
• Cathode rays were deflected by a magnetic
field in the same way as a wire carrying
electric current – which they knew to have a
negative charge. <hmmm…>
• The rays were deflected away from a
negatively charged object…
69
Cathode Ray Tube
71
2+2=…
The particles that compose cathode rays are
negatively charged.
…Thank you to Mr. J. J. Thomson!
In 1897, this English physicist performed the
experiments to measure the charge-to-mass
ratio of an electron!
72
Pardon me, sir, but what does
that mean?
• He was able to measure the ratio of the
charge of the cathode ray particles to their
mass.
• He found that this ratio was always the
same, regardless of the metal used to make
the cathode OR the nature of the gas inside
the tube.
73
Conclusion, please:
• All cathode rays must be composed of
identical negatively charged particles,
which were named
74
Charge and Mass of the Electron
• It was concluded that electrons are present
in atoms of all elements.
• The ol’ cathode ray experiments provided
evidence that atoms are divisible and that
one of the atom’s basic components is the
negatively charged electron. <whoa>
75
Once again… in plain speech,
please.
• Thomson’s experiment also revealed that
the electron has a very large charge to mass
ratio.
• This means he didn’t know the charge and
he didn’t know the mass, but he understood
the ratio!
76
Cathode Ray Tube Experiment
J.J. Thomson
(1906 Nobel Prize in Physics)
Proposed Plum Pudding Model
of the atom
Thomson’s Plum Pudding Model
He believed that the negative electrons were
spread evenly throughout the positive charge
of the rest of the atom.
Kind of like seeds in a watermelon.
But the experiments continued… and
.
78
Still…
Ye olde Plum Pudding model was an
important step forward in our modern
understanding of the atom.
It represents the first time scientists tried to
incorporate the idea that atoms were not
indivisible.
79
The next part of the story!
Robert A. Millikan
80
American physicist
Robert A. Millikan, 1909
His experiments around 1909 measured the
charge of the electron. <amazing!!!>
From this information, scientists were able to
determine the mass of an electron!
81
Oil Drop Experiment
• http://www.kentchemistry.com/links/Atomi
cStructure/Millikan.htm
• Let’s watch it…
• It involved suspending oil drops between 2
charged plates.
82
The oil-drop experiment
83
• So it didn’t measure the mass of the
electron directly but the charge of the
electron.
• Scientists used Thomson’s charge to mass
ratio and the charge Millikan determined,
and they were able to calculate the mass of
an electron.
84
Oil Drop Experiment
R. A. Millikan
(1923 Nobel Prize in Physics)
1.
Measured charge of e-
2.
Calculated mass of e-
Why is it necessary to use experiments
such as those of J.J. Thomson and
Robert A. Millikan to infer information
about electrons?
86
It’s necessary to observe the
behavior of electrons in
experiments like these b/c
electrons themselves are far
too small to see.
87
And then…
Two other inferences were made about atomic
structure:
1) Because atoms are electrically
,
they must contain a positive charge to
balance the negative electrons.
2) Because electrons have so much less mass
than atoms, atoms must contain
that account for most of their
88
mass!
Crazy trivia time
The ratio of the mass of an electron to that of
a proton is roughly the same as the mass of a
house cat to that of an elephant.
89
90
Lord Ernest Rutherford
• Mr. Rutherford was a scientist who liked to
play with radioactive stuff.
• His favorite radioactive thing to play with
was alpha particles (He nuclei).
• One day he decided to shoot a bunch of
alpha particles at a really thin piece of gold
foil...
91
Atoms have small dense
positively charged nuclei.
• It was in 1911, in New Zealand.
• Ernest Rutherford and his associates Hans
Geiger and Ernest Marsden bombarded a
thin piece of gold foil with fast-moving
alpha particles (positively charged particles
with about 4 times the mass of a hydrogen
atom).
92
• They assumed that mass and charge were
uniformly distrubuted throughout the atoms
of the gold foil, like the plum pudding
model.
• They expected the alpha particles to pass
through with only a slight deflection.
• For the vast majority of the particles, this
was the case.
93
HOWEVER…
• When they checked for wide-angle
deflections, they were shocked to find that
roughly 1 in 8000 of the alpha particles had
actually been deflected back toward the
source!
94
Gold Foil Experiment
Ernest Rutherford
(1908 Nobel Prize in Chemistry)
What in the world had happened?
They finally came up with an explanation.
Rutherford reasoned that the deflected alpha
particles must have experienced some
powerful force within the atom.
He figured that the source of this force must
occupy a very small amount of space (because
so few of the particles had been affected).
96
• He concluded that the force must be caused
by a very densely packed bundle of matter
with a positive electric charge. Rutherford
called this positive bundle of matter the
nucleus.
• And that the atom is mostly empty space.
97
• He had also discovered that the volume of a
nucleus was very small compared with the
total volume of an atom.
• In fact, if the nucleus were the size of a
marble, then the size of the atom would be
about the size of a football field!
98
For this we will always know Rutherford as
the man with the gold foil.
And the bad model of the atom, because we
now know that this model isn’t right. Either.
99
So where were the electrons??
• This questions wasn’t answered until
Rutherford’s STUDENT, Niels Bohr,
proposed a model in which electrons
surrounded the positively charged nucleus
as the planets surround the sun.
• (We’ll discuss Bohr’s model a little later on…)
100
Gold Foil Experiment
Ernest Rutherford
(1908 Nobel Prize in Chemistry)
1.
Nucleus positively charged center
2.
Atom mostly empty space
3.
Proton 1840 more massive than e-
Gold Foil Experiment
Ernest Rutherford
(1908 Nobel Prize in Chemistry)
1.
Nucleus positively charged center
2.
Atom mostly empty space
3.
Proton 1840 more massive than e“If the atom is the Houston
Astrodome, then the nucleus is a
marble on the 50-yard line.”
- R. Chang
atomic radius
~100 pm = 1 x 10-10 m
nuclear radius ~0.005 pm = 5 x 10-15 m
Random Interlude:
In 1932, James Chadwick discovered the
neutron, which has no charge at all, by doing
some really complicated experiments that I
don’t understand even a little bit.
Don’t worry though – just remember that
Chadwick discovered it. You’ll be fine!
103
Beryllium Experiment
James Chadwick
1.
Neutron is neutral
2.
Mass p ≈ mass n ≈ 1840 x mass e--
(1935 Nobel Prize in Physics)
a + 9Be
1n
+ 12C + energy
A nucleus contains protons and neutrons
All atomic nuclei are made of two kinds of
particles: protons and neutrons.
Atoms have small, dense positively charged
nuclei.
A proton has a positive charge equal in
magnitude to the negative charge of an
electron.
105
Pssssst…
Atoms are electrically neutral because why??
106
• Because they contain equal numbers of
protons and electrons!
• A neutron is electrically neutral.
107
Masses
A proton has a mass of 1.673 x 10-27 kg.
That is 1836 times greater than the mass of an
electron!
The mass of the neutron is is 1.675 x 10-27 kg
– slightly larger than a proton.
108
Summary of Subatomic Particles
Particle
Symbol
Electron
e
-
Relative Charge
Actual Mass (kg)
-1
9.109 x 10-31
Discoverer
Experiment
JJ Thomson
CRT/oil drop
RA Millikan
Proton
p
+
+1
1.673 x 10-27
Rutherford
Gold foil
Neutron
n
0
0
1.675 x 10-27
Chadwick
Beryllium
Protons determine an atom’s identity
The nuclei of atoms of different elements
differ in their number of protons and therefore
on the amount of positive charge they possess
which influences the pull they will have on
electrons – and then combining with other
elements.
**Physicists have ID’d other subatomic
particles but they have little effect on the
chemical properties of matter.
110
• No two elements have the same number of
protons. The number of protons is unique to
that element alone.
• Charge is determined by the difference of
protons and electrons. Missing an electron?
You have a +1 charged ion.
111
Forces in the nucleus
Nuclear forces are the forces that hold a
nucleus together.
They are the proton-proton, proton-neutron
and neutron-neutron forces. Which is weird
because generally, particles that have the
same electric charge repel each other!
112
However…
• When 2 protons are extremely close to each
other, there is a strong attraction between
them.
• A similar attraction exists when neutrons
are very close to each other or when protons
and neutrons are very close to each other.
113
The radii of atoms are expressed in
picometers
The radius of an atom is the distance from the
center of the nucleus to the outer portion of
the electron cloud.
The electron cloud model is a convenient way
to think of the region occupied by electrons.
114
Electron Cloud Model
115
116
SI Prefixes
Prefix
Symbol
Multipli
er
Tera
T
1012
1 Tm = 1 x 1012 m
giga
G
109
1 Gm = 1 x 109 m
mega
M
106
1 Mm = 1 x 106 m
kilo
k
103
1 km = 1 x 103 m
1m=1m
1 x 10-3 km = 1 m
1m=1m
deci
d
10-1
1 dm = 1 x 10-1 m
10 dm = 1 m
centi
c
10-2
1 cm = 1 x 10-2 m
100 cm = 1 m
milli
m
10-3
1 mm = 1 x 10-3 m
1000 mm = 1 m
micro
µ
10-6
1 µm = 1 x 10-6 m
1 x 106 = 1 m
nano
n
10-9
1 nm = 1 x 10-9 m
1 x 109 = 1 m
pico
pm
10-12
1 pm = 1 x 10-12 m
1 x 1012 = 1 m
The picometer…
Consider that 1 cm is the same fractional part
of 1000 km (about 600 miles) as 100 pm is of
1 cm…
Atomic radii range from about 40 to 270 pm.
118
• Nuclear forces are said to hold protons and
neutrons together. What is it about the
composition of the nucleus that requires the
concept of nuclear forces??
119
Rutherford holds the key
Protons and neutrons are packed very close
together. However, like-charged particles
repel each other, so protons would not be
expected to be close to other protons. The
forces that prevent protons from repelling
each other are the nuclear forces.
120
S.3: Counting Atoms
We can use basic properties of atoms to count the
number of atoms of an element in a sample with
a known mass.
We’ll also become familiar with the mole, a
special unit used by chemists to count atoms and
molecules. (Finally! It’s here!)
121
• All atoms contain the same particles… Yet
all atoms are not the same.
• What makes them different??
122
• Atoms of different elements have different
numbers of protons.
• Atoms of the same element all have the
same number of protons, but not neutrons.
123
• The atomic number (Z) of an element is the
number of protons of each atom of that
element.
• Z = number of protons
124
Look at the periodic table…
An element’s atomic number is usually above
its symbol and the elements are placed in
order of increasing atomic number.
The atomic number identifies an element.
So if the number of protons changes, that
atom becomes a different element.
125
Atomic number
126
The identity of the atom is determined by the
number of protons – not the number of
electrons or neutrons.
The number of electrons & neutrons can vary
and the atom will still be the same element.
But if the number of protons changes, then the
atom becomes an atom of a different element.
127
Isotopes
Isotopes are atoms of the same element that
have different masses.
The isotopes of a particular element all have
the same number of protons and electrons but
different numbers of neutrons.
128
Take hydrogen, for example
All hydrogen atoms have 1 proton, but they
can have different numbers of neutrons – or
no neutrons at all!
Protium is the most common (99.9885% of the H atoms
found on Earth)
Deuterium (0.0115% found on Earth)
Tritium is radioactive and not very common.
129
Most elements…
Most elements consist of mixtures of isotopes.
Tin has 10! (That’s the most of any element)
The atoms in any sample of an element will
most likely be a mixture of several isotopes in
various proportions.
130
Mass number
The mass number (A) of an element is the
total number of protons and neutrons that
make up the nucleus of an isotope.
We have to know the name or atomic number
of the element and the mass of the isotope to
identify the isotope.
131
Mass Number
132
Mass numbers of H isotopes
Protium 1 proton + 0 neutrons
=1
Deuterium 1 proton + 1 neutron = 2
Tritium 1 proton + 2 neutrons
=3
133
The Isotopes of Hydrogen
Designating Isotopes
Hyphen notation = element (hyphen) mass number
hydrogen-3 (tritium)
uranium-235
Nuclear symbol:
mass no. (A)
atomic no. (Z)
235
92
U
135
Atomic number, Mass number and Isotopes
A
X
Z
Mass Number
Atomic Number
Element Symbol
Nuclide symbol
12
C
6
14
C
6
Atomic number (Z)
number of protons in nucleus
Mass number (A)
number of protons + number of neutrons
(Z) + number of neutrons
Isotopes are atoms of the same element (X) with different numbers
of neutrons in their nuclei
To find neutrons:
Mass number – atomic number = number of neutrons
235 (protons + neutrons) -92 protons = 143 neutrons
137
Let’s practice this:
• Write the nuclear symbol for oxygen -16
1) What 3 things do you need?
2) What is the order you write them?
138
Let’s try another one
Tin-121
Write the Nuclide symbol, Atomic No. and
Mass Number.
List the number of protons, neutrons &
electrons found in xenon-113.
139
121
50
Sn
Xenon-113
p+ = 54 e- = 54 n0 = 59
140
How many protons, neutrons, and electrons?
Draw good nuclide symbols.
6 protons, 8 neutrons, 6 electrons
6 protons, 5 neutrons, 6 electrons
Quiz from Tuesday…
While I hand it out – let’s see how you do
with a little target practice…
142
Let’s work on some of these:
Subatomic particle
Sample problem A, pg. 75
Then do the practice questions, pg. 76: 1-3
143
Answers:
1. 35 p, 35 e-, 45 neutrons
1. 13
6 C
3. phosphorus-30
144
Let’s check…
The only stable isotope of naturally occuring
fluorine has a mass number of 19.
How many p+, n0 and e- are in an atom of F?
Write the isotope symbol.
145
19
F
9
146
In chemistry there are many
different concepts of mass…
We need to talk about 3 of them.
147
First, atomic mass…
• Isotopes are atoms of the same element (and
so with the same number of p+ and e-) but
with different masses due to having
different numbers of n0. Right?
148
• The mass of an atom is incredibly small and
it doesn’t make sense to use units like
grams to measure it…
• It’s easier to measure it in atomic mass
units.
149
Scientists like a standard…
• Scientists set up a scale of atomic mass, one
atom was arbitrarily chosen as the standard.
Carbon-12
Chemists use this standard to compare units of
atomic mass. It’s a relative scale… relative
to C-12.
150
Atomic mass unit
It is defined as the mass equal to 1/12th the
mass of one C-12 atom.
Or amu…
(our textbook uses just plain “u”)
151
Atomic mass is the mass of an atom in atomic
mass units (amu)
By definition: 1 atom 12C “weighs” 12 amu
On this scale
1H
= 1.008 amu
16O
= 16.00 amu
Compare that to
average atomic mass
• The average atomic mass of an element is
the weighted average of the masses of its
isotopes on this scale…
• Just like how your grades are weighted.
153
Average atomic mass
This is the number you see on the periodic
table.
154
Average atomic mass
We have to consider that most elements occur
naturally as a mixture of isotopes.
Scientists determine the average mass of a
sample of an element’s isotopes by
determining the percentages of each of the
isotopes and then giving the proper weight to
each value. (Just like your grades!)
155
• Elements rarely occur as only one isotope.
• They exist as mixtures of different isotopes
of various masses.
• This way, the less common isotopes are
accounted for.
156
Don’t confuse mass number with average
atomic mass.
Mass number is the mass of one particular
atom.
Average atomic mass is the average mass of a
group of atoms of the same element that
takes into consideration all the isotopes.
157
• The average atomic mass of an element
depends of both the mass and relative
abundance of each of the element’s
isotopes.
158
The average atomic mass is the weighted
average of all of the naturally occurring
isotopes of the element.
Percent
Isotope Abundance
12C
98.90%
13C
1.10%
Isotopic
Mass, amu
12.00000
13.00335
How do we find
atomic mass?
The average atomic mass of an element
depends of both the mass and relative
abundance of each of the element’s isotopes.
160
Copper consists of 69.15% copper-63, which has an
atomic mass of 62.929 601 amu, and 30.85%
copper-65, which has an atomic mass of 64.927 794
amu. Calculate the average atomic mass for Cu.
Calculating average atomic mass
Step 1: multiply the atomic mass of each
isotope by its relative abundance.
Step 2: add the results.
0.6915 x 62.929 amu + 0.3085 x 64.927 794 amu
= 63.55 amu**
*Our textbook is going to round to 2 decimal places.
162
Average Atomic Mass
Oxygen has three naturally occurring isotopes
in the following proportions:
O-16
99.762 % (15.99491 amu);
O-17
0.038000% (16.99913 amu);
O-18
0.20000% (17.99916 amu).
What is the average atomic mass of oxygen?
Let’s take it to 2 decimal places
1) .99762 x 15.99491 = 15.957
2) .00038 x 16.99914 =
.00646
3) .0020000 x 17.999 =
.03599
15.99 amu
164
Next up…
165
Atoms are very, very small.
For example, 1 atom of H weighs
approximately 1.67 x 10-27 kg.
166
The Mole (mol): A unit to count numbers of particles
Dozen = 12
Pair = 2
Ream of paper
Baker’s dozen
Gross
Fortnight
Googol
Byte
The Mole (mol): A unit to count numbers of particles
Dozen = 12
Pair = 2
Baker’s dozen
Gross
Donkey power
Fortnight
Googol
Mole = 6.0221367 x 1023
Micro World
atoms and
molecules
Macro World
grams
• The answer is that moles give us a
consistent method to convert between
atoms/molecules and grams. It's simply a
convenient unit to use when performing
calculations.
170
The mole (mol) is the amount of a substance that
contains as many particles as there are atoms in
exactly 12.00 grams of 12C
1 mol = 6.0221367 x 1023
Avogadro’s number (NA)
eggs
Molar mass is the mass of 1 mole of shoes in grams
cats
• A mole of carbon atoms is 6.02x1023 carbon
atoms.
• A mole of chemistry teachers is 6.02x1023
chemistry teachers.
172
1 mole of M&Ms…
173
It's a lot easier to write the word
'mole' than to write '6.02x1023’
anytime you want to refer to a large
number of things!
•Basically, that's why this particular unit was
invented.
174
• The abbreviation for mole is mol.
• Molar mass is the weight of one mole of a
chemical compound.
• Molar mass is usually written in units
g/mol.
175
Molar mass
1 mole 12C atoms = 6.022 x 1023 atoms = 12.00 g
1 12C atom = 12.00 amu
For any element
atomic mass (amu) = molar mass (grams)
One Mole of:
S
C
Hg
Cu
Fe
177
The molar mass of He is 4.00 g/mol.
To find how many g of He there are in 2 moles of He,
multiply by the molar mass.
2 .00 mol He x
4.00 g He = 8.00 g He
1 mol He
Conversions with
Avogadro’s Number
• We use Avogadro’s number to find the
number of atoms of an element in moles
• Or
• To find the amount of an element in moles
from the number of atoms.
179
181
Chapter 3
What
Section 3 Counting Atoms
is the mass in grams of 3.50 mol of the
element copper, Cu?
Chapter 3
Sample
The
Section 3 Counting Atoms
Problem B Solution, continued
molar mass of copper from the periodic
table is rounded to 63.55 g/mol.
Chapter 3
Sample
Section 3 Counting Atoms
Problem C
A chemist produced 11.9 g of aluminum, Al.
How many moles of aluminum were
produced?
Chapter 3

Section 3 Counting Atoms
Sample Problem C Solution

Given: 11.9 g Al
Unknown: amount of Al in moles

Solution:
grams Al

´
moles Al
= moles Al
grams Al
The molar mass of aluminum from the periodic
table is rounded to 26.98 g/mol.
1 mol Al
11.9 g Al ´
= 0.441 mol Al
26.98 g Al
Chapter 3
Sample
Section 3 Counting Atoms
Problem D
How many moles of silver, Ag, are in 3.01 
1023 atoms of silver?
Chapter 3




Section 3 Counting Atoms
Sample Problem D Solution
Given: 3.01 × 1023 atoms of Ag
Unknown: amount of Ag in moles
Solution:
moles Ag
Ag atoms ´
= moles Ag
Avogadro's number of Ag atoms
3.01 ´ 10
23
1 mol Ag
Ag atoms ´
=
23
6.022 ´ 10 Ag atoms
0.500 mol Ag
Chapter 3
Sample
Section 3 Counting Atoms
Problem E
What is the mass in grams of 1.20  108 atoms
of copper, Cu?
Chapter 3
Section 3 Counting Atoms
Relating Mass to Numbers of Atoms, continued

Sample Problem E Solution
Given: 1.20 × 108 atoms of Cu

Unknown: mass of Cu in grams

Solution:
moles Cu
grams Cu
Cu atoms ´
´
= grams Cu
Avogadro's number of Cu atoms
moles Cu

The molar mass of copper from the periodic
table is rounded to 63.55 g/mol.
Types of radioactivity
uranium compound
190
Types of radioactivity
The Bohr Model of the Atom
Niels Bohr
(1922 Nobel Prize in Physics)
Thus one mole of ethyl alcohol, C2H6O, weighs 46.069 g. One mole of
water weighs 18.015 g. If we mix 46.069 g of ethyl alcohol with 18.015
g of water, we can be assured that the mixture contains 1 molecule of
ethyl alcohol per molecule of water. Further, we will know that there
are 2 atoms of C and 8 atoms of H per each 2 atoms of O. Thus the
mole allows us to weigh convenient amounts of material containing
known numbers of atoms; i.e., it allows us to count atoms.
193