The Development of Modern
Atomic Theory
In the beginning…
• Democritus- 400 BC, he proposed that the
universe was made of tiny indivisible units.
• He called these atoms, from the Greek word
atomos, meaning unable to be divided or cut.
• He did not have any evidence to support his
theory, but nonetheless, some people
believed.
Aristotle
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-
-
350 B.C.
Aristotle modified a theory of matter and
stated that matter on earth consisted of:
Air
Fire
Earth
Water
“celestial element”
Dalton’s Atomic Theory
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•
•
•
In 1808, John Dalton proposed a revised atomic theory.
According to Dalton:
1) All matter is composed of atoms.
2) Atoms are identical in properties such as mass and
size.
• 3) Atoms cannot be created, destroyed or subdivided.
• 4) Atoms can combine in fixed proportions to form
compounds.
• 5) In chemical reactions, atoms are combined, separated
or rearranged.
Dalton’s Atomic Theory
Why is his theory noteworthy?
-It is the first theory based on experimental evidence…
-He took the laws developed by others and combined them
to form his atomic theory.
The Law of Conservation of Mass
Definition: The mass of reactants must equal the mass of
products in a chemical reaction. Therefore, mass cannot
be created nor destroyed in a chemical reaction.
What does this mean????
Practice problems
1) In a chemical reaction of hydrogen and oxygen gas,
water is made. If you begin with 3.4 g of hydrogen and
1.8 g of oxygen, how much water can be made, using
the law of conservation of mass?
1) If 22.5 g of sodium hydroxide solution is added to 123.5
g of copper (II) nitrate, how much products would be
made?
The Law of Definite Proportions
Defined: When 2 or more atoms combine to form a
compound, they will combine in the same fixed whole
number ratio no matter where they are found or how much
you have of the compound
Example:
The Law of Multiple Proportions
Defined: When two elements can form a series of
compounds, the ratio of the masses of the elements can be
reduced to simple whole numbers.
What did Dalton get wrong?
If you carefully consider all parts of Dalton’s
theory, he got a lot of it right!
However, there are a few things that we now
know are not correct:
1) Atoms can be subdivided.
2) Atoms of an element are not all identical.
Thomson’s Model of the Atom
• 1897- Experiments with electricity led to the
discovery that atoms were NOT indivisible!
• He conducted a famous cathode ray tube
experiment which suggested that these rays
were made of negatively charged particles
that came from inside of atoms.
Cathode Ray Tube Experiment
• Thomson hooked up a
battery to a cathode
metal plate and an
anode metal plate. The
glass tube had a gas at
low pressure, a beam of
light traveled in the tube
from cathode to anode.
Thomson’s Experiments with Cathode
Rays
• To be sure that the
beam was not coming
from only the gas or the
metals, he swapped
them out. The result
was the same each time.
• Then, he put a magnet
to the outside of the
tube. He noticed that
the beam was either
deflected or attracted to
the magnet.
Thomson’s Experiments with Cathode
Rays
• In the paddle wheel
experiment, Thomson
proved the beam had
enough mass to move an
object!
• Paddle Wheel Video
Conclusions from Thomson’s Cathode
Ray Experiments
1) Beam remained regardless of gas and metal… must be
in everything.
2) Beam has a negative charge.
3) Beam has enough mass to move another object.
Electrons are discovered!!!! They are small, have mass,
are in every part of matter… and are negatively
charged!
Thomson’s Plum Pudding Model
• Based on the outcome of his experiment, he
proposed a new model of the atom.
• He reasoned that the electrons were spread
throughout the atom, just like in plum
pudding, a dessert in his time.
• We might use the chocolate chip cookie
model to better visualize it today.
Millikan and Electron Charge
1910- Millikan conducted an oil drop
experiment that allowed him to determine the
exact charge on one electron (1.6 x 10-19C)
Oil Drop Experiment Video
Rutherford’s Model of the Atom
• Ernest Rutherford worked for JJ Thomson.
• He proposed his model of the atom shortly after
Thomson in 1907.
• Radioactivity had just been discovered, and it opened a
whole new world of experiments to researchers.
• Rutherford was fascinated by the emission of alpha
particles and how they were deflected as beams of light
from zinc sulfide.
Gold Foil Experiment
• He decided to aim a beam of positively charged alpha
particles at a thin sheet of gold foil. (0.00004 cm)
• His hypothesis:
• Because of Thomson’s model, Rutherford predicted that
there would not be a large enough mass of positive
charge to cause many particles to bounce back. He
thought they would travel in a straight path through the
foil.
Gold Foil Experiment
• But, the observations did not match his hypothesis.
Most passed straight through, but a few were deflected
by a large amount.
• A few even bounced back!
• He said, “It was quite the most incredible event that has
ever happened to me in my life. It was almost as
incredible as if you fired a 15- inch shell at a piece of
tissue paper and it came back and hit you.”
Gold Foil Experiment
Results of Rutherford’s experiment
• An atom’s positive
charge in concentrated
in the center of the
atom, called the
nucleus.
• Most of the density of
the atom is in the
nucleus, as well as
most of its mass.
• Here is his atomic
model:
What about the neutron?
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•
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James Chadwick discovered the neutron in
1932. He had worked with Rutherford in his
lab from 1911-1913 and then with Geiger.
He bombarded beryllium with alpha particles
and observed a particle coming off with the
same mass as a proton, but no charge.
The neutron is the most difficult particle to
detect. Why?
The Structure of the Atom
An atom consists of a(n):
-nucleus (made of protons and neutrons)
-electron cloud
Current atomic model:
Wave mechanical model (courtesy
of Edwin Schroedinger’s electron probability equation
Atomic Composition
●
Protons (p+)
●
● + electrical charge
●
relative mass = round to 1 a.m.u (atomic mass unit)
Electrons (e-)
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●
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negative electrical charge
relative mass = round to 0 a.m.u
Neutrons (no)
●
●
no electrical charge
mass = round to 1 a.m.u
Atomic Number
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•
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The atomic number (symbolized by Z) is the
number of protons in the nucleus of the
atoms of that element.
The atomic number= # of protons
The number of protons tells us the identity
of that element. If the atomic number
changes, the element does too!
Mass Number
• The mass number of an element is the rounded atomic
mass (decimal number) on the periodic table.
• Example: H = 1.008 amu, so it’s
mass number=1
• The mass number of elements is based on a standard
mass: Carbon-12 has a mass of 12 amu (atomic mass
units).
• Mass number = # of protons + # of neutrons
Mass Number
The number of neutrons= mass number- atomic number
Nuclide
Mass
Number
Protons
Neutrons
Electrons
Oxygen-16
16
8
8
8
Carbon-12
6
6
Uranium-235
92
92
Copper-65
29
29
Atoms are electrically
NEUTRAL….except
●
Neutral Atoms
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Protons = Electrons
Examples: Ca, I, O
Charged Atoms (called IONS)
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Protons do NOT equal Electrons
Since electrons are negative: if an atom has
more electrons than protons, the charge on
the ion is negative
Examples: Ca+2, I-1, O-2
IONS
The charge of an atom can be found by looking at protons
and electrons.
More Electrons(Anions) = Negative Charge
Less Electrons(Cations) = Positive Charge
Protons ALWAYS EQUAL the atomic number! This CAN
NOT change!!!
Ions
Protons Electrons Charge/Symbol
Magnesium 12
10
Fluorine
9
10
Lithium
3
2
Isotopes
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•
Atoms of an element with different mass
numbers and a different number of
neutrons.
Isotope notation:
Carbon-14
or
14
6
C
Isotope Examples
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For the following isotopes, calculate the
number of protons, neutrons and electrons.
Iodine-131
Neon-21
Carbon-14
Isotopes of Hydrogen
• Each element has a different number of isotopes. You
can look up the number of isotopes...they have already
been determined by scientists.
Isotope
Protons
Electrons
Neutrons
Hydrogen-1
Protium
1
1
0
Hydrogen-2
Deuterium
1
1
1
Hydrogen-3
Tritium
1
1
2
Nucleus
Why are the Atomic masses on the
periodic table decimals?
• Any sample of any element is made up of a
mixture of all the isotopes for that element
based upon their fractional abundances
(percentages found in nature).
• Each element has a unique number of isotopes.
• To determine the average atomic mass, you
find the weighted average by using masses
given and their percent abundances.
Example
A sample of copper has 2 isotopes: Copper-63 has a
percentage of 69.15% and Copper-65 has an abundance of
30.85%. Calculate the average atomic mass.
Step 1: Divide each percentage by 100 to get a decimal.
Step 2: Multiply each decimal by its mass.
Step 3: Add all the answers from step 2 to get the average
atomic mass.
69.15 x 63 amu=43.56
30.85 x 65
amu=20.05
100
100
43.56 + 20.05 = 63.61 amu
Example
• Calculate the average atomic mass of neon. There are
3 isotopes of neon:
• Neon-20 has an exact mass of 19.992440 amu and a
percentage of 90.48%
• Neon-21 has an exact mass of 20.993847 amu and a
percentage of 0.27%
• Neon-22 has an exact mass of 21.993366 amu and a
percentage of 9.25%
How do we know about the isotopes
and percentages of each element?
• An instrument called the mass spectrometer
measures the mass to charge ratio and traveling speed
to the detector.
• The detector counts the particles that pass (calculating
a percentage) and the mass.
• A graph is then shown on the computer.
Interpreting mass spectrometry data
This instrument can be used to identify
element based on mass.
What elements are shown to the left?
Abundances can also be determined!
Counting Atoms
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•
•
Atoms are small… like on the order of
picometers.
Remember 1 m = 1012pm
The radius of atoms are represented in
picometers.
Counting Atoms
• Since atoms are so small, chemists like to use a larger
quantity of atoms so that they can be seen!
• That’s where the mole comes in.
• Avogadro’s number is 6.022 x 1023 atoms, molecules or
particles. This is how many items make up 1 mole of a
substance.
• It can be used as a conversion factor:
6.022 x 1023atoms
or
1 mole
1 mole
6.022 x 1023atoms
Using Avogadro’s Number
Examples:
1) How many moles are in 3.45 x 1024 atoms of
zinc?
2) Calculate the number of molecules in a
sample of 3.45 moles of water.
3) How many moles make up 5.90 x 1025 alpha
particles?
Counting atoms
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•
•
Molar mass- the number of grams in 1
mole of a substance
The molar mass is the same number as the
atomic mass on the periodic table.
It can also be used as a conversion factor.
Carbon 12.011 g
1 mole
Finding molar mass of compounds
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When finding molar mass of compounds,
you multiply the atomic mass on the periodic
table for each element by its subscript.
If no subscript is shown, it is understood to
be 1.
Add all the masses together!!
Examples:
KMnO4
Cu2SO3
Using molar mass to find
grams/moles
Set up a conversion factor… the molar mass in
grams always equals 1 mole of the compound.
Examples:
Find the number of moles of H2O2 in 63.50 g.
63.50 g x 1 mole = 1.867 moles
34.02 g
The MOLE Map