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Atomic Structure
Atoms and their structure
Mr. Bruder
Evidence for Atoms



Law of Conservation of Mass
– Mass is not gained or lost in a chemical reaction.
Proposed by Antoine Lavoisier in 1787.
What would happen to the mass reading if the reaction was
done without the balloon (an open system)?
Figure 2.2
2-2
Evidence for Atoms

Law of Definite Proportions
– Proposed by Joseph Proust between
1797 and 1804
– A compound always has the same
relative amounts of the elements that
compose it.
– For example, when water is broken
down by electrolysis into oxygen and
hydrogen, the mass ratio is always 8
to 1.
Figure 1.2
2-3
Dalton’s Atomic Theory

John Dalton (1766-1844) had four theories
1.
All elements are composed of submicroscopic indivisible particles
called atoms
Atoms of the same element are identical. The atoms of anyone
element are different from those of any other element
Atoms of different elements can physically mix together or can
chemically combine w/ one another in simple whole-number ratios
to form compounds
Chemical reactions occur when atoms are separated, joined, or
rearranged. However, atoms of one element are never changed into
atoms of another elements as a result of a chemical reaction
2.
3.
4.
Atoms & Subatomic Particles

Atom- smallest particle of an element that
retains the properties of that element
Proof of Atoms: STM Image of
Gold

The scanning tunneling
microscope, invented in
1981, allows us to
create images of matter
at the atomic level.
Figure 2.4
2-6
Electron
J.J Thomson (1856-1940) – discovered the
electron in 1897
 Electron is the negative charged subatomic
particle
 An electron carries exactly one unit of
negative charge & its mass is 1/1840 the
mass of a hydrogen atom

Cathode Ray

The Cathode Ray tubes pass electricity through
a gas that is contained at a very low pressure
Thomson’s Experiment
Voltage source
-
+
Thomson’s Experiment
Voltage source
-
+
Thomson’s Experiment
Voltage source
-
+
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

+
Passing an electric current makes a beam
appear to move from the negative to the
positive end
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

+
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
Thomson’s Experiment
Voltage source
+
 By adding an electric field
Thomson’s Experiment
Voltage source
+
 By adding an electric field
Thomson’s Experiment
Voltage source
+
 By adding an electric field
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 Atomic Model

Thomson’s Atomic Model

Thomson though electrons were like plums
embedded in a positively charged “pudding”, so
his model was called the “plum pudding” model
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

Mass of Electron
In 1909 Robert Millikan determined the mass of an
electron with his Oil Drop Experiment
 He determined the mass to be 9.109 x 10-31 kg
 The oil drop apparatus

Millikan’s Experiment
Atomizer
+
-
Oil
Microscope
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
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-
+
+
Some drops would hover
Millikan’s Experiment
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-
+
+
+
+
+
+ +
+
Millikan’s Experiment
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+
+
Measure the drop and find volume from 4/3πr3
Find mass from M = D x V
Millikan’s Experiment
-
-
+
+
From the mass of the drop and the charge on
the plates, he calculated the charge on an electron
Proton
In 1886 Goldstein discovered the Proton
 Proton is a positively charged subatomic
particle found in the nucleus of a atom

Radioactivity
Discovered by accident
 Bequerel
 Three types
– alpha- helium nucleus (+2 charge, large
mass)
– beta- high speed electron
– gamma- high energy light

Ernest Rutherford
Rutherford (1871-1937) proposed that all mass
and all positive charges are in a small
concentrated region at the center of the atom
 He used the Gold-Foil Experiment to prove his
theory
 In 1911 he discovered the Nucleus
 Nucleus- central core of an atom, composed of
protons and neutrons
 The nucleus is a positively charged region and it
is surrounded by electrons which occupy most of
the volume of the atom

Rutherford’s Experiment
Used uranium to produce alpha particles
 Aimed alpha particles at gold foil by
drilling hole in lead block
 Since the mass is evenly distributed in
gold atoms alpha particles should go
straight through.
 Used gold foil because it could be made
atoms thin

Lead
block
Uranium
Florescent
Screen
Gold Foil
What he expected
Because
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

+
+
Nuclear Atom Viewed in Cross
Section
Copyright © Cengage Learning. All rights reserved 46
The Neutron



Because the protons in the atom could account for
only about half the mass of most atoms, scientists
knew there was another heavy particle in the
nucleus.
Neutrons were proposed by Ernest Rutherford in
1907 (to account for a mass discrepancy in the
nucleus) and discovered in 1932 by James
Chadwick.
The neutron has about the same mass as a proton
but with no charge.
2-47
Bohr Model


1.
2.
Bohr changed the Rutherford model and
explained how the electrons travel.
Bohr explained the following in his model:
Electrons travel in definite orbits around the
nucleus
His model was patterned after the motion of the
planets around the sun. It is often called the
Planetary model.
48
Bohr’s Model
Nucleus
Electron
Orbit
Energy Levels
The Quantum Mechanical Model
Schrodinger



Things that are very small behave
differently from things big enough
to see.
The quantum mechanical model is
a mathematical solution
It is not like anything you can see.
Modern View
The atom is mostly
empty space
 Two regions
 Nucleus- protons and
neutrons
 Electron cloud- region
where you have a
chance of finding an
electron

Neutron
James Chadwick (1891-1974) – discovered
the neutron in 1932
 Neutron is a subatomic particle with no
charge but their mass is nearly equal to that
of a proton

Quark
Protons & Neutrons can still be broken down into
a smaller particle called the Quark
 The Quark is held together by Gluons

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

Subatomic Particles
Protons and neutrons are located at the center of
an atom called the nucleus.
Electrons are dispersed around the nucleus.
EOS
Chapter 2: Atoms, Molecules, and Ions
55
Atomic Particles
Particle
Charge
Mass (kg)
Location
Electron
-1
9.109 x 10-31
Electron
cloud
Proton
+1
1.673 x 10-27
Nucleus
Neutron
0
1.675 x 10-27
Nucleus
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
Symbols

Contain the symbol of the element, the mass
number and the atomic number
Mass
number
Atomic
number
X
Sub-atomic Particles
Z - atomic number = number of protons
determines type of atom
 A - mass number = number of protons +
neutrons
 Number of protons = number of electrons if
neutral

Symbols
A
X
Z
23
Na
11
Atomic Structure Symbols
Proton = p+
 Electron = e Neutron = n0

Atomic # - Subscript
 Mass # - Superscript

235
92
U
Rules for Atomic Structure
1.
2.
3.


Atomic # = # of Protons
# of Protons = # of Electrons
Mass # = # of Protons + # of Neutrons
# of Neutrons = Mass # - # of Protons
If you know the Mass # & Atomic # you
know the composition of the element
Symbols
 Find
the
–number of protons
–number of neutrons
–number of electrons
–Atomic number
–Mass Number
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
Symbols
 if
an element has 78 electrons and
117 neutrons what is the
–Atomic number
–Mass number
–number of protons
–Complete symbol
Example
Element Atomic Mass # Protons Electro Neutro
K
#
19
ns
19
11
16
5
17
46
35
ns
19
23
35
Isotopes
Isotope- atoms that have the same number
of protons but different number of neutrons
 Since isotopes have a different number of
neutrons the isotope has a different mass
number.
 Isotopes are still chemically alike because
they have the same number of protons and
electrons

Two Isotopes of Sodium
Copyright © Cengage Learning. All rights reserved 68
Examples of Isotopes
Heavy Water


One ice cube is made
with water that
contains only the
hydrogen-2 isotope.
The other ice cube is
composed of water
with normal water
which contains mostly
hydrogen-1.
Which is which?
Figure 2.13
2-70
Isotopes


Isotopes are atoms of the same element with different
masses.
Isotopes have different numbers of neutrons.
11
C
6
12
C
6
13
C
6
14
C
6
© 2009, Prentice-Hall, Inc.
Naming Isotopes
Put the mass number after the name of the
element
 carbon- 12
 carbon -14
 uranium-235

Electrical Charges
Electrical charges are carried by particles of
matter
 Atoms have no net electrical charges
 Given the number of negative charges
combines with the number of positive
charges = Electrically Neutral
 All elements are Electrically Neutral

Atomic Mass vs. Atomic Weight
Atomic Mass is for a single element
 Most elements are Isotopes
 How do we find their mass?
 We use Atomic Weight

Atomic Masses
An atomic mass unit (amu) is defined as
exactly one-twelfth the mass of a carbon-12
atom
1 u = 1.66054 × 10–24 g
The atomic mass of an element is the
weighted average of the masses of the
naturally occurring isotopes of that element
EOS
Chapter 2: Atoms, Molecules, and Ions
75
Measuring Atomic Mass





Unit is the Atomic Mass Unit (amu)
One twelfth the mass of a carbon-12 atom
Each isotope has its own atomic mass. We need
the average from the percent abundance
Each isotope of an element has fixed mass and a
natural % abundance
You need both of these values to find the Atomic
Weight
Calculating Atomic Weight



1.
2.
3.

Cl-35 34.969amu and 75.77% abundance
Cl-37 36.966amu and 24.23% abundance
To solve for Cl-35
AMU x Abundance
34.969 x .7577
= 26.496
You solve for Cl-37
Atomic Weight Cont.

1.
2.
3.




Cl-37
AMU x Abundance
36.966 x .2423
= 8.957
Now you combine your two answers
26.496 + 8.957=
35.453
Look at Cl on the table. What is the Atomic
Weight?
Example

Calculate the atomic weight of copper.
Copper has two isotopes. One has 69.1%
and has a mass of 62.93 amu. The other has
a mass of 64.93 amu. What is the atomic
weight???
Atomic Weight & Decimals
Atomic Weight- of an element is a
weighted average mass of the atoms in a
naturally occurring sample of an element
 Atomic Weights use decimal points because
it is an average of an element
