Chapter 3 Atomic
Structure
3-1 Early Models of the Atom
3-2 Discovering Atomic Structure
3-3 Modern Atomic Theory
3-4 Changes in the Nucleus
3-1 Early Models of the Atom
What are atoms?
What are the postulates of Dalton’s atomic theory?
Ancient Greek – 450BC
 Proposed that all matter is composed of
tiny, invisible particles called atoms
 No one believed him during his lifetime

◦ Including Aristotle

His beliefs were not accepted until the
17th and 18th centuries
Democritus

Was not accepted until 2 discoveries were
made
◦ Lavoisier’s law of conservation of matter
◦ Joseph Louis Proust’s law of constant
composition
 A compound will always contain the same
proportions by mass of elements
◦ Water will always have 88.9% oxygen (O) and 11.1%
hydrogen (H)
Acceptance
English school teacher
 Studied past theories of atoms and laws
of matter
 Formed an atomic theory of matter

John Dalton (1766 – 1844)
Ea element is composed of extremely
small particles called atoms
 All atoms of a given element are identical,
but they differ from those of any other
element
 Atoms are neither created/destroyed in
any chem rxn
 A given compound always has the same
relative #s and kinds of atoms

Dalton’s Atomic Theory of
Matter
The smallest particle of an element that
retains the chemical identity of that
element
 There are 118 elements wh means there
are 118 different kinds of atoms.

Atoms

Atoms are like the words in these slides.
◦ If we broke it all apart, separated and organized
the letters, you would find only 26 piles.
◦ But by taking letters from different piles we can
create millions of very different words

Just like words can be separated into letters,
matter can be separated into atoms.
◦ These separated atoms are called elements

Think of all the words you could make with
the letters A, D, and M….
Atoms


Produces images of atoms
Created in 1981
◦ Nickel
Platinum
Scanning Tunneling Microscope
(STM)

Consumer Tip
◦ “100 Percent Natural”
Chemistry In Action (p93)

Macroscopic – looking a the whole picture
◦ A tree
 It is made of the leaves, branches, trunk, roots

Microscopic – the more detailed vision of
an object and what makes it function
◦ A leaf off a tree and the little veins that carry
the nutrients through it
Macroscopic vs Microscopic
Macroscopic vs Microscopic

Chemists make their
observations in the
macroscopic world
◦ It is the world in wh we
all live

In order to understand
that world, the goal is
to understand the
atoms that the world is
made of
◦ Discoveries/Possibilities
b/c of the study of atoms
 Deciphering the genetic
code
 Designing plastics
 Understanding the hole in
the ozone
 Imprinting data on silicon
chips
3-2 Discovering Atomic Structure
How is atomic structure related to electricity?
What did cathode rays indicate about atoms?
What did Rutherford conclude from his alpha-scattering experiment?


Scientists couldn’t
figure out why atoms
of one element acted
differently than
another element’s
atoms
Michael Faraday
(1791-1867) said
that the structure of
an atom was directly
related to electricity
Electric Charges

Atoms contain
particles that have
electrical charges
Electric Charges


An object will either
have a positive or
negative charge
2 like charges will
repel
◦ Positive w/ positive
◦ Negative w/ negative

2 opposite charges
will attract
◦ Positive w/ negative

Franklin didn’t know
where these charges
came from
Benjamin Franklin

Electric current - A moving stream of
electrical charges
◦ Electricity from wall socket or battery
Studying electrical currents provide keys
to understanding electrical charges
 Mid-1800s, began studying electric
currents in glass tubes w/ little air

Cathode Rays and Electrons

Tube attached on ea end to a battery
◦ Positive and negative
 Negative = cathode
 Positive = anode

Radiation travels from cathode to anode
◦ b/c radiation came from cathode end, called
cathode ray and the tube a cathode ray tube
Cathode Rays and Electrons

Cathode Ray tube being effected by a
magnet
◦ http://www.youtube.com/watch?v=7YHwMWcx
eX8&feature=related
-
+
Battery
Cathode Rays and Electrons

Negative particles within the atom
◦ JJ Thompson (1856-1940)

Mass of 9.11 x 10
-28
gram
◦ 0.000000000000000000000000000911 gram
◦ Robert Millikan (1868-1953)
Electrons

Henry Becquerel (18521908)
◦ Placed uranium on photo
paper and an image
appeared
◦ Uranium was emitting
radiation
Radioactivity:
spontaneous emission
of radiation from an
element
 Marie Curie and
husband Pierre
discovered the
elements of radium and
polonium were also
radioactive

Radioactivity



Thompson said there were electrons in the
atom (neg charge)
Why is the atom neutral then?
Rutherford’s Gold Foil Experiment
◦ http://www.youtube.com/watch?v=5pZj0u_XMbc
◦ Called this center the nucleus
 Has a positive charge
 Very small
◦ If the atom was the size of a football stadium, the nucleus
would be smaller than a dime sitting in the middle
 Electrons would be smaller than Franklin Roosevelt’s eye on
the dime
The Nuclear Atom
The Nuclear
Atom
http://www.ndted.org/EducationResources/Hi
ghSchool/Magnetism/reviewat
om.htm
How far are the
electrons from
the nucleus?
If the earth was the nucleus,
the electrons would cover an
area as large as the distance
b/w the earth and nearest
stars
3-3 Modern Atomic Theory
What are the names and properties of the 3 subatomic particles?
How can you determine the # of protons, neutrons, and electrons in an atom/ion?
What is an isotope? What is atomic mass?

We know atoms are made from protons,
neutrons, and electrons
◦ Recently scientists have found even smaller
particles
 Quarks, Gluons, Mesons, Muons, and others
◦ They don’t seem to impact any Chemistry so
chemists ignore
 Physicists study them
Subatomic Particles

Nucleus
◦ Contains the protons and neutrons
 Protons = positive – p+
◦ Have the same but opposite charge as electrons
 Neutrons = neutral/no charge – n0

Electrons
◦ Negatively charged – e◦ Move in the space outside nucleus – e- cloud
◦ Very small compared to p+
 2000 e- = 1 p+
The Structure of the Atom

Mass
◦ Too small for normal measurements
◦ Has own unit - atomic mass unit (amu)
◦ P+ and n0 = 1 amu, e- = 0 amu b/c so small

Length
◦ Diameter = 0.100 – 0.500 nanometer
 Nanometer = nm = 10-9 meter
◦ If you drew a line across a penny (1.9 cm), you
would touch 810 million copper atoms
◦ If you lined up all 810 million nuclei, you would only
have a line 4 x 10-6 meter long
 4 millionths of a meter
Size of Subatomic Particles

Henry Moseley (1887-1915)
◦ Student of Rutherford
◦ Discovered atoms of ea element contained
differ positive charges
Lead to the idea that an atom’s identity
comes from the # of p+ in nucleus
 Call this # atomic number

Atomic Numbers
The # of protons
 Ea element has a unique atomic #
 Can tell an element’s atomic # from
periodic table

Atomic Number
The p+ are positive
 The e- are negative
 The atom is neutral
 This means, the p+ must equal the e
◦ For N, atomic # = 7
 Means p+ = 7
◦ Means e- = 7
Neutral Atom

How many protons and electrons in:
◦ Oxygen (O)
 8 p+ and e-
◦ Magnesium (Mg)
 12 p+ and e-
◦ Silicon (Si)
 14 p+ and e-

What element has 11 protons?
◦ Sodium
Examples



When an atom gains/loses e-, it will have a
charge
When an atom has a charge, called ion
Charge of ion = #p+ - #e◦ If a magnesium atom loses 2 e-, ionic form has a charge
of:
 #p+ - #e- = 12 – 10 = +2
◦ It is important to add the plus (+) sign into the answer
 Also possible to have a negative (-)
◦ Some people write the charge with the +/- after the # (2+)

After you have calculated the charge, to write it
with the element symbol, add it as a subscript
◦ For our magnesium example: Mg+2
Ions

Write the chemical symbol for the ion w/:
◦ 9 p+ and 10 e F-
◦ 13 p+ and 10 e Al+3
◦ 7 p+ and 10 e N-3

How many p+ and e- are present in:
◦ S-2 ion
 16 p+ and 18 e-
◦ Li+ ion
 3 p+ and 2 e-

Write the chemical symbol for the ion w/:
◦ 12 p+ and 10 e Mg+2
◦ 74 p+ and 68 e W+6
Examples
All atoms of the same element, have the
same # of p+
 They may not have the same # of n0
 If atoms have the same # of p+ but
different # of n0 , we call them isotopes
 Most elements have at least 1 isotope

◦ 1 usually more frequent than another
In nature, it is usually a mixture
 To tell isotopes apart, we use the mass #

Isotopes

Mass # = #p+ + #n0
◦ An atom w/ 17p+ and 18n0 would have an
mass # of 35
 Mass # = 17 + 18 = 35
◦ b/c 17 p+, tells us it is a chlorine atom
 Chlorine – 35

A way to write the element symbol w/
atomic and mass #s would be:
mass #
37
atomic #
17
Mass Number
Cl
element symbol

How many protons, neutrons, and electrons are in the following
ions?
◦ 56
Fe+2
26
 26 p+, 24 e-, and 30 n0
◦
27
13
Al+3
 13 p+, 10 e-, and 14 n0
◦
79
34
Se-2
 34 p+, 36 e-, and 45 n0

Write the complete chemical symbol for the ion w/
◦ 21 p+, 24 n0, and 18 e
45
Sc+3
21
◦ 53 p+, 74 n0, and 54 e 127 I53
Examples
The average mass of all the isotopes of an
element
 Listed in the periodic table

Atomic Mass
Practice Problems
# 1-30
3-4 Changes in the Nucleus
What changes accompany nuclear reactions?
What is radioactivity?
Change the composition of an atom’s
nucleus
 Produces alpha, beta, or gamma radiation

◦ Alpha and beta radiation comes from radiation
emitted from the nucleus
Nuclear Reactions

Almost all atoms have stable nuclei
◦ Not radioactive
Radioactivity could have harmful effects –
good its rare to find in nature
 Why are some more stable than others?

◦ # of p+ and n0 in the nucleus
◦ Some combinations cause instability
Nuclear Stability
In nucleus, p+ and n0 are packed together
in a very small space
 How do p+ stay together in the small
space if like charges repel?

◦ Held there by strong nuclear force
 Can only be found in this situation
◦ Neutrons act like a net to hold the p+ in along
with the strong nuclear force
Nuclear Stability

Pattern of stability
◦ Atomic # 1-20 – nuclei stable, = # of p+ and
n0
◦ Beyond 20 p+ - more n0 needed to keep stable
◦ Atomic # above 83 – radioactive nuclei
 No # of n0 will make it stable

Atoms unstable if too many or too few
neutrons
◦ Atoms w/ too many emit beta radiation
Nuclear Stability

Alpha (α)
◦
◦
◦
◦
◦
Alpha particles have 2 p+ and 2 n0
Identical to Helium – 4 nucleus
Travel only a few cm
Easily stopped by paper or clothing
Usually doesn’t pose a health threat unless
actually enters the body
4
2
He
+2
4
2
He
4
2
α
Types of Radioactive Decay

Beta (β)
◦ High speed electrons (not the ones around the
nucleus)
◦ Comes from charges inside a nucleus
◦ A neutron changes into a p+ and e p+ stays in nucleus
 e- (beta particle) is propelled out of nucleus at
high speed
◦ 100 times more penetrating than alpha
 Able to penetrate 1-2 mm of solid material
 Able to pass through clothing and damage skin
0
-1
e
-
0
-1
e
Types of Radioactive Decay
0
-1
β

Gamma (γ)
◦
◦
◦
◦
Very energetic form of light our eyes can’t see
Doesn’t have any particles
More penetrating than others
Able to penetrate deep into solid material
 Body tissue
◦ Stopped only by heavy shielding
 Concrete or lead
0
0
γ
Types of Radioactive Decay

When an atom emits radiation, it undergoes
radioactive decay
◦ Called decay b/c nucleus is decomposing to form a new
nucleus

The best way to understand the decay is w/ a
nuclear equation
226
88
Ra
222
86
Rn
+
4
2
α
Types of Radioactive Decay

Look at Figure 3-30 on p115 and answer
the questions
◦ Would this protective suit protect the worker
from alpha radiation?
◦ Why would a person working w/ alpha radiation
also need to be concerned w/ gamma
radiation?
◦ Would protective clothing such as this stop
gamma radiation from penetrating the worker’s
skin?
Partner Activity
131
53
I

131
54
Xe
+
Beta decay equation
0
-1
β

Alpha decay of

Alpha decay of

Beta decay of



Alpha decay of
Alpha decay of
Beta decay of
185
79
238
92
24
11
242
94
231
91
233
87
Au
U
Na
Pu
Pa
Fr
Practice Problems

Multiple Choice
◦ all

True/False
◦ all

Concept Mastery
◦ (20-22, 25)

Critical Thinking
and Problem
Solving
◦ 29, 31-33
Chapter 3 Review