Atomic Structure and Nuclear
Chemistry
Chapter 4 and 18
Elements a.k.a atoms

Robert Boyle first defined an element
as a substance which could no longer be
broken down into other substances
 Each element has unique properties
 Many early scientists speculated how the
element or atom was structured
 Theory of atomic structure evolved from
early thoughts to today’s atom
What do you know about the atom?
Take a moment and create a concept web
about the atom.
 Work on your own.
 You have about 5 minutes.
 Jot down everything you can connect to
atoms.

Atoms
Timeline of the Atom
Democritus (460-370 B.C.) understood that if you cut a stone in two pieces, each piece
contained the same material as the original stone. He also believed that you could do
this an infinite number of times. He called these infinitesimally small pieces of matter
atomos, meaning "indivisible.“
Rutherford’s gold foil
experiment leads to
atomic nucleus and in 1919
the introduction of the
proton
1808 Dalton first
proposed a
theory on
atoms
Discovery
of the
electron by
J.J.
Thompson
in the late
1890’s
1910 Lord
Kelvin’s
“plum
pudding
theory
1912 –Bohr
model of the
hydrogen atom
Mid-1920’s –
wave
mechanical
model
1932 –
Rutherford and
his coworker
Chadwick
identified the
neutron
Who started…
John Dalton
(1766 – 1844)
was an English
scientist who
made his living
as a teacher in
Manchester.
Dalton’s Atomic Theory (p.88)


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

Elements are composed of atoms
All atoms of a given element are identical
Atoms of a given element are different from those of any other
element
Atoms of one element combine with atoms of other elements to
form compounds

Law of Constant Composition: all samples of a
compound have the same proportion of the elements as
in any other sample of that compound
Atoms are indivisible in a chemical process.
 all atoms present at the beginning of a chemical process must
also be present at the end of the process.
 atoms are not created or destroyed, they must be conserved.
 atoms of one element cannot be turned into atoms of another
element
Atomic Structure History

Discovery of the Electron
 1st atomic particle identified
 In 1897, J.J. Thomson used a cathode ray tube to
deduce the presence of a negatively charged particle.
 Cathode ray tubes pass electricity through a gas that
is contained at a very low pressure. This creates a
beam of negatively charged particles bent by an
electric field.
Conclusions from the Study of the
Electron
 Cathode
rays have identical properties regardless of the
elemental gas used to produce them.
 Therefore, all elements must contain identically
charged particles (electrons).
 Atoms are neutral, so there must be positive particles in
the atom to balance the negative charge of the electrons
 Electrons have so little mass that atoms must contain
other particles that account for most of the mass
 An electron is a tiny, negatively charged particle
 The
next step…. What is the positive charge?
Models of the atom
William Thomson’s (Lord Kelvin’s)
Atomic Model
Lord Kelvin believed that the electrons
were like plums embedded in a
positively charged “pudding,” thus it
was called the “plum pudding” model
(easier to think of as “chocolate chips"
in chocolate chip cookie dough.)
Rutherford’s Gold Foil Experiment
Rutherford’s Gold Foil Experiment
Rutherford shot α (alpha) particles at a thin
sheet of gold foil (think: bullet = alpha
particles, target atoms = gold foil)
 α particles are positively charged
 gold atoms are about 50 larger than α
particles.
 Particles were fired at a thin sheet of gold
foil
 Particles hit on the detecting screen (film)
were recorded

(a) The results that the metal foil experiment would have
yielded if the plum pudding model had been correct.
(b) Actual results known as Rutherford’s model.
Over 98% of the  particles
went straight through
About 2% of the  particles
went through but were
deflected by large angles
About 0.01% of the 
particles bounced off the
gold foil
Most of the volume of the
atom is empty space
Rutherford’s Conclusion: A Nuclear Model
The atom contains a tiny dense center called the
nucleus
The nucleus is essentially the entire mass of the
atom (extremely dense)
The nucleus is positively charged
The electrons move around in the empty space of
the atom surrounding the nucleus
Finally, the neutron..




Discovered in 1932 by Chadwick based on the idea from
Rutherford
Has no charge
Is located in the nucleus
Mass a mass of 1 amu (actually, it’s slightly larger than a
proton but for our work the mass is the same)
The Modern Atom



We know atoms are composed of three main
atomic particles - protons, neutrons and electrons
The nucleus contains protons and neutrons
The radius of the atom is about 100,000 times larger than
the radius of the nucleus
Summary of Atomic Particles
Particle
Charge
Mass #
Location
Electron
-1
0
Electron cloud
Proton
+1
1
Nucleus
0
1
Nucleus
Neutron
Going beyond the electron, proton, and neutron
Describing an Atom
How many protons, neutron, and electrons does an atom have?
Atomic Structure - protons
The number of protons in an atom of a given element is
the same as its atomic number (Z).

(Z) is found on the Periodic Table, whole # for each
element
# of
protons
Atomic #
(Z)
6
6
Phosphorus
15
15
Gold
79
79
Element
Carbon
Atomic Structure - neutrons

Mass number = protons + neutrons; always a whole
number.

# of Neutrons = mass number - # of protons

Atomic mass to mass number– decimal number in
each element’s box on the periodic table. If you
round the atomic mass of an element to the closest
whole number you get the mass # for that element.
Atomic Structure - Electrons


# of Electrons = # of protons if the atom is neutral
If the chemical symbol is written with a charge,
representing an ion, the charge indicates the number of
electrons that have been added or removed from the
atom.




If the ion has a positive charge (cation), subtract that charge
from the # of protons to get the number of electrons.
If the ion has a negative charge (anion), add that charge number
to # of protons to get the number of electrons.
# of Electrons = # protons – charge
Charge = # protons - # electrons
Representing atomic particles in atoms


The number of each type of atomic particle (proton,
neutron, electron) is determined by using symbols.
There are several different ways to write an element:



Atomic symbols
Nuclear symbols
Atomic Symbols include the element symbol and a charge
if any.



C – neutral carbon
C+4 – carbon cation
C-4 – carbon anion
Nuclear Symbols
Mass number
(p+ + no)
Charge (if any)
238
92
Atomic number
(number of p+)
U
+
Element symbol
Element name
Cl-1 –38
Mass number
Fluorine-18
Element symbol with charge
Mass number
Atom
Silver – 109
Pb-208
C-14
# protons
# neutrons
# electrons
Isotopes
atoms of an element with the same number of
protons but different numbers of neutrons
Isotope
Protons
Electrons
Neutrons
Hydrogen–1
(protium)
1
1
0
1
1
1
1
1
2
Hydrogen-2
(deuterium)
Hydrogen-3
(tritium)
Nucleus
Two isotopes of sodium.
Isotopes Examples
 3517Cl

17Cl
37
H-1 H-2 H-3
Copper – 63 Copper – 65
Isotopes and Atomic Mass
AMU




When we think about the mass of an atom, we use
atomic mass units (amu).
A proton is 1 amu
A neutron is 1 amu
Add up protons and neutrons to get the mass number
(not atomic mass) Why?


Most elements in nature have isotopes
All these isotopes contribute to the average atomic mass
(listed on the table)
Determining Average Atomic Mass
The average atomic mass seen on the periodic table is a
combination of all an element’s isotopes and their
abundance.
 To determine the average atomic mass for an element, you
must
1. Multiply the percentage (percent abundance) of each
isotope of the element by its mass number.
2. Add the products of the multiplications together.
3. Divide by 100.
4. Your answer should be very close to the atomic mass of
the element for that element

Average Atomic Mass Examples

Find the average atomic mass of each of the following
elements from their percentages and mass numbers.

69.17% 63Cu and 30.83% 65Cu

5.85% Fe-54, 91.75% Fe-56, 2.12% Fe-57 and 0.28% Fe-58
Nuclear Reactions

What you just did was write nuclear reactions.


Typical reactions are decay reactions and capture reactions.
What did you notice about the products of the reactions?




The products of the reactions are isotopes of the element.
Nuclear reactions produce different particles that are not elements.
You need these particles to balance out the protons and neutrons in
the nuclei.
In a nuclear reaction, the atomic number (Z) and the mass number
(A) are conserved.
Radioactive decay

Radioactive decay is a natural process.



A nuclei may spontaneously kick out a particle, forming a new
element.
There are four common particles:





Only a handful (~200) of the known isotope nuclei (2000) do not
decay
The alpha particle - α (gold foil)
The beta particle – β
The gamma particle – γ
The positron particle
In many cases, the process does not stop at one step, but
rather a combination of steps. This is known as a decay series.
Capture Reactions
Nuclear Transformations

Nuclear transformation is the changing of one element to
another (modern alchemy!!!) using larger atoms

Rutherford observed 1st transformation in 1919

Marie Curie and her husband another transformation (14
years later)

Electron capture is the “natural” nuclear transformation

Man-made elements are made by bombarding two nuclei
together
The nuclear particles
Radioactivity Review

Radioactivity is expressed as either a decay process or a
capture process

Decay processes can be connected as a chain

Nuclear transformations use larger atoms to create
different elements than what you started with
Review of nuclear symbols


Mass number is in upper left of symbol
Atomic mass is in lower left of symbol
238
92

U
Nuclear particles are written in the nuclear format
The alpha particle (α)



The alpha particle is actually a helium nucleus.
It is the weakest of the decay particles.
A alpha particle has a mass number of 4 and an atomic
number of 2.



If an α particle is added to a nuclei, the mass number will increase
by 4 and the atomic number by 2.
If an α particle is released from a nuclei, the mass number will
decrease by 4 and the atomic number by 2.
For example:
The Beta particle (β)

The beta particle is an electron.


It has no mass.
It does have a charge

Examples include:

Sometimes a nucleus will grab an electron that is close. This is
called electron capture.
The Gamma Particle (γ)





This is also know as the gamma ray.
This is a high energy photon of light.
Picked up by specific detectors.
It is the strongest (most dangerous) of the decay particles.
Example:
The Positron


The positron is similar to a beta particle, but has a
positive charge.
Example:
Neutron Emission


Neutron emission or decay does not change the element,
only the mass
Example:
Balancing Nuclear Equations

Mass # and the atomic # totals must be the same for
reactants and the products.
 3919K
 3517Cl + ___
 20682Pb
 0-1e + ___
Writing Balanced Nuclear Equations

Alpha decay of Cu-68

Gamma emission of Thorium-235

Positron emission of P-18

Astatine-210 releasing 3 neutrons

Electron capture with Ti-45
Nuclear Chemistry and half-life
Balancing Nuclear Equations
Half-life
Half-Life and Nuclear Stability

Radioactive isotopes or nuclides all decay because they are
unstable, some just breakdown much faster than others

Geiger counter or scintillation counters are used to detect
particles.

Half-life – amount of time for half of the original sample to
decay

For two samples of the same isotope, regardless of the
sample size, after one half-life, only half of the original amount
of sample remains.
Sample Half-lives

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Isotopes
Carbon – 14
Sodium – 24
Bismuth – 212
Polonium – 215
Thorium – 230
Thorium – 234
Uranium – 235
Uranium – 238
Half-Live
5730 years
15 hours
60.5 seconds
0.0018 seconds
75400 years
24.1 days
7.0 x 108 years
4.46 x 109 years
Working with half-life

A material has t1/2 = 10 minutes. If you begin with 16g,
how long will it take to decay to 2 g?





Begin with 16 g, 1 half life gets you to 8 g.
2 half lives get you to 4 g.
3 half lives get you to 2 g.
So, 3 x 10 minutes = 30 minutes.
A material has t1/2 = 150 years. If you begin with 100 g,
how long will it take to decay to 3.125 g?

A material has t1/2 = 15 minutes. How much material is
left after 75 minutes if you begin with 100 g?


Calculate the number of half-lives used = 75/15 = 5
Run through 5 half lives:






After 1 – 50g
After 2 – 25 g
After 3 – 12.5 g
After 4 – 6.25 g
After 5 – 3.125 g
A material has t1/2 = 6.2 years. How much material is left
after 24.8 years if you begin with 14 g?

What is the half-life of a material that decays from 16 g to
2 g over 20 minutes?

Determine the number of half lives:






16 → 8
8→4
4→2
3 half lives spent
Divide the time by the number of half-lives: 20/3 = 6.67
minutes
What is the half-life of a material that decays from 125 g
to 3.9 g over 100 years?
Uses of nuclear chemistry
Uses of Nuclear Chemistry

Medicine




X-rays and MRIs
Chemical tracers
Energy
Destruction
Fission versus Fusion

Fusion – combining two smaller nuclei into one heavier,
more stable nucleus.
3 He + 1 H  4 He + 0 e
2
1
2
1

Fission – splitting a large unstable nucleus into two nuclei
with smaller mass numbers.
209 Po  125 Te + 84 Ge
84
52
32
Sample fission reaction
Sample fusion reaction