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Unit 4 – Atomic Structure
Bravo – 15,000 kilotons
Atoms, Ions, and the Periodic Table
What is an atom?
It is smallest particle of an element that retains
the elements properties.
But how did we come to know all the information
we have about these tiny particle?
Democritus (460-370 BC)
Democritus (460-370 BC)
•
•
•
•
Matter is made of tiny, solid, indivisible particles which he called
atoms (from atomos, the Greek word for indivisible).
Different kinds of atoms have different sizes and shapes.
Different properties of matter are due to the differences in size,
shape, and movement of atoms.
Democritus’ ideas, though correct, were widely rejected by his
peers, most notably Aristotle (384-322 BC). Aristotle was a very
influential Greek philosopher who had a different view of matter.
He believed that everything was composed of the four elements
earth, air, fire, and water. Because at that time in history,
Democritus’ ideas about the atom could not be tested
experimentally, the opinions of well-known Aristotle won out.
Democritus’ ideas were not revived until John Dalton developed his
atomic theory in the 19th century!
John Dalton (1766-1844)
John Dalton (1766-1844)
–
–
–
–
–
All matter is composed of extremely small
particles called atoms.
All atoms of one element are identical.
Atoms of a given element are different from
those of any other element.
Atoms of one element combine with atoms of
another element to form compounds.
Atoms are indivisible. In addition, they
cannot be created or destroyed, just
rearranged.
• Dalton’s theory was of critical importance. He was
able to support his ideas through experimentation,
and his work revolutionized scientists’ concept of
matter and its smallest building block, the atom.
• Dalton’s theory has two flaws:
– In point #2, this is not completely true. Isotopes
of a given element are not totally identical; they
differ in the number of neutrons. Scientists did
not at this time know about isotopes.
– In point #5, atoms are not indivisible. Atoms are
made of even smaller particles (protons, neutrons,
electrons). Atoms can be broken down, but only in
a nuclear reaction, which Dalton was unfamiliar
with.
Discovery of the Electron
JJ Thomson (1856-1940)
Discovery of the Electron
JJ Thomson (1856-1940)
• Discovered the electron, and determined that it had a negative
charge, by experimentation with cathode ray tubes. A cathode ray
tube is a glass tube in which electrons flow due to opposing charges
at each end. Televisions and computer monitors contain cathode ray
tubes.
• Thomson developed a model of the atom called the plum pudding
model. It showed evenly distributed negative electrons in a uniform
positive cage.
• Diagram of plum pudding model:
Discovery of the Nucleus
Ernest Rutherford (1871-1937)
http://www.mhhe.com/physsci/chemistry/animations/chang_2e/rut
herfords_experiment.swf
Discovery of the Nucleus
Ernest Rutherford (1871-1937)
• Discovered the nucleus of the atom in his famous Gold Foil
Experiment.
• Alpha particles (helium nuclei) produced from the radioactive
decay of polonium streamed toward a sheet of gold foil. To
Rutherford’s great surprise, some of the alpha particles
bounced off of the gold foil. This meant that they were
hitting a dense, relatively large object, which Rutherford
called the nucleus.
Rutherford then discovered the proton, and next, working with a colleague,
James Chadwick (1891-1974), he discovered the neutron as well.
Models of the Atom - Niehls Bohr
Models of the Atom - Niehls Bohr
• Developed the Bohr model of the atom (1913) in
which electrons are restricted to specific
energies and follow paths called orbits a fixed
distance from the nucleus. This is similar to
the way the planets orbit the sun. However,
electrons do not have neat orbits like the
planets.
• Diagram of Bohr model:
Quantum Mechanical Model
Quantum Mechanical Model
• This is the current model of the atom. We now know
that electrons exist in regions of space around the
nucleus, but their paths cannot be predicted. The
electron’s motion is random and we can only talk about
the probability of an electron being in a certain region.
Modern Atomic Theory
 All matter is composed of atoms
 Atoms cannot be subdivided, created, or
destroyed in ordinary chemical reactions.
However, these changes CAN occur in nuclear
reactions!
 Atoms of an element have a characteristic
average mass which is unique to that
element.
 Atoms of any one element differ in
properties from atoms of another element
Discovery of the Electron
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.
Conclusions from the Study of
the Electron
 Cathode rays have identical properties regardless
of the element used to produce them. All elements
must contain identically charged 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
Thomson’s Atomic Model
Thomson believed that the electrons were like plums
embedded in a positively charged “pudding,” thus it was
called the “plum pudding” model.
Rutherford’s Gold Foil Experiment
 Alpha particles are helium nuclei
 Particles were fired at a thin sheet of gold foil
 Particle hits on the detecting screen (film) are
recorded
Rutherford’s Findings
 Most of the particles passed right through
 A few particles were deflected
 VERY FEW were greatly deflected
“Like howitzer shells bouncing off
of tissue paper!”
Conclusions:
 The nucleus is small
 The nucleus is dense
 The nucleus is positively charged
Atomic Particles
Particle
Charge
Mass #
Location
Electron
-1
0
Electron cloud
Proton
+1
1
Nucleus
0
1
Nucleus
Neutron
The Atomic
Scale
 Most of the mass of the
atom is in the nucleus
(protons and neutrons)
 Electrons are found
outside of the nucleus (the
electron cloud)
 Most of the volume of
the atom is empty space
“q” is a particle called a “quark”
About Quarks…
Protons and neutrons are
NOT fundamental particles.
Protons are made of
two “up” quarks and
one “down” quark.
Neutrons are made of
one “up” quark and
two “down” quarks.
Quarks are held together
by “gluons”
Atomic Number
Atomic number (Z) of an element is the
number of protons in the nucleus of each atom
of that element.
Element
# of protons
Atomic # (Z)
6
6
Phosphorus
15
15
Gold
79
79
Carbon
Mass Number
Mass number is the number of protons and
neutrons in the nucleus of an isotope.
Mass # = p+ + n0
Nuclide
p+
n0
e-
Mass #
Oxygen - 18
8
10
8
18
Arsenic - 75
33
42
33
75
Phosphorus - 31
15
16
15
31
Isotopes
Isotopes are atoms of the same element having
different masses due to varying numbers of neutrons.
Isotope
Protons
Electrons
Neutrons
Hydrogen–1
(protium)
1
1
0
Hydrogen-2
(deuterium)
1
1
1
Hydrogen-3
(tritium)
1
1
2
Nucleus
Atomic Masses
Atomic mass is the average of all the naturally
isotopes of that element.
Carbon = 12.011
Isotope
Symbol
Composition of
the nucleus
% in nature
Carbon-12
12C
6 protons
6 neutrons
98.89%
Carbon-13
13C
6 protons
7 neutrons
1.11%
Carbon-14
14C
6 protons
8 neutrons
<0.01%
Isotopes
•
Isotopes are atoms of an element with the same number of protons
but different numbers of neutrons.
• Change in # of n0 = Change in the Mass #
•
Most elements on the periodic table have more than one naturally
occurring isotope.
•
There are a couple of ways to represent the different isotopes.
One way is to put the mass after the name or symbol: Carbon-12 or
C-12
•
Another way is :
mass # 12
atomic # 6
C
Determining Average Atomic Mass
• The atomic mass on the periodic table is
determined using a weighted average of all the
isotopes of that atom.
• In
–
–
–
order to determine the average atomic mass,
convert the percent abundance to a decimal
multiply it by the mass of that isotope
values for all the isotopes are added to together to get
the average atomic mass
Formula:
(Mass1 * Abundance1) + (Mass2 * Abundance2) + (Mass3 * Abundance3) + etc….
Example of Average atomic mass
calculation
Given:
12C = 98.89% at 12 amu
13C = 1.11% at 13.0034 amu
Calculation:
(98.89%*12 amu) + (1.11%*13.0034 amu) =
(0.9889*12 amu) + (0.011*13.0034 amu) =
12.01 amu
Now you try one:
• Neon has 3 isotopes: Neon-20 has a mass of 19.992
amu and an abundance of 90.51%. Neon-21 has a
mass of 20.994 amu and an abundance of
0.27%. Neon-22 has a mass of 21.991 amu and an
abundance of 9.22%. What is the average atomic
mass of neon?
• The answer is: (0.9051 * 19.992 amu) + (0.0027 *
20.994 amu) + (0.0922 * 21.991 amu) =
20.179 amu
Now compare this mass for Neon to the mass on the
periodic table!
The Mole
1 dozen = 12
1 gross = 144
1 ream = 500
1 mole = 6.02 x 1023
There are exactly 12 grams of
carbon-12 in one mole of carbon-12.
Avogadro’s Number
6.02 x 1023 is called “Avogadro’s Number” in
honor of the Italian chemist Amadeo Avogadro
(1776-1855).
I didn’t discover it. Its
just named after me!
Amadeo Avogadro
MOLE
MOLAR MASS
Molecular Mass
Atomic Mass (really amu’s)
Formula Mass
6.02 x 1023 particles
Molecules
Atoms
Formula Units
Calculations with Moles:
Converting moles to grams
How many grams of lithium are in 3.50 moles of
lithium?
3.50 mol Li
6.94 g Li
1 mol Li
=
24.29 g Li
Calculations with Moles:
Converting grams to moles
How many moles of lithium are in 18.2 grams of
lithium?
18.2 g Li
1 mol Li
6.94 g Li
=
2.62
mol Li
Calculations with Moles:
Using Avogadro’s Number
How many atoms of lithium are in 3.50 moles of
lithium?
3.50 mol Li 6.02 x 1023 atoms Li
1 mol Li
= 2.11 x 1024 atoms Li
Calculations with Moles:
Using Avogadro’s Number
How many atoms of lithium are in 18.2 g of
lithium?
18.2 g Li 1 mol Li
6.94 g Li
6.02 x 1023 atoms Li
1 mol Li
(18.2)(6.022 x 1023)/6.94
= 1.58 x 1024 atoms Li
Nuclear Symbols
Mass number
(p+ + no)
235
92
U
Atomic number
(number of p+)
Element symbol
Types of Radioactive Decay
alpha production (a): helium nucleus
238
4
234
92 U  2 He  90Th
0
beta production (b):  1 e
234
234
90Th  91Pa

0
1e
4
2+
He
2
Alpha
Radiation
Limited to
VERY large
nucleii.
Beta
Radiation
Converts a
neutron into
a proton.
Types of Radioactive Decay
gamma ray production (g):
238
4
U

92
2 He

234
90Th
 2 00 g
0
positron production 1 e :
22
0
Na

11
1e

22
10 Ne
electron capture: (inner-orbital electron
is captured by the nucleus)
201
0
201
Hg

e

80
1
79 Au
 00 g
Types of Radiation
Deflection of Decay Particles
attract
Opposite charges_________
each other.
repel
Like charges_________
each other.
Nuclear
Stability
Decay will occur in
such a way as to
return a nucleus to
the band (line) of
stability.
Half-life Concept
Sample Half-Lives
A radioactive nucleus reaches a stable state
by a series of steps
A
Decay
Series
Nuclear Fission and Fusion
•Fusion: Combining two light nuclei to form
a heavier, more stable nucleus.
3
1
4
0
2 He  1H  2 He  1e
•Fission: Splitting a heavy nucleus into two
nuclei with smaller mass numbers.
1
235
142
91
1
0 n  92 U  56 Ba  36 Kr  30 n
Energy and Mass
Nuclear changes occur with small but measurable
losses of mass. The lost mass is called the mass
defect, and is converted to energy according to
Einstein’s equation:
DE = Dmc2
Dm = mass defect
DE = change in energy
c = speed of light
Because c2 is so large, even small amounts of
mass are converted to enormous amount of
energy.
Fission
Fission Processes
A self-sustaining fission process is called
a chain reaction.
Neutrons
Causing
Event
Fission
subcritical
<1
critical
=1
supercritical
>1
Result
reaction stops
sustained reaction
violent explosion
A Fission Reactor
Fusion