Download Unit 3 - Structure of the Atom

Unit 3
The Structure of the Atom
Early Theories of Matter
Do you recognize the these names?
Aristotle
or
Democritus
Early Theories of Matter

Atomic Theory


Matter is composed of empty
space through which atomos
(indivisible particle) move
Atoms – small indivisible
particles whose shape
determines the properties of
the matter
460-370 B.C.
4.1 Early Theories of Matter
Aristotle

Elemental Theory


everything in the world was
made up of some combination
of four elements: earth, fire,
water, and air
elements were acted upon by
the two forces of gravity and
levity
gravity was the tendency for
earth and water to sink
 levity the tendency for air and
fire to rise

384-322 B.C.
ARISTOTLE WINS!
Because Aristotle was
considered one of the
greatest thinkers of his
time, Aristotle sets back
chemistry over 2,000
Years!
Early Theories of
Matter




Atoms are small solid
indivisible particles
Different elements are made
of different atoms, but atoms
of the same element are
identical
Atoms combine in simple
whole number ratios to make
compounds
In a reaction, atoms are
rearranged
1766-1844
Early Theories of Matter

How big are atoms?
Reeeeeeeeeeeeeaaaaaaaaaaaaaaaaaaly small!
In the year 2010, it was
estimated that 7 billion
people lived on Earth!
The number of atoms of
copper atoms in ONE
solid copper penny is 50
billion times bigger!
World Population
7 000 000 000
Atoms in a penny
29 000 000 000 000 000 000 000
Can you count to a million this
semester?
1000000 dollars 2 s
1 dollar
1 min 1 hr 1 day
60 s 60 min 24 hr
=
23 days
At 4 hrs a day it would take you
138 days!!!
1 billion would take you 368 years!
How long would it take everyone on
Earth to count the atoms in a
penny?
1,764,167 years
Can We See Atoms?
NO!
While we can’t see them,
we can sense them!
Scanning
Tunneling
Electron
Microscope
(STM)
Scanning Tunneling Electron
Microscope (STM)
Gold atoms
Silicon atoms
Not only can the
probe sense the
atoms electron
cloud, it was
discovered that it
can also manipulate
the atoms!
Scanning Tunneling Electron
Microscope (STM)
“Corrals”
“Round”
Iron atoms on copper
“Stadium”
Iron atoms on copper
Scanning Tunneling Electron
Microscope (STM)
“IBM Scientists Playing”
“Atom” in Japanese
Iron atoms on copper
“IBM”
Xenon on nickel
“carbon monoxide man”
Carbon monoxide on platinum
Subatomic Particles and the
nuclear atom


Thomson (1897) –
Cathode-Ray
Experiments: discovered
the electron .
Established the charge to
mass ratio for these
particles
1856-1940
Cathode Ray Tube
Subatomic Particles and the
nuclear atom

Millikan (1909) – Oil-Drop
Experiment: determined
the mass of the electron.
All of the change in charges came in
multiples of 1.602  10-19!
This must be the actual charge of the electron!
Subatomic Particles and the
nuclear atom
Combined with the charge/mass ratio
from Thomson, he was able to accurately
calculate the mass of
a single electron to be 1/1840 the
mass of a hydrogen nucleus (proton)!
Even with his primitive equipment, his value
was within 1% of the accepted value that
we use today!
Subatomic Particles and the
nuclear atom

Rutherford (1911) –
Gold Foil Experiment:
discovered the nucleus
and basic structure of
the atom.


Electrons are outside
the positively charged
nucleus.
Most of the atom is
empty space.
1871-1937
Gold Foil Experiment
Subatomic Particles and the
nuclear atom


Rutherford (1920) – Concluded there must
be a particle in the nucleus carrying a
charge equal but opposite of the electron
and he called it a proton.
James Chadwick (1932) discovered the existence of a
neutrally charged particle in
the nucleus called the
neutron.
Subatomic Particles and the
nuclear atom
Particle
Symbol
Relative
Location Electrical
Charge
(amu)
Actual
Mass
(g)
Relative
Mass
Electron
e-
Outside
Nucleus
1-
1/1840
9.11×10-28
Proton
p+
Nucleus
1+
1
1.673×10-24
Neutron
n0
Nucleus
0
1
1.673×10-24
How Atoms Differ

Moseley (1913) – discovered each element has a
unique positive charge in their nuclei.



Each element has a different number of protons.
Atomic number - # protons
Atoms are neutral

# protons = # electrons
How many protons
does chlorine have?
17
17
How many electrons?
17
How Atoms Differ


Thomson (1910) – It’s discovered that neon
consists of atoms with 2 different masses.
WHY?
Isotope – atoms with the same number of
protons, but different numbers of neutrons.
Neon-20
Neon-22
p+ = 10
p+ = 10
n0
n0 = 12
= 10
How Atoms Differ
If not all atoms of an element are identical, how can we tell
them apart?



Mass number – the total number of particles in the nucleus
Mass number = # protons + # neutrons
MASS NUMBER IS NOT ON THE PERIODIC TABLE!!!
Mass number
7
Mass number
6
How Atoms Differ
Hyphen Notation
Element name – mass number
Symbol Notation
Mass #
Atomic #
Potassium-41
p+ = 19
41
19
symbol
K
e- = 19
p+ = 19
n0 = 22
e- = 19
n0 = 22
What about ions (atoms with a charge)?
37
17
Cl
p+ = 17
-
e- = 18
n0 = 20
46
20
Ca
2
p+ = 20
e- = 22 18
n0 = 26
How Atoms Differ


Atomic mass – because the mass of atoms
is so small (proton = 1.67×10-24g) we
simplify atomic masses by measuring
them in atomic mass units.
Atomic mass unit (amu) – 1/12 the
mass of carbon-12.


Hydrogen-1 = 1.007825 amu
Silicon-30 = 29.974 amu
How Atoms Differ


If elements have isotopes with different atomic
masses, what is the atomic mass on the periodic
table?
Atomic Mass – weighted average mass of the
isotopes of an element.
This found by summing the mass contribution of
each isotope of the element.
% abundance
)
100
% abundance
 (mass second isotope) (
)
100
average atomic mass  (mass first isotope) (
There are 2 naturally occurring isotopes of
copper – copper-63 and copper-65. If copper-63
has a mass of 62.930 amu and 69.17%
abundance and copper-65 has a mass of 64.928
amu and 30.83% abundance, what is the
average atomic mass of copper?
1. First, calculate the mass contribution of each
isotope to the average atomic mass, being sure
to convert each percent to a fractional
abundance.
For copper-63:
Mass contribution = (62.930 amu)(.6917) = 43.52868
43.53 amu
For copper-65:
Mass contribution = (64.928 amu)(.3083) = 20.017302
20.02 amu
2. Finally, the average atomic mass of the element
is the sum of the mass contributions of each
isotope.
43.53 amu
 20.02 amu
63.55 amu
Unstable Nuclei and
Radioactive Decay
Can chemical reactions change the identity
of an atom?
NO!
Why Not?
You can’t mess with the nucleus!
Does this mean the nucleus is not affected
by any reactions?
NO!
Ch 23 - Unstable Nuclei and
Radioactive Decay
Why do they change?
Stability!
Unstable systems, like atoms with the wrong
number of neutrons or a pencil sitting on
its tip, gain stability by losing energy!
The pencil loses
energy when it falls,
but gains stability
sitting on the table
How is carbon-14
formed?
This can be used to find out how old
something is in carbon-14 dating!
Carbon-14 Dating
How do we see radiation?
Just look around you!
You were thinking of radioactive
particles!
Remember Rutherford’s Alpha
Particles
23.1-3 - Unstable Nuclei and
Radioactive Decay
Types of radioactive decay
Alpha decay ( 42 He )
1.


Mass # drops by 4 and atomic number drops by 2
+2 charge
226
88
Ra 
Rn  He
222
86
4
2
Beta Decay ( 01 e or -10 )
2.


Mass # stays same and atomic number increases by 1
-1 charge
14
6
C N e
14
7
0
-1
23.1 Unstable Nuclei and
Radioactive Decay
3.
Gamma decay ( γ )
0
0




No charge
No change in mass number or atomic number
Usually accompanies alpha or beta decay
Most of the energy lost in radioactive decay is
from gamma decay
238
92
U
Th  He  2 
234
90
4
2
0
0
Detecting Radiation
23.1 Unstable Nuclei and
Radioactive Decay

Balancing nuclear reactions

Mass number and atomic number must be
conserved.
238
92
U
234
90
Th
 2 He
4