Download The History of the Atom

THE HISTORY OF THE ATOM
How did we learn about the atom?
WARM UP-IF YOU DID NOT FINISH THE
SCIENTISTS WORKSHEET FINISH NOW!
1 .Chemistry is the
study of matter,
brainstorm the
different types of
matter and list them
in your warm-up.
2. How do you compare
Physical and chemical
properties
THE ATOMIC THEORY OF MATTER

All matter is made up of fundamental particles.

What does “fundamental” mean?
S.MORRIS 2006
HISTORY OF THE ATOM
460 BC
Democritus develops the idea of atoms
he pounded up materials in his pestle and
mortar until he had reduced them to smaller
and smaller particles which he called
ATOMA
(greek for indivisible)
THE HISTORY OF THE ATOM
How did we learn about the atom?
Warm- up
• Chemistry is the study
of matter, brainstorm
the different types of
matter and list them in
your warm-up.
• How do you compare
Physical and chemical
properties
THE ATOMIC THEORY OF MATTER

All matter is made up of fundamental particles.

What does “fundamental” mean?
THE GREEK PHILOSOPHERS, 400 B.C.
 Democritus
coined the
term, “atom”
 Derived
from the
Greek word atomos,
meaning “indivisible.”
 Are
atoms indivisible?
THE 1ST SCIENTIFIC THEORY, 1803

Dalton’s Atomic Theory
1.
All matter is made up of atoms.
2.
Atoms are indivisible and
indestructible.
3.
A) Atoms of the same element are
identical.
B) Atoms of one element are
different from atoms of another
element.
4.
5.
Still True Today?
1. Yes
2. No: subatomic
particles
3. A) No: isotopes
B) Yes: differ by #
protons
Atoms can physically mix together or
can chemically combine in wholenumber ratios to form compounds.
4. Yes
Chemical reactions occur when
atoms are separated, joined, or
rearranged.
5. Yes
1ST DISCOVERY OF SUBATOMIC PARTICLES,
1897

J. J. Thomson

Atoms are divisible and can be
broken down into subatomic
particles.

Discovered the electron and
is accredited with discovery
of the proton.
Plum Pudding Model of the
Atom: Tiny particles of negative
charge embedded in a ball of
positive charge.
HOW DID THOMSON DISCOVER THE
ELECTRON?

Passed an electric current through gases in a
cathode-ray tube producing a glowing beam.

If electrically charged plates are placed near the cathode tube,
the cathode ray will be attracted towards the positive plate and
repelled from the negative plate.
Therefore, the ray must be made of negatively-charged
particles.

HOW DID THOMSON KNOW THAT
THERE WERE PROTONS?
 Thomson
inferred the existence of positively
charged particles since matter is neutral!

You do not get a shock every time you touch
matter!
Ball of Positive
Charge
+
Negative e= Neutral atom
MILLIKAN’S OIL DROP
EXPERIMENT, 1916
 Found
the mass and charge of the electron.
Force
upwards from
the
negatively
charged
plate.
Force of Gravity
HERE’S WHAT MILLIKAN’S
APPARATUS REALLY LOOKED LIKE…
CHADWICK, 1932


Discovered the neutron.
Isotopes: Atoms of an element that are
chemically alike but differ in mass.

Same number of protons and electrons, different
number of neutrons!

Ex. Carbon-14 and Carbon-12
SUBATOMIC PARTICLES
Particle
Symbol
Relative
Charge
Relative
Mass
Actual
Mass (g)
Proton
p+
1+
1
1.67 x 10-24
Neutron
no
0
1
1.67 x 10-24
Electron
e-
1-
1/1840
9.11 x 10-28
Which particle(s) make up most of the mass of the atom?
OKAY, SO NOW WE HAVE
SUBATOMIC PARTICLES…
How are the particles arranged in the
atom???
RUTHERFORD’S GOLD FOIL
EXPERIMENT, 1911
HERE’S WHAT YOU NEED TO
REMEMBER:
Rutherford shot alpha particles (2+ charge)
at gold foil.

1)
2)
Most of the alpha particles
went straight through…
=> Atoms are mostly empty space.
A few bounced back…
=> Small, dense positively-charged nucleus.
BOHR’S PLANETARY MODEL, 1913

Electrons orbit the nucleus like planets orbit the sun.
BOHR: ELECTRONS IN ORBITS
•
Bohr studied how atoms emit or absorb light.
Negative electrons circle a positive nucleus.
•
Electrons overcome the attraction of the nucleus
because they are in motion.
Given more energy, electrons will move to an
orbit farther away from the nucleus.
Electrons can occupy only certain orbits
called energy levels.
Energy Levels: regions of space in which ecan move about the nucleus of an atom.
ELECTRON CLOUD MODEL - SCHRODINGER,
1926
a.k.a. Quantum Mechanical Model,
Wave Model
WHY IS IT CALLED THE ELECTRON
CLOUD MODEL?




Electrons do not orbit.
Electrons can only have certain energies (energy
levels).
Cannot pinpoint the exact location of electrons (only
probability). Therefore, the probability of finding an
electron is represented by an electron cloud.
Heisenberg’s Uncertainty Principle: It is impossible
to know the velocity and the position of a particle at
the same time.
THE GREEK PHILOSOPHERS, 400 B.C.
 Democritus
coined the
term, “atom”
 Derived
from the
Greek word atomos,
meaning “indivisible.”
 Are
atoms indivisible?
THE 1ST SCIENTIFIC THEORY, 1803

Dalton’s Atomic Theory
1.
All matter is made up of atoms.
2.
Atoms are indivisible and
indestructible.
3.
A) Atoms of the same element are
identical.
B) Atoms of one element are
different from atoms of another
element.
4.
5.
Still True Today?
1. Yes
2. No: subatomic
particles
3. A) No: isotopes
B) Yes: differ by #
protons
Atoms can physically mix together or
can chemically combine in wholenumber ratios to form compounds.
4. Yes
Chemical reactions occur when
atoms are separated, joined, or
rearranged.
5. Yes
1ST DISCOVERY OF SUBATOMIC PARTICLES,
1897

J. J. Thomson

Atoms are divisible and can be
broken down into subatomic
particles.

Discovered the electron and
is accredited with discovery
of the proton.
Plum Pudding Model of the
Atom: Tiny particles of negative
charge embedded in a ball of
positive charge.
HOW DID THOMSON DISCOVER THE
ELECTRON?

Passed an electric current through gases in a cathode-ray
tube producing a glowing beam.

https://www.youtube.com/watch?v=2xKZRpAsWL8

If electrically charged plates are placed near the cathode tube,
the cathode ray will be attracted towards the positive plate and
repelled from the negative plate.
Therefore, the ray must be made of negatively-charged
particles.

HOW DID THOMSON KNOW THAT
THERE WERE PROTONS?

Thomson inferred the existence of positively charged
particles since matter is neutral!

You do not get a shock every time you touch
matter!
Ball of Positive
Charge
+
Negative e= Neutral atom
MILLIKAN’S OIL DROP
EXPERIMENT, 1916

Found the mass and charge of the electron.
Force
upwards from
the
negatively
charged
plate.
Force of Gravity
HERE’S WHAT MILLIKAN’S
APPARATUS REALLY LOOKED LIKE…
CHADWICK, 1932


Discovered the neutron.
Isotopes: Atoms of an element that are
chemically alike but differ in mass.

Same number of protons and electrons, different
number of neutrons!

Ex. Carbon-14 and Carbon-12
SUBATOMIC PARTICLES
Particle
Symbol
Relative
Charge
Relative
Mass
Actual
Mass (g)
Proton
p+
1+
1
1.67 x 10-24
Neutron
no
0
1
1.67 x 10-24
Electron
e-
1-
1/1840
9.11 x 10-28
Which particle(s) make up most of the mass of the atom?
RUTHERFORD’S GOLD FOIL
EXPERIMENT, 1911
HERE’S WHAT YOU NEED TO
REMEMBER:
Rutherford shot alpha particles (2+ charge)
at gold foil.

1)
2)
Most of the alpha particles
went straight through…
=> Atoms are mostly empty space.
A few bounced back…
=> Small, dense positively-charged nucleus.
BOHR’S PLANETARY MODEL, 1913

Electrons orbit the nucleus like planets orbit the sun.
BOHR: ELECTRONS IN ORBITS
•
Bohr studied how atoms emit or absorb light.
Negative electrons circle a positive nucleus.
•
Electrons overcome the attraction of the nucleus
because they are in motion.
Given more energy, electrons will move to an
orbit farther away from the nucleus.
Electrons can occupy only certain orbits
called energy levels.
Energy Levels: regions of space in which ecan move about the nucleus of an atom.

https://www.youtube.com/watch?v=pNTMYNj2Ul
k
ELECTRON CLOUD MODEL - SCHRODINGER,
1926
a.k.a. Quantum Mechanical Model,
Wave Model
WHY IS IT CALLED THE ELECTRON
CLOUD MODEL?




Electrons do not orbit.
Electrons can only have certain energies (energy
levels).
Cannot pinpoint the exact location of electrons
(only probability). Therefore, the probability of
finding an electron is represented by an electron
cloud.
Heisenberg’s Uncertainty Principle: It is
impossible to know the velocity and the position of
a particle at the same time.
Bohr’s Atom
electrons in orbits
nucleus
HELIUM ATOM
Shell
proton
+
-
N
N
+
electron
What do these particles consist of?
-
neutron
OKAY, SO NOW WE HAVE
SUBATOMIC PARTICLES…
How are the particles arranged in the
atom???
ATOMIC STRUCTURE
I. Subatomic Particles
SUBATOMIC PARTICLES
ATOM
NUCLEUS
ELECTRONS
PROTONS
NEUTRONS
POSITIVE
CHARGE
NEUTRAL
CHARGE
NEGATIVE CHARGE
equal in a
Atomic
Most Number
of the atom’s mass.
neutral atom
equals the # of...
QUARKS
ATOMIC STRUCTURE
II. Masses of Atoms
o
Mass Number
o
Isotopes
o
Relative Atomic Mass
o
Average Atomic Mass
MASS NUMBER
mass
# = protons + neutrons
 always a whole
number
 NOT on the
Periodic Table!
© Addison-Wesley Publishing Company,
Inc.
ATOMIC NUMBER

# of protons
Every element has a different number of protons
 In a neutral atom, # protons = # electrons

Example:
9
Be
4
ISOTOPES

Atoms of the same element with different mass
numbers.
 Nuclear symbol:
Mass #
Atomic
#
 Hyphen notation: carbon-12
12
6
C
ISOTOPES
© Addison-Wesley Publishing Company, Inc.
ISOTOPES

Chlorine-35

atomic #:

mass #:

# of protons:

# of electrons:

# of neutrons:
17
35
17
17
18
35
17
Cl
ISOTOPES

Chlorine-37

atomic #:

mass #:

# of protons:

# of electrons:

# of neutrons:
17
37
17
17
20
37
17
Cl
Warm Up for Wednesday
Complete the Pre-assessment and turn to bin
• Name the three
subatomic particles and
give there charges:
• Fill in the Chart:
Element
Protons
Electrons
Neutrons
Protons
electrons
neutrons
Carbon
Arsenic
Argon
Uranium
•
•
•
•
What are Isotopes?
Compare C-12 to C-14
What are Ions?
Compare Cl to Cl-1
C-12
C-14
Cl
Cl-1
Average Atomic Mass
• Weighted average of the relative
abundance and all the isotopes
Recall: Atomic Mass Unit
• The unit that describes the mass of an
atom
• Symbol: amu
Average Atomic Mass
• Relative abundance and isotopes averaged
• on the Periodic Table the entire number
• round to 2 decimal places on yours
Avg.
Atomic
Mass
(mass)(%) (mass)(%)

100
Average Atomic Mass
• EX: Calculate the avg. atomic mass of oxygen if its
abundance in nature is 99.76% 16O, 0.04% 17O, and
0.20% 18O.
Avg.
(16)(99.76 )  (17)(0.04)  (18)(0.20)

Atomic 
100
Mass
16.00
amu
Example #1
• The mass of a Cu-63 atom is 62.94 amu, and
that of a Cu-65 atom is 64.93 amu. The
percent abundance of Cu-63 is 69.17% and
the percent abundance of Cu-65 is 30.83%.
What is the average atomic mass of Cu?
Example #1
• The mass of a Cu-63 atom is 62.94 amu, and that of a
Cu-65 atom is 64.93 amu. The percent abundance of
Cu-63 is 69.17% and the percent abundance of Cu-65 is
30.83%. What is the average atomic mass?
• Step 1: Find the contribution of each isotope:
Cu-63: (62.94 amu) x (0.6917) = 43.535598 amu
Cu-65: (64.93 amu) x (0.3083) = 20.017919 amu
• Step 2: Add the relative abundances from each isotope
together.
43.535598 amu + 20.017919 amu = 63.553517 amu
Round answer to two numbers after the decimal:
63.55 amu
Example #2
• Calculate the average atomic mass of
chromium. It is made up of isotopes with
the following percent compositions and
atomic masses: 83.79% with a mass of 51.94
amu; 9.50% with a mass of 52.94 amu;
4.35% with a mass of 49.95 amu; 2.36 %
with a mass of 53.94 amu.
• Calculate the average atomic mass of chromium. It is made up of
isotopes with the following percent compositions and atomic masses:
83.79% with a mass of 51.94 amu; 9.50% with a mass of 52.94 amu;
4.35% with a mass of 49.95 amu; 2.36 % with a mass of 53.94 amu.
• Step 1: Find the contribution of each isotope:
(51.94 amu) x (0.8379) = 43.520526 amu
(52.94 amu) x (0.0950) = 5.0293 amu
(49.95 amu) x (0.0435) = 2.172825 amu
(53.94 amu) x (0.0236) = 1.272984 amu
• Step 2: Add the relative abundances from each isotope together.
43.520526 amu + 5.0293 amu + 2.172825 amu + 1.272984 amu =
51.995635 amu
Round answer to two numbers after the decimal: 52.00 amu
Example #3
• Calculate the average atomic mass of iron if its
abundance in nature is 15% iron-55 and 85%
iron-56.
Example #3
• Calculate the average atomic mass of iron if its
abundance in nature is 15% iron-55 and 85% iron-56.
• Step 1: Calculate the contribution of each isotope:
Iron-55: (55 amu) x (0.15) = 8.25 amu
Iron-56: (56 amu) x (0.85) = 47.6 amu
• Step 2: Add the contribution from each isotope
together:
8.25 amu + 47.6 amu = 55.85 amu
The element with an atomic mass of 10.8 amu is Boron.
IONS
 Ion:
formed when an atom gains or loses
an ELECTRON and becomes charged.

Gained an electron  (-) chargeNa1+


(Has more electrons than protons)
Lost an electron  (+) charge

(Has more protons than electrons)
O2-
IONS



Examples: Na1+  lost 1 electron
O2-  gained 2 electrons
To find # of electrons, start with # of protons and do the
opposite what the charge says.
Example: Na1+#p = 11 (-1) = 10 electrons
O2- #p = 8 (+2) = 10 electrons
WARM UP FOR THURSDAY, SEPTEMBER 11

Identify the number of protons, neutrons, and
electrons for the following (write the complete
isotopic symbol next to your answers)
N-3
Carbon-16
Mg+2
55Fe+2








CFrequency
HARACTERISTICS OF WAVES (CONT’D)
of Wave A: ____. Frequency of Wave B: _____
As wavelength increases, frequency ________________.
As wavelength decreases, frequency _________________.
Wavelength and frequency have a/an _____________
relationship!
As frequency increases, the energy of the wave
_____________.
As the wavelength increases, the energy of a wave
__________.
Energy and frequency have a/an _____________
relationship!
Energy and wavelength have a/an _____________
relationship!
ELECTROMAGNETIC SPECTRUM
FLASHCARDS
 Arrange
in order from least energy 
most energy.
 Hint:
The more energy an
electromagnetic wave has, the more
dangerous it is for us.
WAVE ENERGY & HEALTH RISK
LOW ENERGY, LOW FREQUENCY, LONG
WAVELENGTH
- AM/FM
 Radio Waves
- Microwave ovens; Cell phones
 Microwaves
- Warm objects; Night vision
 Infrared
goggles
 Visible Light Red
 Violet
 Ultraviolet
- Sunburns; skin cancer
 X-Rays
- Goes through skin, reflected by
 Gamma Rays
bone
- Causes cancer; Used in radiation
therapy to kill cancer cells; given off
HIGH ENERGY,byHIGH
FREQUENCY,
SHORT
nuclear
reactions
WAVELENGTH
ELECTROMAGNETIC SPECTRUM
BOHR’S PLANETARY MODEL
1913

Electrons orbit the nucleus like
planets orbit the sun. [WRONG!]
BOHR: ELECTRONS IN ENERGY LEVELS
•
Bohr studied how atoms emit or absorb light.
Here’s what he discovered that still holds true
today…

Electrons can occupy only certain energy levels.

Energy Levels: Regions of space in which e- can
move about the nucleus of an atom.
 Given more energy, e- will move to an energy level
farther away from the nucleus.
 Eventually, electrons will fall back to lower energy
levels, releasing electromagnetic radiation (light).
EVIDENCE FOR ENERGY LEVELS
Evidence for Energy Levels
 EACH ELEMENT HAS ITS OWN EMISSION
SPECTRUM (WAVELENGTHS OF THAT ARE
RELEASED).
(Argon)
BOHR MODEL – TERMS (IN PACKET)




Photon: A particle of electromagnetic radiation
with no mass that carries a quantum of energy.
Quantum (pl. quanta): “a packet of energy”; the
amount of energy absorbed or released when an
electron moves between energy levels.
Ground State: lowest possible energy of an e-.
The principal quantum number (n) = 1. [Energy
level = 1]
The most energy is released when electrons
fall to the n=1 energy level.
BOHR MODEL
OF THE HYDROGEN
ATOM
Examples: What type of
light is released?
1) n=5 to n=3
2) n=6 to n=2
3) n=1 to n=3
FLAME TEST LAB


Safety
 Bunsen Burners (fire) –
keep all papers away from
flame, wear goggles
 Gas valve – ON and OFF
positions
 Chemicals – wash hands if
get on skin
Disposal
 Please put well plates
back in baggies,
throw away Q-tips, wipe
down lab tables.
IS AN ELECTRON A PARTICLE OR A WAVE?

Wave-Particle Duality: electrons have both
wave-like and particle-like properties.
Video –
Double Slit
Experiment
(5 min)
Warm-Up
1.) Put in order from lowest to highest frequency: Radio
waves, gamma rays, IR, X-rays, UV, Microwaves
2.) If violet light of wavelength 4.1 x 10-7 m is released,
what transition did the electron make?
n = __________ to n = __________
3.) What is the difference between nuclear fussion and
nuclear fission?
Did you finish your pre-unit notes on Nuclear Chem?
FINISH NOW!
Nuclear Chemistry
Nuclear Decay
Henri Becquerel – discovery of Uranium
 Marie Curie - radioactivity

Nuclear Reaction

Nuclear Reactions: a nucleus loses or gains protons
and neutrons.

Why do some nuclei undergo radioactive decay?
1) They are too big (too many protons)!
- All elements with atomic numbers of 84
or higher are radioactive!
2) There are too many neutrons compared to
protons.
Alpha (α) Decay



Alpha (α) Decay: Nucleus releases an alpha particle.
Alpha particle: helium nucleus = 4 He
2
 2+ charge
 Lowest energy radiation.
 Can be stopped by a
sheet of paper or skin.
 Reduces mass number by
4 and atomic number by 2.
A new element is created
(transmutation) because the
atomic number changed.
Alpha Decay Problems
Example: 238
U
92
 Practice:
1. 234 Th 


90
2. 218
84 Po 
3. Write an equation that represents the
alpha decay of Rn-222.
Beta (β) Decay




no
p+
e-
Beta (β) Decay: A neutron in the nucleus is converted
into a proton and an electron. The electron is created
INSIDE the nucleus and is emitted as a beta particle.
Beta particle = a fast moving electron sent shooting out
of nucleus
 Negative charge
 Can be stopped by aluminum foil or a piece of wood.
Increases the atomic number by 1 and does not change
the mass number.
A new element is created (transmutation) because the
atomic number changed.
Beta Decay Examples
238
92

Example:

Practice:
218 Po
1.
84
U


2. Write the equation for the beta decay of
Pb-214.
Gamma Decay





Gamma (γ) Decay: A gamma ray (high energy
electromagnetic wave) is released from the nucleus.
 Highest energy radiation.
 Most dangerous.
 Can be stopped by several cm of lead or several
meters of concrete.
The nucleus goes from an excited state to a normal
(unexcited) state.
Almost always occurs with alpha or beta decay.
Does not change mass number or atomic number.
No new element is created
Gamma Decay Examples

* excited state
238
92
218
84
U*

Po* 
How would alpha, beta, and gamma
radiation be affected by an electric
field?
Positive (+) Plate
Lead Box with
Radioactive
Material
Negative (-) Plate
Nuclear Fission



Nuclear Fission: A large nucleus is split into two or
more nuclei.
 A neutron is sent into a uranium nucleus. The U
nucleus splits into two smaller nuclei and three
neutrons are released to hit other U nuclei.
 Nuclear chain reaction: continuous series of fission
reactions
 Lots of energy released.
Atom bomb = uncontrolled fission reaction.
Nuclear power plants use controlled fission reactions
to make electricity.
Nuclear Fission
Fission Chain Reaction
Uncontrolled Fission
Reaction – Hydrogen Bomb
A controlled fission chain
reaction in a nuclear reactor
Nuclear Fusion

Nuclear Fusion: Two or more small nuclei
combine to form a larger nucleus.



Occurs on the sun and other stars (not on Earth!).
Produces enormous amounts of energy
Why can’t we use this energy as a source of
electricity???
To start a fusion reaction, temperature must be
200 million Kelvin (360 million degrees Fahrenheit).
 Difficult to initiate and contain this reaction due to high
temperature required (Remember Spiderman?…).

An attempt at controlling a
fusion reaction on Earth…
Nuclear Fusion
2H
1
+
3
1
H

4
2
He + 1 n
0
Half-Life



Radioactive Half-Life (t1/2): the amount of time it takes
for half of the atoms to undergo decay.
Use: dating fossils (carbon-14 dating), geological
formations and human artifacts
Example Half Lives:
 Potassium-40: Half-life = 1.25 years
 Carbon-14: Half-life = 5739 years
 Uranium-238: Half-life = 4.5 billion years
 Rubidium-87: Half-life = 48 billion years
Half-Life
Example Problems
1)
The half-life of carbon-14 is 5739 years. You start
with a sample of 16 grams of C-14.
a) How much will you have after 5739 years?
b)
How much will you have after 11,478 years?
c)
What fraction of the original amount will be left
after 4 half-life periods?
d)
After 7 half-life periods?
Half-Life
Example Problems
2)
A radioactive substance with a half-life of 8 years is
allowed to decay. If the sample started out with 80 grams
and ended with 10 grams, how many years have passed?
Start Amount Half Life End Amount Total Time -
3)
You started with a sample of 70 grams of a radioactive
material. After 30 years, you only had 8.75 grams left.
What is the half-life of this substance?
Start Amount Half Life End Amount Total Time -
Half-Life
Example Problems
4)
You have 60 grams of Potassium-40 (half-life = 1.25
years). If this sample has been decaying for 10 years,
how much did you start with?
Start Amount Half Life End Amount Total Time -
Uses of Nuclear Radiation




Dating Fossils
Smoke Detectors
 Contain Americium which gives off alpha particles –
producing a current. Smoke interrupts the current.
Detection of Diseases
 Radioactive tracers
Radiation Therapy for Cancer
Why is radiation
dangerous?
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Alpha and Beta particles are charged.
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They can ionize, or change the number of
electrons in the atoms in your body.
Gamma Rays are energy (not charged).
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They can give enough energy to the electrons in
atoms to allow them to escape the atom (leaving
an ion behind).
How much radiation have
you been exposed to?
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Have you ever….
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Traveled in an airplane?
Gone through luggage inspection at the airport?
Worn a luminous watch?
Smoked?
Watched TV?
Sat in front of a computer screen?
Been near a smoke detector?
Had an x-ray?
Lived within 50 miles of a nuclear power plant?
Most organisms are adapted to survive low
levels of background radiation.
Nuclear Waste

Fission
 The products of fission reactions are often also
radioactive.  used fuel rods must be stored in
shielded containers away from ground water…
forever…
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Fusion
 Produces very fast-moving neutrons. Shielding
material in the reactor would have to be replaced
periodically.