Transcript History of Atomic Structure

History of Atomic Structure
Chemistry
Democritus (460-370 BC)
• Goal of Greek philosophers was to explain the
natural world
 Believed that all materials could be broken down smaller and smaller parts
until you reach a point
 This point was what Democritus called “atomos” which in Greek meant
“indivisble” or “uncuttable”
• This particulate view of nature was not too popular among
many Greek thinkers due to its rejection by Aristotle
 Matter is made up of 4 elements: fire, air, water, and earth
 Matter was continuous all of one piece
• Unfortunately for mankind, the ideas of Democritus would
not come back into the public domain for nearly 2,000 years!
John Dalton: Atomic Theory
1.) All matter is made up of atoms. Atoms are
indivisible
2.) All atoms of an element are identical in every
respect (have the same masses and same
properties)
3.) Atoms of different elements are different (have
different masses)
4.) Compounds are formed by a combination of 2 or
more different kinds of atoms (same ratio of atoms)
5.) A chemical reaction is a rearrangement of atoms
Dalton’s Atomic Theory
Dalton’s theory seemed to be quite similar to
Democritus don’t you think?
One part of Dalton’s atomic theory has been
rejected
 Make a prediction of which statement you think is incorrect
and a hypothesis as to why you think that
 We will discuss this more later
J.J. Thomson: 3 experiments, 1 big idea
• Do atoms have parts?
 Thomson suggested that they do
 Advanced the idea that cathode rays are really streams of very
small pieces of atoms
 3 experiments led him to this
Thomson’s
st
1
experiment
• Had already been found that cathode
rays deposited an electrical charge
• What Thomson was interested in was
whether or not he could separate the
charge from rays by bending them with
a magnet
• Found that when rays entered slit and
into electrometer, it measured a large
amount of negative charge
• Electrometer did not register much
charge if rays were bent so they would
not enter slit
• Thomson concluded that the negative
charge and the cathode rays must
somehow be stuck together (you
cannot separate the charge from the
rays
Electrometer: device for
measuring electrical charge
Thomson’s 2nd experiment
• Prior to Thomson’s 2nd experiment, all
attempts had failed when trying to bend
cathode rays with an electric field
• A charged particle will normally curve as
it moves through an electric field, but
not if it’s surrounded by a conductor
• Thomson realized that others had failed
probably because traces of gas
remaining in the tube were being turned
into electrical conductors
• To test this idea, he took great pains to
extract nearly all the gas from a tube,
and found that now the cathode rays did
bend in an electric field after all
Notice which way the cathode ray is being
bent after passing through electric field.
Why is this?
After 2 experiments….
“I can see no escape from the conclusion
that cathode rays are charges of negative
electricity carried by particles of matter.
But, what are these particles? Are they
atoms, or molecules, or matter in a still
finer state of subdivision?”—J.J. Thomson
Thomson’s 3rd Experiment
Sought to determine the basic properties of the
particles
 Thomson calculated the ratio of the mass of a particle to its
electric charge (m/e)
What he concluded was astounding:
 The mass to charge ratio for cathode rays turned out to be far smaller than
that of a charged hydrogen atom—more than 1,000 times smaller
2 possibilities:
1.) the cathode rays carried an enormous charge
2.) they were amazingly light relative to their charge
2nd possibility was
eventually proven
Summary of Thomson’s Findings
• Discovered the electron (negatively charged
subatomic particle) in 1897
• Developed the “plum pudding” model in 1904
 Probably easier to call it the “watermelon” model
Seeds: negative particles
Fruit: positively charged
low density material
Ernest Rutherford
• Student of J.J. Thomson (1894)
Early studies based on radioactivity
• Discovered that some materials emit radiation
 Alpha particle: Eventually was confirmed to be helium nuclei, which
meant an alpha particle was simply 2 protons and 2 neutrons
 Like Thomson, Rutherford was interested in the deflection patterns of alpha
particles when exposed to electric & magnetic fields
 Found the deflection patterns by measuring alpha particles position on
photographic film
 By accident, he noticed that if the alpha particles passed through a thin sheet
of mica, the images on the film were blurred
Rutherford and alpha particles
• The images were sharp if the mica was not
present
 Something about the mica sheet was causing the alpha
particles to scatter at seemingly random small angles,
resulting in blurred images
Rutherford & Geiger
• Wanted to study the effects of alpha particles with
matter
 Had to develop a way to count individual alpha particles when they hit the screen
 Found that a screen coated with zinc sulfide emitted a flash of light each time it was hit
by an alpha particle
 Rutherford & Geiger had to sit in a dark room for hours and individually count the
flashes of light
Rutherford asked Geiger to measure the angles of deflection
when the alpha particles were passed through a thin (.00004 cm)
sheet of gold foil
 When the gold sheet was bombarded with alpha particle, Geiger found that the
scattering small
 Results were consistent with Rutherford’s expectations. Because he knew that alpha
particles had a considerable mass and moved quite rapidly, he anticipated that virtually
all of the alpha particles would go through the metal foil without much disruption
Rutherford, Geiger, & Marsden
Marsden was a graduate student working in Rutherford’s
lab and Geiger had suggested that a research project should
be given to Marsden
 Rutherford: “why not let him see whether any alpha particles can be
scattered through a large angle?”
 Marsden had found that a small fraction (perhaps 1 in 20,000)
of the alpha particles were scattered through angles larger
than 90 degrees
Rutherford was in awe
 “It was quite the most incredible event that has ever happened to me in
my life. It was almost as incredible as if you fired a 15-inch shell at a piece
of tissue paper and it came back and hit you!”
What could the results from this gold
foil experiment mean?
The + charge in an atom of gold was known, but
if this charge were collected on a sphere the size
of an atom (like Thomson suggested), the
repulsion would be far too weak
 To explain the force experienced by the alpha particle, the charge and
much of the mass would have to be collected in a much smaller sphere
 Published results in 1911 and proposed a model for the atom that is still
accepted today
 All of the positive charge and essentially of the mass of the atom is
concentrated in an incredibly small fraction of the total volume of the
atom, which he called the nucleus (Latin for “little nut”)
Further findings from Rutherford’s
experiment
• Most of the alpha particles were able to pass right through
• A small fraction came close to the nucleus of a gold atom as
they passed through and were slightly deflected from the
positive-positive repulsion of the alpha particle and nucleus
• But occasionally, an alpha particle would run directly into the
nucleus and would result in a great repulsion that deflected
the alpha particle through an angle of 90 degrees or more
 By carefully measuring the fraction of the alpha particles deflected through large
angles, Rutherford was able to estimate the size of the nucleus (JUST LIKE WE
DID!)
 Found that the radius of the nucleus is at least 10,000 times smaller than the
radius of the atom
 The vast majority of the atom is therefore empty space!!
Someone throw out the plum
pudding!!
Rutherford revised Thomson’s plum pudding model,
showing how electrons could orbit a positively charged
nucleus, like planets orbiting a sun
Because the majority of the “plum-pudding”
atom would be electrically neutral (no charge),
the alpha particles would have no problem
shooting through
Rutherford’s atom
Along comes Niels Bohr
Also student of Thomson and worked with Rutherford
• Bohr, and many other, knew that Rutherford’s model
made no sense based on one specific reason….
 We knew that any charged body (electron) that was in a state of
motion other than at rest or in uniform motion in a straight line, will
emit energy
 Thus the electrons in this “solar system” model would be constantly
emitting energy
 IF that were the case, the electrons would eventually run out of
energy and spiral down into the nucleus and the entire atom would
collapse!
Bohr
Had trouble making sense of line spectra data
with Rutherford’s atom
Bohr’s Model
• Similar nucleus to Rutherford’s with both
protons and neutrons inside
• Negative particles (electrons) are in specific
orbits around the nucleus
Main problem with Bohr model:
• Only accounted for hydrogen atom
• Couldn’t explain multi-electron atoms
(anything other than hydrogen)
Subatomic Particles
• Based on evidence from Rutherford and
Thomson
Location of Subatomic Particles
Atomic Numbers
What is atomic mass?
Also referred to as:
 Atomic weight
 Average atomic mass
 Relative atomic mass
Important characteristic for elements:
 Each element has its own atomic mass
Atomic mass is an average
Atomic mass
• Average of the masses of a number of
different atoms
• Special kind of average called a weighted
average
• Different than the usual average you’re
probably familiar with in math
Understanding Weighted Averages
Even though these are different
models and have different
features, they are both lemonas
due to their distinct lemon-like
shape
Using this analogy, the models of the lemona are similar to the isotopes of an element
• 29 protons makes both atoms
copper—even though they
differ in their numbers of
neutrons
• Just like the “lemon-like” shape
of a car makes it a lemona—
even though they differ in their
features
Average vs. Weighted Average
What is the average weight of the 2
cars?
4,000+5,000
•
= 4,500
2
• Regular Average
What would happen if we added in extra information?
Weighted Average
What is the average weight of lemonas, taking
into account the amount of each model?
4,000 𝑥 .95 + (5,000 𝑥 .05)
= 4,050 𝑙𝑏𝑠
Because there are so many more GXs than GXLs, the weighted
average is much closer to the actual weight of a lemona GX
Using Weighted Average with Different
Atoms
• Atomic mass: a weighted average of the masses for
all the isotopes of a certain element
Mass = 63 amu
Mass = 65 amu
If we pulled out a random sample of 100 copper atoms, we would find that 69% of
them would be Cu-63 and 31% of them would be Cu-65
• 69% : Cu-63
• 31% : Cu-65
Atomic Mass of Copper
69% : Cu-63
31% : Cu-65
63 𝑥 .69 + 65 𝑥 .31 = 63.62 𝑎𝑚𝑢
Why are the 2 atomic masses different then?
Difference between mass number and
atomic mass
Mass number: protons + neutrons
1 proton or 1 neutron = 1 amu
• Therefore, if you have 6 protons and 6
neutrons, your atom is going to weigh 12
amu
• If you have 6 protons and 7 neutrons, your
atom is going to weight 13 amu
Practice
Gallium has 2 stable isotopes, and the masses of Gallium-69
(60.11% abundant) and Gallium-71 (39.89% abundant) are
68.926 amu and 70.925 amu, respectively. Calculate the average
atomic mass of Gallium
Practice
Rubidium has 2 isotopes: Rubidium-85 (atomic mass of 84.911
amu) and Rubidium-87 (atomic mass of 86.909 amu). The
atomic mass of Rubidium reported on the periodic table is 85.47
amu. Based on this information, which of the isotopes of
Rubidium is more abundant? How do you know?
•
This is more of a thought problem. No real calculations are really necessary
Practice
Magnesium has 3 stable isotopes. Calculate its average
atomic mass, using the information in the chart below.
Isotope
Mass
Abundance
Mg-24
23.985 amu
78.99 %
Mg-25
24.986 amu
10.00 %
Mg-26
25.983 amu
11.01 %