Chapter 7: Models of Atomic Structure
Investigating the Atom
During the 19th century, when many chemists were
studying elements, other scientists were interested in
the way electricity interacts with matter. Because they
knew that air conducts electricity better if it is trapped
and its pressure is reduced, they used gas discharge
tubes in their investigations.
Gas discharge tubes are made of glass and can trap
air and other gases. They are equipped with electrodes
at either end: The positive charged electrode is called
the anode, and the negatively charged electrode is called
the cathode. Pumping air out of the tube can reduce the
pressure inside the tube. (Modern gas discharge tubes
are used in television picture tubes.)
Scientists found that when they reduced the
pressure in the tube, the gases inside began to glow! The
reason was that cathode rays from the cathode were
traveling through the tube toward the anode and were
carrying a negative charge. The search for the electron
began! (Electron is the word they used to describe tiny,
negatively charged bits of matter).
Scientists were puzzled. If atoms contain
negatively charged electrons, why did they usually have
no charge?
Joseph John Thomson (1856-1940) inferred that
atoms must possess something with a positive charge to
balance the negatively charged electrons. He
experimented with tubes of hydrogen gas and concluded
that anode rays must be made up of positive particles.
Thomson called them protons.
Thomson concluded the following:
-all atoms contain both protons and electrons
-all protons are identical, and all electrons are
identical. Electrons differ from protons,
however.
-an electron has a negative charge. A proton
has a positive charge. A proton has a positive
charge.
-an electron has the same amount of charge as
a proton, even though the charges are opposite
in kind.
-a proton has much more mass than an electron.
1938 - The first electron microscope was made by Albert
Prebus and James Hillier at the University of Toronto.
This gave scientists more knowledge about the structure
of the atom.
In 1803, Dalton had said that atoms were indivisible.
By Thomson’s time (1900’s), it was evident that this was
not true! Atoms could be torn apart by high-energy
electricity. Protons and electrons were known to be
sub-atomic particles
Thomson’s view of subatomic particles:
Negative electrons were like raisins in a muffin:
They could be pulled out of the positively charged
dough
In 1895, x-rays were discovered. This led to the study
of radioactivity. During that time, a student of
Thomson’s named Ernest Rutherford was teaching at
McGill University. He won the Nobel Peace Prize in
Chemistry for his work in radiation.
In 1909, Rutherford probed atoms with radioactive
particles. He discovered that atoms must also contain:
-a nucleus: a tiny core that is very small in volume,
dense compared to the rest of the atom, and intensely
positive.
-an electron cloud: an “envelope” that is very large
in volume, light compared to the nucleus, and negatively
charged.
Because the protons in an atom’s nucleus account for less
than half the mass calculated for the nucleus,
Rutherford inferred that the nucleus must also contain
additional, uncharged (neutral) particles. He called them
neutrons and concluded that each neutron would have
roughly the same mass as a proton. His ideas would not
be confirmed until the 1930’s!
In 1912, a Danish scientist, Niels Bohr, began working
with Rutherford in Manchester, England. He was trying
to explain the question put forth by doubters of
Rutherford’s ideas: if positive and negative charges
attract, why didn’t the negatively charged electrons
crash into the positively charged nucleus?
Bohr came up with a model that compared the nucleus to
the sun and the electrons to the planets. Why doesn’t
the gravitational pull of the sun cause the planets to
spiral inward and crash? Because they revolve at just
the right speed to remain in orbit. Similarly, the atom’s
positively charged nucleus exerts a strong force of
attraction on negative electrons. The electrons don’t
spiral inward and crash because they are moving rapidly
in fixed regions around the nucleus. These 3-D, sphereshaped regions are called electron shells. Bohr
concluded that the more energy an electron had the more
distant shell (energy level) it could occupy.
Bohr-Rutherford Model of Atomic Structure
1st
shell
2nd
shell
3rd
shell
This was the first modern view of the atom!
In 1913, Henry Mosely began working in Rutherford’s
laboratory. He discovered a very regular pattern in the
way different elements responded to x-rays. This
pattern provided information about the atomic nucleus of
each atom. He determined an increasing number of
protons as you move through Mendeleeve’s table. #of
protons=Atomic #
Atomic Number: number of protons in an atom.
Example: Fluorine (9) –has nine protons in the nucleus
Therefore, it also has 9 electrons
How many neutrons does it have? You cannot tell by the
atomic number.
You must look at the mass number (based on atomic
mass): gives the total # of protons and neutrons.
Ex: the atomic mass of F = 19
Atomic #
-
= 9
____
10 neutrons
Proper notation:
F
Atomic Mass Units: because atomic mass is not always a
whole number, it is stated in atomic mass units (u).
Atomic mass is not the same as the mass expressed in
grams. It is a ratio. All atomic masses are defined
relative to the mass of carbon12.