CHM 1020--Chemistry for Liberal Studies--Fall 2000
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CHM 1020--Chemistry for Liberal Studies--Fall 2000 |
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Dalton postulated an atom that was a “hard structureless sphere”, something that was not divisible into something simpler. He held to this notion even in spite of later evidence that suggested some elements were made of diatomic molecules. Common gaseous elements such as oxygen, nitrogen, hydrogen, chlorine, are actually O2, N2, H2, and Cl2, for example. He refused to believe that.
But
other evidence was accumulating in the early part of the 19th century
that indicated there must be some sub-structure to atoms.
One was in the discovery of electricity, and the observations
that passing electric currents through substances could cause chemical change.
A few weeks after the Italian scientist Volta discovered the battery,
Nicholson and Carlisle, carried out the decomposition of water to hydrogen
and oxygen by passing an electric current through it.
This was in 1800, three years before Dalton stated his atomic theory.
The British chemist Humphry Davy and his student Michael Faraday
proceeded to experiment with electric current, discovering a number of new elements.
Passing an electric current through molten salts, such as potassium hydroxide
and sodium hydroxide, produced the highly reactive metals potassium and sodium.
This process of splitting compounds into elements with an electric current
was known as electrolysis.
Faraday
postulated that the current was carried in the molten salts by charged atoms
which later were called ions. In
an electrolytic cell, the positive electrode is called the anode, and
the ions attracted to it are called anions. Anions therefore carry a negative charge.
The negative electrode is called a cathode, and the ions attracted
to it are called cations. Cations
therefore carry a positive charge. By
measuring the mass of an element produced by a given amount of current, one
can get a mass to charge ratio that was proportional to the known relative
atomic weights measured by other means.
Towards
the end of the 19th century, experiments with higher voltage electrical
sources established that electrical current could actually flow in a vacuum,
provided sufficient gas was removed from the vacuum tube.
It was established using barriers in the tube that something was flowing
from the negative electrode (the cathode) to the positive electrode.
(Figure 3.4) This something was called a cathode ray, and experiments
designed to study the nature of the cathode ray by the English physicist Joseph
Thomson established the nature of the particles in this “ray”.
(See Figure 3.5). They are negatively charged, and by measuring the extent of
their deflection in a magnetic field it could be determined that mass to
charge ratio was much smaller (about 1/2000) than that of the lightest then
known substance, hydrogen. This
observation strongly suggested one was dealing with a particle much smaller
than an atom. The particles in
the cathode ray later were named electrons.
Therefore
atoms are not indivisible. They
can be broken down into a small negatively charged particle, an electron,
and a positive ion. Thus
another of Dalton’s tenets had to be modified.
One
can deduce a great deal from dealing with ratios, such as the ratios
of reacting substances in a chemical reaction or the mass to charge ratios in
the various electrical experiments. A
clever experiment just after the turn of the century provided us with a measure
of the size of the charge on the electron, allowing us to actually calculate
the size of its mass and therefore the mass of various ions.
This is the oil drop experiment of Robert Millikan.
(See Figure 3.7). Tiny oil
droplets are sprayed into a chamber and aquire a negative charge of static electricity. The mass of an individual particle can be measured by the rate
at which it settles in the chamber by the pull of gravity.
A positively charged plate placed at the top of the chamber can attract
the droplets upward, opposing the pull of gravity.
The voltage required to just arrest the fall of the particle will give
the mass to charge ratio of the particle, and combining this value with the
mass determined by the rate of fall allows the calculation of the charge.
The charges so measured were all multiples of a single value, which we
now know represents the charge on the electron, and which we represent simply
as –1.
(We
might think of this discovery as sort of an atomic theory for electrical charge.
The atomic theory says that matter comes in discrete units called atoms.
This observation tells us that there is a “smallest unit” of electrical
charge—the size of the charge on the electron).
With
this discovery of the electron, Thomson developed a revised model of the atom,
a model in which electrons were stuck in the atom like raisins in a plum pudding,
and his model was called the pudding model.
Other
observations around the turn of the century produced other problems with Dalton’s
theory.
The
first Nobel prize in physics was awarded in 1901 to the German scientist named
Wilhelm Roentgen. While studying
the effect of cathode rays on substances, he discovered a new kind of ray, called
an x-ray, which could travel through walls, and which when passing through
his hand onto a film could produce an image of the bones in his hand.
This discovery caused quite a public stir.
The discovery of rays that could penetrate clothing raised great privacy
issues. It was know that the rays could not pass through heavy metals
such as lead, and lead lined underwear became a popular item.
Interestingly,
the discovery of x-rays, and experiments designed to study this phenomenon,
lead to an even more interesting discovery.
It was know that some chemicals when exposed to strong sunlight would
produce a fluorescence, a glow of a different wavelength of light.
You have probably all seen examples of this when certain minerals are
exposed to ultraviolet light (“black light”).
A Frenchman, Henri Becquerel, was studying this phenomenon to see if
x-rays could be produced by fluorescence.
He would expose minerals to sunlight while the mineral was placed on
an envelope containing photographic film.
If x-rays were produced, they should produce an image of the mineral
on the film.
It
was a cloudy day in Paris one day, and Becquerel put aside his minerals and
x-ray film in the drawer without carrying out the experiment.
For some reason, he later developed the film, and found the image of
the mineral exposed on the film!. Something
was being emitted by the mineral (a compound of uranium) that passed through
the envelope and exposed the film. Further
studies showed this to be a characteristic of uranium, and the phenomenon became
known as radioactivity, the spontaneous emission of radiation from certain
unstable elements.
Becquerel
shared the 1906 Nobel prize in physics with Marie and Pierre Curie, who following
up on Becquerel’s discovery were able to isolate other radioactive elements
radium and polonium. After her
husband's death in a traffic accident, Marie Curie continued her work with radioactive
elements, and she won the 1911 Nobel prize in chemistry, the first person for
50 years to win two prizes.
Ultimately,
three types of radioactivity were characterized.
Particle |
Charge |
Mass |
|
alpha |
2+ |
4 |
|
beta |
1- |
0
(1/1837) |
|
gamma |
0 |
0 |
The
charge and charge to mass ratio could be determined by experiments with electrical
plates and photographic film, as in Figure 3.9.
The particles causing exposure of X-ray film in Becquerel’s experiments
were the alpha particles.
Ernest
Rutherford, a New Zealander who worked in Cambridge, England, used these particles
to study the nature of matter. He
devised a way to create a beam of alpha particles from a sample of radioactive
mineral, and he and his students proceeded to study how bombarding various types
of matter with these “rays” affected the matter.
One of his more famous experiments was that of bombarding thin gold foil
with the alpha particles. (See
Figure 3.10)
Rutherford
deduced that for most of the particles to pass through the gold atoms, most
of the gold atoms must be made of empty space.
But for some to be deflected, the mass of the gold must be contained
in a very small region at the center of the atom called the nucleus.
(See Figure 3.11) While an atom has a size of about 10-10
m, the diameter of the nucleus is about 10‑14 m, about 1/10,000
the size. If the atom is the size
of Campbell Stadium, the nucleus is about the size of the coin tossed at the
beginning of a football game.
For
the nucleus to repel the positively charged alpha particle, it must be positively
charged. According to Rutherford,
then, the space around the nucleus must be taken up by the very light, negatively
charged, electrons. The characteristic
thing about the atom of an element was therefore not its atomic mass, but its
nuclear charge, also called its atomic number.
Rutherford visualized these electrons spinning around the nucleus much
as planets spin around the sun, hence his model of the atom was called the “planetary
model”.
The smallest nucleus was that of hydrogen, with a mass of 1 and a charge of plus 1. Therefore, the hydrogen atom consisted this nucleus and one electron surrounding it. The nucleus of the hydrogen atom is called a proton.
Discovery of another elementary particle by Chadwick, one of Rutherford’s students, allows us to develop a rather simple model for the atom. Chadwick discovered the neutron, which has a mass of 1, equivalent to the proton, but lacks an electrical charge.
The alpha particle,
then, would be composed of two protons, imparting the 2 positive charges and
2 mass units, and 2 neutrons, adding 2 mass units for a total of 4.
If the alpha particle picks up in addition two electrons, it becomes
a helium atom.
Therefore, according to Rutherford, the nucleus of atoms contain protons and neutrons. The protons give a positive charge to the nucleus, and the number of protons determines the atomic number of the element, which is the important variable to its identity. The neutral atom contains enough electrons around the nucleus to balance the charge. An ion is created when an electron is taken away (forming a positive ion) or is added (forming a negative ion). Isotopes are explained by atoms that have the same number of protons, and hence the same atomic number, but different numbers of neutrons, and therefore different masses.
We can easily summarize this atomic structure in the symbolism we use for the elements as follows:
A q
X
Z n
Where:
The
mass number is the atomic mass rounded to the nearest whole number.
If you examine atomic weights, you will find some that are very near
whole numbers, some that are not. One
reason for deviation from whole numbers is the the fact that the atomic weights
are averages for a mixture of isotopes.
A given isotope will have a mass near a whole number but will deviate
by a small amount for a reason we will see in the next chapter.
Fill
in the following table to see how one describes a given atom or ion in terms
of this sub-atomic particle model:
Atom or ion |
#
of protons |
#
of neutrons |
#
of electrons |
|
7Li |
3 |
4 |
3 |
|
16O |
8 |
8 |
8 |
|
35Cl- |
17 |
18 |
18 |
|
23Na+ |
11 |
12 |
10 |
|
56Fe2+ |
26 |
30 |
24 |
|
32S2- |
16 |
16 |
18 |
The
next step in our model development comes from studies of the way that light
interacts with atoms.
If
we put compounds in flames, they produce colors which are characteristic of
the elements in the compounds.
See figures 3.13 and 3.17.
To
understand Bohr’s argument, we need to diverge a little bit and talk about light
as a form of energy called electromagnetic radiation.
Unfortunately, your textbook does not take any space to this topic.
As
light passes through space it disturbs the electrical and magnetic fields of
space. This disturbance gives it
the property of a wave, just as the
surface of a lake is disturbed.
Wavelength
x frequency = velocity,
or ln=c
C
is the velocity of light, which in a vacuum is 3.00 x 108 ms-1.
Visible
light spans from l
of 400 nm to 700 nm.
Spectrum
of rainbow: ROY G. BIV (700nm®400nm)
Slightly
longer wavelengths are the infrared;
slightly shorter is the ultraviolet.
Known
spectrum extends from l
of 10-14 m to 104 m as follows:
Cosmic
rays—gamma rays—X-rays—UV—visible—infrared—microwaves—TV—radio.
.
Frequencies
vary from 1022 Hertz to 104 Hertz.
(AM radio frequencies are in kilohertz;
FM frequencies are in megahertz)
Do
some conversions:
Wavelength
of WFSQ signal (88.9 MHz)
Frequency
of blue light at 680 nm.
Quantum
theory: sort of an atomic theory for energy. Idea that energy comes in definite packages called quanta,
and that atoms exist only in discrete energy states.
A quantum of electromagnetic radiation is considered to be a particle
called a photon.
Particle-wave
duality
seems to our senses to be a contradiction in terms.
Yet not only does electromagnetic radiation seem to behave that way,
matter at the very tiny level of electrons and below also behaves that way.
Through
a combination of experiments by Max Planck and Albert Einstein, we can put a
quantitative figure on this particle, and relate the particle and energy properties
of light. The basic relationship
is:
E = hn
(or the counterpart E =
hc/l).
h
= Planck’s constant, which has the value 6.63 x 10-34 Js
When
you consider the energy required to do something to a molecule, we can show
that microwave and infrared radiation only have enough energy to enhance the
motions (vibrations) of molecules. Infrared
radiation is radiant heat, and associated
with motions in molecules. When one gets to ultraviolet light, then the energy per quantum
begins to be enough to lead to breakage of chemical bonds in molecules.
That is why from UV to cosmic radiation, the photons are considered hazardous.
You
can explore these relationships further at this web
site.
With this understanding of the nature of light, let us return to Bohr’s model of the hydrogen atom. Most of his quantitative work was with the simplest atom, hydrogen. He suggested that in the planetary model of hydrogen, in which an electron circled the nucleus in an orbit, there were only certain discrete orbits allowed, and each had a particular energy. (The further the electron from the nucleus, the higher its energy, because it takes work to pull the negatively charged electron from the positively charged nucleus). (This was a break from classical physics, which had shown that a moving charge radiates energy, and if the electron behaved that way it would lose energy and collapse into the nucleus).
The allowed orbits constituted energy levels. When hydrogen is put in a flame, all the atoms are excited, and electrons are promoted to higher energy levels. They don’t remain there for long, but fall back to the lower levels, ultimately coming to the lowest, most stable level, called the ground state. As the electrons drop from a high to a lower energy level, they give off a particle of electromagnetic radiation with the energy corresponding to the difference in energy levels in the atom. Therefore by calculating the energies of the photons in the line spectrum, Bohr could calculate the energy differences in the levels of the hydrogen atom, and develop a picture of these allowed energy levels for the electron in hydrogen.
Bohr
used classical mechanics to calculate the kinetic and potential energy an electron
would have in a certain set of defined orbits.
He had to assume that some orbits are stable, and could have only certain
definite energy values defined by the relationship:
E = -2.18 x 10-18J x 1/n2,
where n =1, 2, 3, 4, 5, etc.
When
an electron drops from energy level 2 (-2.18 x 10-18J/4) to energy
level 1 (-2.18 x 10-18J/1), the difference in energy is released
as a quantum of light with the energy E2 – E1, which determined
its frequency and wavelength.
The
wavelengths of light emitted by excited hydrogen, in both the ultraviolet, the
visible, and the infrared regions of the spectrum, corresponded with very high
precision and accuracy (five to six significant figures) to those calculated
by Bohr.
In
spite of this almost perfect fit between theory and experiment, there were some
serious shortcomings of Bohr's theory:
