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Democritus and Dalton - The Atom as an Indivisible Sphere

The Greek philosopher, Democritus, was the first to envision matter as being composed of tiny particles. He named
these particles atoms, from the Greek word atomos, which means “indivisible”. In 1807, an English schoolteacher,
John Dalton, rediscovered Democritus’ ideal of the atom. Experimental data at the time had shown that elements
combined in the same percent by mass each time the same compound was formed. Dalton proposed that the atom
was the smallest particle of matter, an indivisible sphere. He further stated that all atoms of the same element had
the same mass and the same chemical behaviour, while atoms of different elements had different masses and differed
in their chemical behaviour. Dalton theorized that atoms of different elements are able to combine in fixed number
ratios to produce specific compounds. Though Dalton based his model on chemical evidence, he knew nothing about
the electrical nature of the atom.

The Discovery of the Electron

      The first hint that matter was electrical in nature came in 1834 when Michael Faraday, a British physicist, found
that chemical changes occur when an electrical current is passed through certain chemical solutions. Later in the
nineteenth century, some studies of the effects of electricity on matter were done with gas discharge tubes. These are
glass tubes that contain a gas at a low pressure. Each is fitted with a pair of metal wires called electrodes, one at each
end of the tube, that can be connected to a source of electricity. When this is done, electricity flows through the tube
and the gas within glows. (Modern neon signs work this way.)
      The physicists who first studied this phenomenon did not know what caused the tube to glow, but tests soon
revealed that something was moving from the negative electrode, which they called the cathode, to the positive
electrode, which they called the anode. The physicists called these emissions rays, and because the rays came from the
cathode, they called them cathode rays.
      Simple experiments soon showed that the cathode rays were electrically charged particles, rather than just light
waves. For example, cathode rays could make a paddle wheel turn when it was placed in their way inside the tube.
When a metal plate outside the tube was given a positive charge, the cathode rays bent toward it, which meant that the
particles in the cathode rays were negatively charged. Remember, "unlike charges attract."

The early experiments with gas discharge tubes gave only qualitative information. In 1897 the British physicist J..J.
Thomson constructed a special gas discharge tube to make quantitative measurements of the properties of cathode
rays. In some ways, the apparatus he used was similar to a television picture tube.

       During a typical experiment, a high voltage was passed between the cathode and the anode of the tube. Because
the anode had a hole in it, some of the cathode ray particles passed through and traveled on to strike the face of the
tube at point 1, where they produced a bright spot on a phosphor (a chemical that glows when struck by a beam of
cathode rays) that coated the tube's inner surface. Thomson also fitted the neck of the tube with a magnet and a pair of
metal plates, which could be given electrical charges. He found that if the metal plates were left uncharged while the
cathode rays passed between the poles of the magnet, the path of the negatively charged particles was curved down-
ward to point 2 by the magnetic field. On the other hand, when the magnet was removed and the metal plates were
given opposite electrical charges, the cathode ray particles were deflected upward (toward the positive plate) to point

Thomson found that by carefully controlling the charge on the plates when the plates and the magnet were both
around the tube, he could make the cathode rays strike the tube at point 1 again. In other words, he was able to cancel
the effect of the magnetic field by applying an electric field that tended to bend the path of the cathode rays in the
opposite direction.
      By measuring the strengths of the magnetic and electric fields needed to balance each other, Thomson was able
to calculate the first bit of quantitative information about a cathode ray particle - the ratio of its charge to its mass
(often expressed as e/m, where e stands for the charge and m stands for the mass). The charge-to-mass ratio has a
value of -1.76  1O8 coulombs/gram, where the coulomb (C) is the SI unit of electrical charge, and the negative sign
reflects the negative charge on the particle. Notice that he wasn't able to measure just the charge or just the mass;
instead, his calculations only gave the value of their ratio.
      One of the most important discoveries that Thomson made in his experiments was that he always obtained the
same value for the charge-to-mass ratio, regardless of the materials used to construct the tube or the nature of the
residual gas in the tube through which the cathode rays passed. This suggested that cathode ray particles are found in
everything, and that they are basic, fundamental particles of matter. They are, in fact, the electrons that we discussed
GROUP#2: Determination of the Charge and the Mass of the Electron

   In 1909 a researcher at the University of Chicago, Robert Millikan, designed a clever experiment that enabled him
to measure the electron's charge. During an experiment he would spray a fine mist of oil droplets
above a pair of parallel metal plates, the top one of which had a small hole in it. As the oil drops settled, some would
pass through this hole into the space between the plates, where he would irradiate them briefly with X rays. The X
rays knocked electrons off molecules in the air, and the electrons became attached to the oil drops, which thereby
were given an electrical charge. By observing the rate of fall of the charged drops both when the metal plates were
electrically charged and when they were not, Millikan was able to calculate the amount of charge carried by each
drop. When he examined his results, he found that all the values he obtained were whole-number multiples of -1.60 
10-19 C. He reasoned that since a drop could only pick up whole numbers of electrons, this value must be the charge
carried by each individual electron.

   When atoms lose or gain electrons, as they do when they become ions, their electrical charges always change by
multiples of 1.60  10-19 C. Therefore, when we specify an ion's charge, we are really giving it in units of this size.
For example, an ion with a charge of 1 + has an actual charge, expressed in SI units, of + 1.60  10-19 C, and an ion
with a charge of - 2 has an actual charge of -3.20  10-19 C. These actual charges are rarely needed, however.
Normally we only need to know the number of charges and their sign.

Once Millikan had measured the electron's charge, its mass could then be calculated from Thomson's charge-to-mass
                                           e/m = -1.76  108 C/g

Solving for the mass, m, and substituting -1.60  10-19 C for e gives

                                        m = -1.60  10-19 C
                                             -1.76  108 C/g
                                          = 9.09  10-28 g
More precise measurements have since been made, and the mass of the electron is currently reported to be
9.109389710-28 g.

The Discovery of the Proton and Thomson’s Model of the Atom

    After the discovery of the electron, gas discharge tubes were modified for additional experiments. In one series,
the cathode was shaped like a disk with a hole in its centre, and the tube's inner surface behind the cathode was coated
with a phosphor . In the new experiments, the phosphor was placed where the electron beam could not touch it, yet it
still glowed. Obviously, something was moving through the hole in the cathode, in a direction opposite that of the
cathode rays. It didn't take long to find out that a stream of particles was moving through the hole in the cathode and
that these particles had a positive charge. The beam, for example, could be deflected toward a negatively charged
plate positioned outside the tube. The mechanism for making these positive particles, as the experimenters reasoned
things out, was probably collisions between electrons in the cathode rays and atoms of the residual gas in the tube.
Such a collision knocked off an electron from the atom, leaving behind a particle with a positive charge.

    Unlike cathode rays, the masses of the positive partices varied according to the gas present in the tube. The
lightest positive particles (hydrogen ions, H+) were observed when the tube had hydrogen gas in it, yet these particles
were still about 1800 times as heavy as an electron. When other gases were used, their masses always seemed to be
whole-number multiples of the mass observed for hydrogen atoms. This suggested the possibility that clusters of the
positively charged particles made from hydrogen atoms made up the positively charged particles of other gases. In
other words, since the lightest of all positive particles came from hydrogen, it seemed likely that all gases - and all
matter, in fact - were made of combinations of the particles in hydrogen. The hydrogen atom, minus an electron, thus
seemed to be a fundamental particle in all matter, and this particle was therefore named the proton, after the Greek
proteios, meaning "of first importance."
             By 1903, the evidence was becoming quite clear that atoms were not solid, indivisible spheres as Dalton
had proposed. Atoms were made up of even smaller subatomic particles. Thomson’s model of the atom incorporated
the newly discovered evidence that atoms are composed of electrons and protons. In his plum pudding or blueberry
muffin model, Thomson suggested that the bulk of an          atom was composed of positive charges, with electrons
dispersed throughout, like plums in a pudding or blueberries in muffin, in order to give the atom an overall neutral
GROUP#3: The Discovery of the Atomic Nucleus

    The effect of the cathode ray (an electron beam) on matter that led to the discovery of the proton suggested many
similar experiments. The subatomic "bullets" did not all have to be made by gas discharge tubes, either. Elements had
been discovered that spontaneously sent out showers of subatomic particles in a phenomenon called radioactivity.
Some radioactive elements send out electrons, but others send out much larger particles having masses four times
those of the proton and bearing two positive charges. These were called alpha particles.
    Early in this century, Hans Geiger and Ernest Marsden, working under Ernest Rutherford at Great Britain's
Manchester University, studied what happened when alpha rays hit thin metal foils (See Figure 6). Most of the alpha
particles sailed right on through as if the foils were virtually empty space. A significant number of alpha particles,
however, were deflected at very large angles. Some were even deflected backward, as if they had hit stone walls.
Rutherford was so astounded that he compared the effect to that of firing a 15-in. artillery shell at a piece of tissue
paper and having it come back and hit the gunner! He reasoned that only something extraordinarily massive,
compared with the alpha particle, could cause such an occurrence. From studying the angles of deflection of the
particles, Rutherford determined that whatever it was in the foil that was so massive had to be positively charged.
However, since most of the alpha particles went straight through, he further reasoned that the metal atoms in the foils
must be mostly empty space. Rutherford's conclusion was that virtually all of the mass of an atom must be
concentrated in a particle having a very small volume located in the center of the atom. He called this massive particle
the atom's nucleus.

The Discovery of the Neutron and the Rutherford Model of the Atom

    From the way alpha particles were scattered by a metal foil, Rutherford and his students were able to estimate the
number of positive charges on the nucleus of an atom of the metal. This had to be equal to the number of protons in
the nucleus, of course. But when they computed the nuclear mass based on this number of protons, the value always
fell short of the actual mass of the atom. In fact, Rutherford found that only about half of the nuclear mass could be
accounted for by protons. This led him to suggest that there were other particles in the nucleus that had a mass close
to or equal to that of a proton, but with no electrical charge. This suggestion initiated a search that finally ended in
1932 with the discovery of the neutron by Sir James Chadwick, a British physicist. Chadwick received the 1935
Nobel Prize in physics for his discovery of the neutron. Table 1 summarizes the properties found for the three
principal subatomic particles.

Table 2: Properties of Subatomic Particles
PARTICLE              MASS (amu)              MASS (g)                  ELECTRICAL CHARGE
Electron             0.0005485712           9.1093897  10-28           - 1 or - 1.60  10-19 C
Proton               1.00727605             1.6726231  10  -24
                                                                        +1 or +1.60  10-19 C
Neutron              1.008665               1.674954  10 -24
                                                                         0 (no charge)
    The picture of the atom that emerged from the work of Rutherford and his students - a small dense nucleus
containing protons and neutrons surrounded by electrons in the remaining volume of the atom - raises important
chemical questions. Although the nucleus determines the mass of the atom and the number of electrons needed to give
the atom electrical neutrality, the nucleus does not play a direct part in chemical reactions. When two or more atoms
join together to form a compound, the nuclei of the atoms stay relatively far apart. Only the atoms' outer reaches - the
regions inhabited by electrons - come in close contact. The chemical properties of the elements, then, must be
determined by the electrons of the various atoms, and the similarities and differences in these properties must have
something to do with the way these electrons are distributed around the particular nuclei.
    How electrons are distributed about the nucleus is called the atom's electronic structure or electron configuration.
The basic clue to the electronic structures of the various elements comes from the study of the light emitted when
atoms of the elements are excited, or energized.
GROUP#4: Atomic Spectra - Some Background

      A continuous spectrum contains light of all colors. It is formed when the light from the sun, or any other object
that's been heated to very high temperatures (such as the filament in an electric light bulb), is split by a prism and
displayed on a screen. A rainbow after a summer shower is a continuous spectrum that most people have seen. In this
case, the colors contained in sunlight are spread out by tiny water droplets in the air.
    A somewhat different kind of spectrum is produced if we examine light that is given off by a gas such as hydrogen
when an electric discharge passes through it. This discharge is an electric current that excites, or energizes, the atoms
of the gas, and they emit this energy in the form of light as they return to a lower energy state. When a narrow beam
of this light is passed through a prism a continuous spectrum is not produced. Instead, only a few colors are observed,
displayed as a series of individual lines. This series of lines is called the element's atomic spectrum. Each element has
its own unique atomic spectrum that is as characteristic as a fingerprint.

The Energy of a Light Wave - A little Physics!

    In 1900, Max Planck, a German physicist, launched one of the greatest upheavels in the history of science when he
proposed that electromagnetic radaiation is emitted only in tiny packets of energy later called photons. Each photon
pulses with a frequency, v, and each travels at the speed of light. Planck proposed and Albert Einstein confirmed that
the energy of a radiation is proportional to its frequency, not its intensity or brightness as had been believed up to that
                                 energy of a photon = E = h  v
where h is Planck’s constant or 6.63  10-34 J s. The energy of one photon is called one quantum of energy. Planck’s
and Einstein’s discovery was really quite surprising. If a particular event requiring energy, such as photosynthesis in
green plants, is initiated by the absorption of light, it is the frequency of the light that is important, not its intensity or

The Significance of Atomic Spectra on the Development of an Atomic Model

    Planck described a simple relationship between the frequency of light and its energy, E = hv. The fact that excited
atoms emit light of only certain characteristic frequencies tells us that only certain characteristic energy changes take
place within the atom. For instance, in the spectrum of hydrogen there is a red line that has a wavelength of 656 nm
and a frequency of 4.571014 Hz. Thus,
                                               E = 6.63  10-34 J s  4.571014 Hz
                                           E = 3.03  10-19 J
The energy of each photon of this light is 3.03  10-19 J. Therefore, when a hydrogen atom emits light of this
frequency, the energy of the atom decreases by 3.03  10-19 J . What is very special here is that whenever a hydrogen
atom emits red light, its frequency is 4.57  1014 Hz and the energy change within the atom is always exactly 3.03 
10-19 J, never more and never less! Atomic spectra, then, tell us that when an excited atom loses energy, not just any
arbitrary amount is lost. Only certain specific energy changes can occur, which is why only certain specific
frequencies of light are emitted. This is the only way atomic spectra can be explained.
    How is it that atoms of a given element always undergo exactly the same specific energy changes? The answer
seems to be that in an atom an electron can have only certain definite amounts of energy and no others. In the words
of science, we say that the electron is restricted to certain energy levels. We also say that the energy of the electron is
quantized, meaning once again that the electron's energy in a particular atom can have only certain values and no
    The energy of an electron in an atom might be compared to the potential energy of a ball on a staircase. The ball
can only come to rest on a step, and on each step it will have some specific amount of potential energy. If the ball is
raised to a higher step, its potential energy will be increased. When it drops to a lower step, its potential energy
decreases. But each time the ball stops, it stops on one of the steps, never in between. Thus, the ball at rest can only
have certain specific amounts of potential energy, which are determined by the energy levels of the various steps of
the staircase. So it is with an electron in an atom. The electron can only have energies corresponding to the set of
electron energy levels in the atom. When the atom is supplied with energy, as in a gas discharge tube, an electron is
raised from a low-energy level to a higher one. When the electron drops back, energy equal to the difference between
the two levels is released and emitted as a photon. Because only certain energy jumps can occur, only certain frequen-
cies can appear in the spectrum. The existence of specific energy levels in atoms, as implied by atomic spectra, forms
the foundation of all theories about electronic structure. Any model of the atom that attempts to describe the positions
or motions of electrons must also account for atomic spectra and the arrangement of electrons in an atom.
     GROUP#5: The Bohr Model Of The Atom - The Early Quantum Model

         Current theories of electronic structure are referred to as quantum mechanics because they predict that electrons
     are found in quantized energy levels. The first model of the atom to meet with some success imagined the electron to
     be revolving about the nucleus in orbits of fixed energy. The discovery that the energy of electrons is quantized led
     to attempts to develop theoretical models of the way electrons behave in atoms. The goals were to explain how
     electrons move, where they are located, and how they change energy to give off photons of light. Physicists were
     faced with a problem, however. None of the physical laws that seemed to govern the motion of large objects, like
     baseballs or people, were able to account for the strange behaviour of electrons.
         In 1913 Niels Bohr (1885 - 1962), a Danish physicist, proposed a theoretical model for the hydrogen atom. He
     chose hydrogen because its atoms are the simplest, having only one electron about the nucleus, and because it
     produces the simplest spectrum with the fewest lines. In his model, Bohr imagined the electron to move around the
     nucleus following fixed paths, or orbits, much as a planet moves around the sun. His model also restricted the sizes of
     the orbits and the energy that the electron could have in a given orbit. The equation Bohr derived for the energy of the
     electron includes a number of physical constants such as the mass of the electron, its charge, and Planck's constant. It
     also contains an integer, n, that Bohr called a quantum number. Each of the orbits can be specified by its value of n.
     If all the constants are combined, Bohr's equation is
     where E is the energy of the electron and k is the combined constant (its value is 2.18  10-18 J). The allowed values of
     n range from 1 to , with all integers permitted (i.e., n could equal 1, 2, 3, 4, . . . , ). Therefore, the energy of the
     electron in any particular orbit could be calculated.
         Because of the negative sign in Bohr’s equation, the lowest (most negative) energy value occurs when n = 1,
     which corresponds to the first Bohr orbit. This lowest energy state is called the ground state. According to Bohr's
     theory also, this orbit brings the electron closest to the nucleus.
         When the hydrogen atom absorbs energy, as it does in a discharge tube, the electron is raised from the orbit having
     n = 1 to a higher orbit, n = 2, n = 3 or even higher. Then, when the electron drops back to a lower orbit, energy is
     emitted in the form of light. Since the energy of the electron in a given orbit is fixed, a drop from one particular orbit
     to another, say from n = 2 to n = 1, always releases the same amount of energy, and the frequency of the light emitted
     because of this change in energy is always precisely the same.
         Bohr's model of the atom was both a success and a failure. It successfully predicted the frequencies of the lines in
     the hydrogen spectrum, so there seemed to be some validity to the theory. Nevertheless, the model was a total failure
     for atoms with more than one electron! Still, though the theory met with only limited success, introduction of the
     ideas of quantum numbers and fixed energy levels was an important step forward.

     The Wave Nature of Matter - Development of the Wave Mechanical Model of the Atom

         Bohr's efforts to develop a theory of electronic structure were doomed from the very beginning because the
     classical laws of physics - those known in his day - simply do not apply to particles as tiny as the electron. Since all of
     the objects that had been studied by scientists until that time were large and massive in comparison with the electron,
     no one had yet detected the limits of classical physics. Classical physics fails for atomic particles because matter is
     not really as our physical senses perceive it. Under appropriate circumstances, small bits of matter, such as an
     electron, behave not like solid particles, but instead like waves. This idea was first proposed in 1924 by a young
     French graduate student, Louis de Broglie. De Broglie’s ideas were confirmed by experiments that showed that a
     beam of electrons could be bent, or diffracted, by passing it through a crystal in just thesame way that light is
     diffracted by passing it through a glass prism. Bohr was unable to accurately define the arrangement of electrons
     because he still considered them as particles, goverened by the same classical laws of physics that control the motion
     of planets.
         Light waves are characterized by their wavelengths and their frequencies. The same is true of matter waves. De
     Broglie suggested that the wavelength of a matter wave, , is given by the equation
                                              = h
     where h is Planck's constant, m is the particle's mass, and v is its velocity.

The concept of a particle of matter behaving as a wave rather than as a solid object may at first seem difficult to
comprehend. Your text book certainly seems solid enough, and if you drop it on your toe, it surely doesn't seem to be a
wave, at least not as we generally think of waves in the ocean. The reason for the book's apparent solidity is that in de
Broglie's equation the mass appears in the denominator. This means that heavy objects have extremely short wavelengths.
The peaks of the matter waves for heavy objects are so close together that the wave properties go unnoticed and can't even
be measured experimentally. But tiny particles with very low masses have much longer wavelengths; therefore, their wave
properties become an important part of their overall behavior. You may not feel the teeth of a comb if they are very thin
and placed extremely close together - the comb may feel as if it is completely solid. In a similar fashion, you are unable to
notice the wave properties of the matter around you. The theory of electronic structure based on the wave properties of the
electron is

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