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The Atomic Theory of Matter
The understanding of the nature of matter as atoms was first set down by the British scientist
John Dalton in the early 19th century.
Dalton’s theory can be summarized by the following:
1)
All matter is made of atoms. Atoms are indivisible and indestructible.
2)
All atoms of a given element are identical in mass and properties, atoms of
different elements are different
3)
Compounds are formed by a combination of two or more different kinds
of atoms. Atoms combine in fixed whole-number ratios to make these
compounds.
4)
A chemical reaction is a combination, separation, or rearrangement of atoms.
In 1897, the British physicist J. J. Thompson carried out a series of
experiments that showed that atoms were not indivisible particles.
Thompson used a cathode ray tube, which is shown below, to produce a
beam that was deflected by a magnet towards the positive pole of the
magnet.
From his experiments, Thompson concluded that
1. A cathode ray consists of a beam of negatively charged particles (electrons).
2. Electrons are found in all atoms.
3. Since the atom is neutral, the total charges of the electrons must be equal to the total positive
charge of the rest of the atom.
Thomson was able to calculate the huge ratio of the electron's charge qe, to its mass, me. In his
theory (plum pudding or raisin pie model), the atom looks more like a positively charged ball
that has electrons embedded in it.
In 1911, Ernest Rutherford tested this atomic theory, through the famous gold-foil experiment.
Alpha particles (helium atoms that have lost two electrons and have a double positive
charge) were directed in a narrow beam at a very thin sheet of gold foil.
They expected all the alpha particles to pass straight through the gold, with little or no deflection.
To everyone's surprise, a small fraction of the alpha particles deflected, or bounced off the gold
foil at very large angles, even straight backwards. Based on the experimental results, Rutherford
suggested a new theory of the atom. He proposed that:
1. The atom is mostly empty space
2. Almost all the mass, and all the positive charges are concentrated in a small region at the center
of the atom. He called this region the nucleus.
3. The nucleus is so dense and tiny in size, compared to the rest of the atom.
4. Electrons must be orbiting somewhere in the space around the nucleus.
In 1913, the Danish physicist Niels Bohr proposed yet another modification to the theory of
atomic structure based on a phenomenon called line spectra.
When matter is energized (heated for example), it gives off light. When normal
white light, such as that from the sun, is passed through a prism, the light
separates into a continuous spectrum of colors:
Figure 1: Continuous (white light) spectra
When elements were excited by heat or electricity, they gave off distinct colors rather than
white light. This phenomenon is seen in modern-day neon lights, tubes filled with gaseous
elements (most commonly neon). When an electric current is passed through the gas, a distinct
color is given off by the element. When light from an excited element is passed through a
prism, only specific lines (or wavelengths) of light can be seen. These lines of light are called
line spectra. Bohr worked with the hydrogen spectrum which gave a lavender glow, and when
diffracted it produces the emission line spectrum below.
Figure 2: Hydrogen line spectra
Each element has its own distinct line spectra. For example:
Figure 3: Helium line spectra
Figure 4: Neon line spectra
To Bohr, the line spectra phenomenon showed that electrons can emit energy in very precise
quantities (he described the energy emitted as quantized). Bohr hypothesized that electrons
occupy specific energy levels. When an atom is excited, (such as during heating), electrons can
jump to higher levels. When the electrons fall back to lower energy levels, precise quanta of
energy are released as specific wavelengths (lines) of light.
Under Bohr's theory (also known as the planetary model):
1.
An electron's energy levels (also called electron shells) can be imagined as concentric circles
around the nucleus.
2. If the atom is unexcited, (in the ground state), the electron of hydrogen occupies the lowest
energy level possible (the electron shell closest to the nucleus).
3. Energy of an electron is constant in one of its allowed orbits. As long as an electron remains in
its orbit, it neither absorbs nor radiates energy.
4. When an electron is excited, the electron will absorb specific quanta of energy, "jump" to a
higher energy level, but after a very short time, it will spontaneously "fall" back to a lower
energy level, giving off a specific quantum of light energy.
5. The electron could only "jump" and "fall" to precise energy levels, thus emitting a specific
number of radiations (specific wavelengths) and hence produce a line spectrum of sharp
colored lines.
6. The electron cannot exist in between orbits.
Although the Bohr model adequately explained how atomic spectra worked, it failed to explain
certain aspects of multi-electron atoms.
In 1924, the French physicist Louis de Broglie suggested that, like light, electrons could act as
both particles and waves. De Broglie's hypothesis was soon confirmed in experiments. Another
question quickly followed: If an electron traveled as a wave, could you locate the precise
position of the electron within the wave? A German physicist, Werner Heisenberg, answered
no in what he called the uncertainty principle:
We can never know both the energy and position of an electron in an atom. Therefore,
Heisenberg said that we shouldn't view electrons as moving in well-defined orbits around the
nucleus.
With de Broglie's hypothesis and Heisenberg's uncertainty principle in mind, an Austrian
physicist named Erwin Schrodinger derived a set of equations or wave functions in 1926 for
electrons. According to Schrodinger, you could describe only the probability of where an
electron could be. The distributions of these probabilities formed regions of space around the
nucleus that were called orbitals. Orbitals could be described as electron density clouds. The
densest area of the cloud is where you have the greatest probability of finding the electron and
the least dense area is where you have the lowest probability of finding the electron.
Below is a visual summary of the evolution of the atomic models through history