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Chemical experiments from the 17th to the 19th centuries revealed the behavior of elements and chemical reactions. Scientists thought that atoms were indivisible; however, they should have realized, based on Mendeleev’s organization of the elements in columns with similar characteristics that there was a more fundamental substance in matter that was the basis of the arrangement of the elements in columns. The 19th century paradigm of indivisible atoms was upended in 1897 when J. J. Thomson bent the path of particles in a cathode ray tube with a magnet. Thomson stated, “I can see no escape from the conclusion that they [cathode rays] are charges of negative electricity carried by particles of matter.” Thomson developed a theory of atomic structure called the plum pudding model with electrons embedded in the larger atom, similar to raisins in plum pudding.
In 1898, Marie and Pierre Curie discovered polonium and radium, which produced alpha rays (helium nuclei with two protons and two neutrons). In 1911, Ernest Rutherford passed alpha radiation through thin gold sheets and measured the deflection of the alpha particles. Two undergraduate students, Hans Geiger and Frank Marsden, measured alpha particle deflections and found that one alpha particle out of 8,000 deflected by more than 900 and the rest of the alpha particles passed straight through the gold sheets.
Figure 2-8. Helium atom ground state. Credit: Yzmo. Used here per CC BY-SA 3.0.
Based on the experiments, Rutherford concluded that most of the atom is empty space, and that a very small positive nucleus in the center of the atom, 8,000 times smaller than the atom, holds the majority of the mass of the atom. Rutherford stated, “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.” [1] This experiment disproved the plum pudding atom model. The positive charges were in a dense mass in the center of the atom. Rutherford conceived of protons forming the nucleus of the atom, but he did not know realized that there were also neutrally charged particles in the nucleus called neutrons (Figure 2-8).
In the nucleus, there is a balance of the strong nuclear force keeping the nucleus together, and the weak nuclear and electromagnetic forces pushing it apart. The distribution of elements in the universe is a function of this balance of forces.
Electrons have the opposite electrical charge from the nucleus. Because opposites attract, scientists could not understand why negatively charged electrons did not fall into the positively charged nucleus. When there is a scientific problem such as this, it is actually an opportunity for discovery. Niels Bohr suggested in 1913 that electrons could only be located certain orbitals (distances) around the nucleus and thus could not crash into the nucleus. In 1919, Irving Langmuir proved this concept when he discovered that electrons orbit the nucleus in clusters within electron shells (Figure 2-9). This arrangement of electrons is the basis for the great diversity of elements and for the formation of bonds between elements in molecules.
Figure 2-9. Electron orbitals (clouds) where electrons exist at various energy levels. Credit: Geek3. Used here per CC BY-SA 4.0.
Bohr discovered that electrons jump instantaneously from one orbital to another when they absorb energy from a photon. Bohr also proposed that when electrons are within an orbital, they are completely stable and emit no energy; thus, they do not fall into the nucleus because of a loss of energy. Scientists define the electron orbitals as the locations at which there is a 95% probability of finding the electron. Because electrons only jump from one orbital to the next, which is a higher energy orbital, and never exist at intermediate energies, scientists had to develop a new branch of physics to describe their behavior called quantum mechanics where quantum refers to the fact that the electrons can only exist at discrete (certain) values.
Figure 2-10. A wave function for two identical fermions (electrons). Credit: Timothy Rias. Used here per CC BY 3.0.
As with photons, electrons act as waves (Figure 2-10) and as particles. In fact, de Broglie found that even atoms and all of matter behave like waves. The wavelength of matter is related to the momentum of matter. Electrons form standing waves which are similar to the string of a violin, where the waves can only exist at certain points in the string, just as electrons only exist at certain distances from the nucleus of the atom. The Stern-Gerlach experiments showed that particles also have angular momentum (spin). Schrodinger then derived extremely complex quantum mechanics equations for the behavior of the waves. These equations were not elegant, which should have been a sign that they could be improved. In 1928, Paul Dirac merged space and time into a four-dimensional equation according to Einstein's special theory of relativity (E = mc2) and then merged this equation with the principles of quantum mechanics. His equation was elegant and symmetrical and is shown below. It explained many relationships in electricity and magnetism. It also predicted the existence of antimatter.
There are several paradoxes in quantum mechanics. (1) The uncertainty principle is that the position and momentum of a particle cannot both be known. (2) Schrodinger's cat paradox is that systems can have different states if they are unobserved. This is expressed as “a cat can be both dead and alive as long as nobody is observing it.” (3) On the other hand, a particle cannot change its state while it is observed (Quantum Zeno effect). (4) Two particles with no observable contact can accomplish tasks together. (5) Spherical wave functions produce linear paths of particles. (6) Matter can act as both a wave and a particle. (6) A charged particle is influenced by an electromagnetic field at a point where the electromagnetic field strength is zero. All these strange behaviors are necessary to the function of atoms and particles. Particle physicists will never have a lack of phenomena to explore. Wouldn't Dalton have been surprised?
The last part of this section reviews some basic concepts in chemical bonding that will be useful later in this chapter and in future chapters. If you have already taken chemistry, then you can skip it. Electronegativity determines whether an atom will take an electron, give an electron, or share an electron in a chemical bond. There are two primary types of chemical bonds, covalent and ionic. Ionic bonds form between atoms of different electronegativity. For example, chlorine and sodium have extremely different electronegativities. Sodium gives away its outer electron and chlorine takes the electron. As solids, sodium and chlorine form ionic bonds in table salt. When dissolved in liquid they separate and become separate ions in water. The elements of life (hydrogen, carbon, oxygen, and nitrogen) form strong covalent bonds and share electrons because they have similar electronegativities. These strong covalent bonds are essential to life. Carbon, in particular, has four open electron positions in its outer shell and thus forms numerous bonds with other atoms. When atoms share electrons, the orbitals encompass parts of both atoms. The orbitals encompassing two or more atoms act as a bond between the atoms and form a molecule. The ways in which electrons and orbitals interact determines all of chemistry and the formation of molecules and solids. Molecules conform to the shape that has the lowest total energy in their chemical bonds.
The configurations in which bonds form enable the functions of molecules. For example, water has its hydrogen atoms grouped on one side of the oxygen atom, which gives it have a positive charge on one side and a negative charge on the other side of the molecule. This makes it a "polar" molecule and gives it many important characteristics such as cohesion, surface tension, and the capability of being a solvent. Proteins in cells have specific configurations that enable them to act like little machines in the cell. Some molecules can switch from one shape to another from the impact of a photon. This happens in the eye when photons cause chromophores to change shape (isomerization). This begins a series of shape changes in proteins that results in a signal sent to the brain.
The fundamental configurations of electron clouds and the behavior of electrons govern the activity of elements in chemical reactions. Similar configurations of outer electron shells in the same columns of the Period Table explain the similar behavior of different elements. The position in the Periodic Table determines electronegativity. All of the chemical properties or atoms and molecules on earth are dependent on the increasing and repeating structures of electrons in atoms. The tremendous chemical and biological diversity on Earth comes this building block of electron behavior and structures in atoms. Scientists estimate that there are between 1018 and 10200 possible chemicals in the universe.
References
References
[1] Jeremy Bernstein and David Cassidy, Hitler’s uranium club, the secret recordings at Farm Hall (New York: Copernicus Books, 2001), p. 5.
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Electron wavefunctions in hydrogen atom at different energy levels. Highest probability of finding electron in light areas. Credit: PoorLeno. Public domain.
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