The History of Chemical Symbols
The one or two letter shorthand used to represent chemical elements is a familiar feature in modern science. The tradition of using symbols to represent elements is quite ancient. Long before those interested in studying the composition and behavior of matter were known as chemists, mystical practitioners of alchemy devised symbols often to obfuscate their experimentation, and to cloak their work in secrecy. Their coded imagery drew inspiration from astrology, as well as ancient writing systems like the hieroglyphs. Alchemists linked certain metals with celestial bodies to describe their behavior, such as the connection between the rapidly moving planet Mercury and the metallic liquid quicksilver.
As chemistry became an experimental science and new methods produced scores of newly discovered elements, the need for a shorthand technique to describe chemical changes became apparent. One of the first chemists to attempt to introduce a symbolic system for identifying the elements was John Dalton, known for his relative mass scale of the atomic weights. His symbols, introduced in 1808 in his “New System of Chemical Philosophy”, consisted mainly of circles, some with inscribed alphabetic letters and others with dots or lines within the circles. Compounds were written as combinations of circles representing the constituent atoms. His system did not lend itself to ready memorization and did not catch on with his contemporaries.
Our modern method of using one or two letter shorthand for the elements was devised in 1813 by Jöns Jakob Berzelius, citing the ease of implementation, particularly for typesetters. Due to the common employment of Latin terminology in scientific communication, Berzelius suggested using the first or first two letters of the element’s Latin name as the symbol for that atom. In the case of confusion or duplication of the letters, exceptions includes the use of Hg (hydrargyrum for Mercury and plumbum for lead). Some modern modifications have been introduced for new elements, especially those named in honor of famous scientists. Berzelius is also responsible for the use of subscripts in a chemical formula to designate the ratio of atoms.
One of the primary causes of change in physical systems is the tendency toward minimum potential energy. Objects roll downhill, objects above the Earth fall, stretched rubber bands contract, objects with like charges separate, and objects with unlike charges move together. All of these changes involve a decrease in potential energy.
In many situations, the potential energy of a system increases somewhat, at first, in order to achieve a position from which the potential energy will significantly decrease. A siphon is an example of this. In a siphon, water will run up the hose as long as the final position of the water is lower in potential energy than the original position.
In ionic bonding, the metallic atom must lose one or more electrons in order to form a bond. This loss of electrons by metallic atoms requires an input in energy. The necessary ionization energy (energy to remove an electron) must be provided to form a cation (positive ion). Suppose we use a sodium atom as an example.
An input of energy ionizes the sodium atom to a sodium ion, . In the presence of chlorine atoms, the electron can then add to a chlorine atom to form a negative chloride ion. In this process, energy is given off. The electron affinity of chlorine is . That means that adding an electron to a chlorine atom is a reduction in potential energy.
The sodium ions and the chloride ions have opposite charges, and therefore, are attracted to each other. When the ions move closer together, potential energy is again lowered. As the oppositely charged ions move together to form a crystal lattice, energy is given off. For each mole of sodium ions and chloride ions that move together to form a lattice, of energy are given off.
The overall process of removing an electron from a sodium atom (), adding the electron to a chlorine atom (), and the ions moving together in a lattice structure () has a net energy output of ; therefore, like a siphon, the process occurs because it has a net lowering of potential energy.
In the case of covalent bonding, there is no electron losing, gaining, or transferring. In covalent bonds, the bonding electrons are shared. The potential energy lowering that occurs in covalent bonding can be represented by showing the relationship between the potential energy of the system and the distance between the nuclei of the bonding atoms.
When individual non-metallic atoms are completely separated, there is an attraction (electron affinity) between the electrons of each atom and the protons in the nucleus of the other atom. This attraction and the distance between the atoms cause potential energy to exist. As the atoms move closer together, the potential energy of the system decreases because the distance between the attracting objects is becoming less.
As the distance between the two nuclei decreases, the potential energy becomes less and less. Since these are non-metallic atoms, their electron outer energy levels are not full and this allows the electron clouds of the two atoms to overlap, and the atoms may continue to move closer together. When the atoms reach position B in the diagram, the potential energy of the system is at its lowest possible for these two atoms. If the nuclei continue to move closer together, the potential energy increases dramatically due to the repulsion of the positively charged nuclei. In position A, the nuclei are too close together so potential energy is high. In position C, the atoms are too far apart so the potential energy is high. In position B, the atoms are the proper distance apart for lowest potential energy. Since these two atoms reach lowest potential energy in a position where their electron clouds are overlapped, they will be bonded. The distance between the nuclei in position B is known as the bond length. If the atoms attempt to move either closer together or farther apart, the potential energy increases. The tendency toward minimum potential energy causes these atoms to remain in the bonded position.