In a compound, atoms are held together by an energy force called a chemical bond, which is somewhat analogous to a piece of elastic holding two balls together. Of course, the bond does not really consist of a material substance, but like a piece of elastic, it represents a certain amount of potential energy.
To carry the analogy further, when either a piece of stretched elastic or a chemical bond id broken, potential energy is converted to kinetic energy. A chemical bond if apparently an energy relationship between atoms.
The amount of energy in chemical bonds is variable and is dependent upon the number and kinds of atoms which are associated. As we mentioned previously, the number of electrons in the outermost shell or energy level of an atom does not exceed eight, and in atoms that have more than one shell, the presence of eight electrons in the outermost shell represent a stability.
As a general rule, we can predict that atoms will interact to form compounds when all participants are able to achieve stability by doing so. In essence, we are saying that all chemical reactions involve an exchange of energy.
This energy becomes kinetic as electrons interact in achieving stability, which means that bonds are either broken or formed, as the case may be. For our purposes, there are two types of atomic interactions, based on the manner in which electrons of one atom relate to those of another atom.
Let us suppose that an atom has a second shell containing eight electrons, and that a single electron occupies a third shell. It has a tendency to give up this single electron in achieving stability, and it will do so under certain conditions to any atom which will accept the electron.
In contrast, if we should find an atom with only seven electrons in its outermost shell, then it has a tendency to accept a single electron in achieving stability. If two such atoms are brought together, they make this exchange, and the result is an ionic or electrovalent bond.
A number of similar interacting pairs of atoms comprise an ionic or electrovalent compound. For example, the sodium atom (nNa23) has eleven electrons, two of which form the first shell, eight of which form the second shell, while the remaining electron occupies a third shell.
The chlorine atom (17C135) with seventeen electrons has seven of these in its third shell. The chlorine atom readily accepts the outermost electron of sodium- this electron actually transfers to the chlorine atom.
The resulting compound, sodium chloride, actually consists of two types of stable but electrically imbalanced atoms. An atom achieving structural stability through the loss or gain of electrons is called an icon.
Hence, sodium chloride consists of two types of ions, the sodium ions and the chloride ions. These are represented by the symbols Na+ and CI-, respectively, and the sodium chloride is represented by the formula Na’CL. In such reactions as we have described, the total number of positive charges carried by one ion equals the total number of negative charges carried by the other.
Since opposite charges attract each other we should expect that the positive and negative ions exert a mutual attraction. This force of attraction is termed the ionic or electrovalent bond.
Atoms do not always react in one-to-one ration in attaining stability. For example, consider the interaction between calcium (20Ca40) and chlorine in the formation of calcium chloride (Ca++2C1~). Let us remember that we can get a reaction between calcium and chlorine if all participating atoms achieve stability. This is accomplished quite readily in the interaction of calcium and chlorine.
In other words, the calcium atom has two electrons to donate in achieving stability and, in a manner of speaking; it does not “care” whether it donates both electrons to a single atom or to a pair of atoms.
To summarize ionic or electrovalent interactions, it is meaningful to say that atoms with more than four electrons in the outermost shell have a tendency to accept additional electrons, whereas those with fewer than four trend to give them up, thus presenting a satisfied shell to the out die.
From a purely physical standpoint, atoms which have interacted to form ionic bonds may be only loosely associated, especially if the compound they form is dissolved in some liquid such as water.
Under certain conditions, atoms may satisfy their outer most orbits by sharing electrons. In this case, an energy bond is formed. All atoms held together as unit by covalent bonds are called a molecule. A compound which is composed of similar molecules is called a covalent or molecular compound. Using chlorine again as an example, let us consider how the atoms of this element might interact to achieve stability.
Chlorine has seven electrons in its outer orbit, but eight are required for stability. Whenever a chlorine atom comes close to another chlorine atom, each “tries” to wrest an electron from the other. However, it is a drawn match, since each atom holds on to its electrons with equal tenacity. As a result, both atoms end up sharing a pair of electrons which sometimes circle the nucleus of one atom, and sometimes that of the other.
By virtue of this arrangement, both atoms achieve stability, and comprise a molecule, which we symbolize as Cl2. Thus, chlorine does not exist naturally in the form of individual atoms but, rather, as molecular chlorine.
In similar fashion, many elements exist as molecules; for example, two hydrogen atoms interact to form a molecule of hydrogen (H2). Now let us consider the carbon atom, which has four electrons in its outer shell.
We might question whether it tends to give them up take on four more in achieving stability. Actually, it usually does neither, but participates in the formation of molecules through sharing of its electrons.
Because of its outer shell configuration, carbon is a most versatile atom, and can form an almost infinite number of different arrangements with other atoms, for example, carbon reacts in the presence of hydrogen to form methane.
In this case, each hydrogen atom shares a pair of electrons with the carbon atom. A number of ways may be used to present the relationship of shared electrons.
The mostly used method of representing molecular relationships is the second one shown above. Each line represents a pair of shared electrons, which form a chemical bond between the two atoms involved.
Ionic Dissociation and Electrolytes
Water has been termed the universal solvent because it dissolves more substances found in nature than does any other known liquid. In living systems, practically all chemical reactions take place in an aqueous medium.
For this reason, it is important that we understand some of the fundamental types of interactions between water and the compounds that go into aqueous solution. Let us consider what happens when we mix a spoonful of sodium chloride with a point of water.
Not only does sodium chloride dissolve in water, but it also is associates, thus releasing Na+ and CI- ions into the solution. Since there are equal numbers of positively and negatively charged ions, the solution remains electrically neutral.
However, the very fact that there are ions of opposite charge in the water means that an electric current can pass through the solution. Thus compounds that ionize in water are known as electrolytes; and those which do not ionize in water are called none electrolytes.
Fundamentally, electrolytes fall into three groups known as acids, bases, and salts. There are more technical definitions of acids, bases, and salts than the ones we have given and the chemist is often obliged to formulate such definitions.
However, these will serve us as beginning concepts, and the terms are used according to these definitions. The strength of an electrolyte is determined by the degree to which it ionizes. For example, hydrochloric acid is a strong acid because it dissociates almost completely into hydrogen and chloride ions.
In contrast, an organic acid such as acetic acid is considered to be relatively weak because only a small percentage of its molecules dissociate into ion pairs. Similarly, there are strong and weak bases and salts.