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Everything we see, everything we touch and taste, most everything in our natural world is composed of matter. Matter surrounds us and interacts with us, but what is it really? Matter is anything that occupies space and has mass. Each bit of matter has a characteristic density, hardness, taste and color. Matter can undergo transformations including melting, boiling, and changes in composition. But, what is it made of and how does that determine the above properties? To begin, consider this analogy.

The writing on this page can be broken down into its basic parts. Each sentence is composed of words. Each word contains 26 letters of the English alphabet, and each letter is simply a combination of curves, lines, and dots. Building these curves, lines, and dots into words and sentences can convey ideas such as the nature of matter. Similarly, matter is composed of atoms, and each atom can be broken down into subatomic particles. The way these subatomic particles combine determines the properties and behaviors of the atoms and the matter they comprise.

Surprisingly, in this vast cosmos we live in, there are only 109 natural atoms (also called elements). All of them are combinations of three subatomic particles, protons, neutrons, and electrons. Protons are positively charged particles that have a tiny mass. Neutrons are of similar mass but have no charge. Electrons are particles of very little mass with a charge equal to and opposite to that of the proton. It is convenient when talking about such a small charge and mass to use relative units, and the weight of one proton is defined as one atomic unit (also be referred to as a Dalton). For a similar reason, the charge on a proton is positive and given the value of +1. The charge on an electron is negative and given the value -1.

Protons and neutrons combine to form a nucleus in the center of the atom. The positive charge of the nucleus attracts and holds electrons. The electrons move very fast (near the speed of light) and, as best as we can measure, form a cloud of charge around the nucleus. Electrons want to be around as much positive charge as possible, and protons want to collect as much negative charge as possible. The ideal state for the atom, its most stable state, is to balance these two forces such that the atom has an equal amount of positive and negative charge. As an example, here is the construction of two of the lighter elements, hydrogen and helium. Hydrogen has one proton (one positive charge), no neutrons, and one electron (one negative charge). Helium has two protons, two neutrons, and two electrons (Figure 2.1).

The number of protons in an atom increases its weight and changes its fundamental properties, differentiating one element from the next. The number of protons in an atom's nucleus delineates its properties, and scientists coined the term atomic number to express this idea. An atom containing 6 protons is a carbon atom no matter how many neutrons or electrons it contains.

The addition of a neutron adds to an atom's weight and can affect its stability, but it does not change its properties. In fact, atoms of carbon containing seven neutrons (carbon 13) and eight neutrons (carbon 14) exist in nature. Carbon 12, Carbon 13, and Carbon 14 are three isotopes of carbon.

Bonding of atoms

The electron count of an atom affects its reactivity. Atoms want to be "full" of electrons. If they are not, they will try to borrow them from other atoms. As an example, consider hydrogen and helium. Hydrogen is a very reactive molecule, always searching for an extra electron, while helium having two electrons is "full" and not very reactive.

So what determines whether an atom is full or not? The way electrons associate with the nucleus. The discipline of quantum mechanics elucidated many of the rules for this association. When electrons come in association with the nucleus of an atom, they want to house themselves in very precisely defined spaces. These acceptable spaces are called orbitals. Each orbital defines a region where the electron is likely to be found and contains two electrons when full. Atoms have discrete energy levels or shells that contain these orbitals. Electrons must occupy orbitals in these energy states and are not allowed to exist outside the shells. A way you can imagine this is as the difference between sitting on an inclined plane vs. a set of steps. When on a plane, you can have any amount of potential energy, but you can only have discrete amounts with a set of steps. As one moves away from the nucleus, these shells require higher energy electrons to fill the orbitals. Atoms prefer to be in the lowest energy level possible; thus, electrons will fill the open orbitals in the lowest energy shell first, then move to the next shell, and so on.

As an analogy, orbitals can be thought of as houses where two electrons can live. These houses are organized into subdivisions (the shells) encircling a lake (the nucleus) at the bottom of a valley. Paths away from the hill have large steps instead of a simple incline. The most desirable location is on the lakefront, and electrons fill that house first. Further sets of electrons will fill the closest available house to the lake, but they must jump up a step, a quantum of energy to do so. More energy is required to live in the houses further from the lake because you have to walk up the steps to get there (Figure 2.2).

The rules for these associations determine the number and types of orbitals present in each shell. The first shell, the K shell, contains one s orbital that is spherical. The next shell, the L shell, contains four orbitals, one s orbital, and three p orbitals. The M shell contains 1 s, 3 p, and 5 d orbitals. Higher shells will contain more orbitals, but for the most part, atoms important in biological systems do not fill beyond the M shell (Figure 2.3).

Let's now go back to hydrogen and helium. Hydrogen has one electron that fills the s orbital of the K shell. There is room for one more electron in the orbital, and hydrogen seeks it out. In its elemental form, each hydrogen atom will pair up with a second hydrogen atom, forming H₂. Helium contains two electrons that fill the s orbital of the K shell. The K shell is full, and helium does not easily add more electrons. Elements with completely full shells (helium, argon, neon, and krypton, for example) are very non-reactive

In contrast to helium, most atoms have partially filled outer shells and are constantly searching for ways to fill them. This tendency makes them reactive with other atoms, and that is why they associate and build larger structures called molecules. Atoms fill these orbitals by sharing electrons between different nuclei and overlapping their orbitals. The number of protons and electrons in an atom dictates how much sharing each atom is willing to do and dictates its behavior in chemical reactions. Table 2.2 lists some common elements in organisms and their electron-pairing behavior.

All sorts of combinations between these molecules are possible. Carbon can combine with four hydrogens to form methane. Nitrogen can combine with another nitrogen molecule to form nitrogen gas. Pure oxygen can share electrons with another oxygen atom, forming O₂. Two hydrogens can share their electrons with an oxygen atom forming water. Figure 2.4 shows some simple molecules.

The term covalent bond is used to describe this type of electron sharing. Covalent bonds are stable and bind molecules together rather tightly. You can view this electron sharing as a combination of the individual shared electrons forming a distributed electron cloud around the two molecules.

Covalent bonds are not the only ways to interact

As we have learned above, covalent bonds connect atoms forming the monomers and polymers of the cell, but there are other forces at work. Three of the more essential bonding forces are ionic interactions, hydrogen bonds, and hydrophobic interactions, as shown in Figure 2.5.

Ionic interactions are the attraction of opposite charges for one another. Negatively charged groups are attracted to positively charged groups.

Hydrogen bonds involve hydrogen, which is not very good at holding onto electrons when covalently bonded with certain other atoms (usually oxygen or nitrogen). It will not get its fair share of time with the electrons and will be attracted to high electron densities around other nearby atoms. The hydrogen atoms are attracted to the extra electrons and tend to stay in the vicinity of the electron density. Not surprisingly, oxygen and nitrogen atoms have these high electron densities. This is not a covalent bond and hydrogen bonds tend to be weak interactions, but large numbers of them can add significantly to the stability of a protein or a DNA double helix.

If you want to see hydrogen bonds in action, fill a glass of water to overflowing. Note how the water clings to itself, even though it is above the height of the glass. Hydrogen bonds between the water molecules cause them to stick together. Now add a few drops of detergent. The detergent will disrupt the bonding between the water molecules. Don't do this unless you have a towel nearby!

Hydrophobic interactions consist of the congregation of non-polar groups together. The presence of a hydrophobic (water-hating) molecule in a water solution causes an undesirable ordering of the water around it. Water hates this and will minimize this interaction by pushing the hydrophobic groups together and away from the water. Think back to the behavior of oil in water. When you place some oil in a bottle of water, the oil tends to separate from the water and form one large bleb (that is a scientific term, really, look it up). If you shake the bottle, the oil disperses, but in a little while, it all congregates again. This behavior is hydrophobic interaction at work. The same process is functioning at the molecular level. Hydrophobic groups in biological systems consist of alkyl chains (see later in the chapter for a definition of this term). These groups are pushed together and hide away from the water in the environment. This grouping can lead to the formation of large structures spontaneously. Hydrophobic interactions largely drive membrane formation and protein folding.

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