2: Structure of Matter and Principles of Adhesion

Structure of Matter and Principles of Adhesion

Key Terms

Around 460 B.C., the Greek philosopher Democritus proposed that all matter was composed of indivisible particles called átomos (á = “un”; temno = “to cut”; meaning “uncuttable”), which is the origin of the name atoms. We know that an atom consists of a nucleus surrounded by a cloud of negatively charged electrons, as depicted in the electron cloud model of an atom (Figure 2-1). Except for the hydrogen atom, where there are no neutrons, the nucleus contains a mix of positively charged protons and electrically neutral neutrons. The electrons of an atom exist in different clouds at the various energy levels. An atom becomes a negative ion when it gains electron(s) or a positive ion when it loses electron(s).

Two or more atoms can form an electrically neutral entity called a molecule. Attraction between atoms and between molecules result in materials we can see and touch. Consider water as an example. Chemically, the basic unit of water is a molecule made of two hydrogen atoms and one oxygen atom. If each molecule attains a kinetic energy that is higher than the attraction between these molecules, they appear in the vapor form. As the surrounding temperature decreases, the level of kinetic energy within individual molecules decreases and the attraction between them becomes more prominent, so that they condense to a liquid form. Further cooling yields a solid called ice, where the kinetic energy is so low that the molecules are immobilized by the attraction between them.

The transformation between vapor, liquid, and solid is called the change of state. A change from the solid to the liquid state will require additional energy—kinetic energy—to break loose from the force of attraction. This additional energy is called the latent heat of fusion. The temperature at which this change occurs is known as the melting temperature or fusion temperature. When water boils, energy is needed to transform the liquid to vapor, and this quantity of energy is known as the heat of vaporization. It is possible for some solids to change directly to a vapor by a process called sublimation as seen in dry ice; this, however, has no practical importance as far as dental materials are concerned.

Interatomic Bonds

The preceding brief focus on change of state raises a question concerning the types of forces holding these atoms and molecules together. The electronic structure of an atom is relatively stable if it has eight electrons in its outer valence shell, as noble gases do, except for helium, which has only two electrons. Other atoms must lose, acquire, or share electrons with yet other atoms to achieve a stable configuration—that is, eight electrons in the outer shell. These processes produce strong or primary bonds between atoms. The bonding of atoms within a molecule also creates new but much weaker forces holding the molecules together. These are often called secondary bonds.

Primary Bonds

The formation of primary bonds depends on the atomic structures and their tendency to assume a stable configuration. The strength of these bonds and their ability to reform after breakage determine the physical properties of a material. Primary atomic bonds (Figure 2-2), also called chemical bonds, may be of three different types: (1) ionic, (2) covalent, and (3) metallic.

Ionic Bonds

The classic example of ionic bonding is the bond between the Na+ and Cl of sodium chloride (Figure 2-2, A). Because the sodium atom contains one valence electron in its outer shell and the chlorine atom has seven electrons in its outer shell, the transfer of the sodium valence electron to the chlorine atom results in the stable compound Na+Cl. In dentistry, ionic bonding exists in some dental materials, such as in gypsum structures and phosphate-based cements.

Covalent Bonds

In many chemical compounds, two valence electrons are shared by adjacent atoms (Figure 2-2, B). By virtue of sharing electrons, the two atoms are held together by covalent bonds to form a molecule that is sufficiently stable, and electrically neutral in a definite arrangement. The hydrogen molecule, H2, exemplifies covalent bonding. The single valence electron in each hydrogen atom is shared with that of the other combining atom, and the valence shells become stable. Covalent bonding occurs in many organic compounds, such as in dental resins, where they link to form the backbone structure of hydrocarbon chains (Chapter 6).

Metallic Bonds

The third type of primary atomic interaction is the metallic bond (Figure 2-2, C). The outer shell valence electrons can be removed easily from metallic atoms and form positive ions. The free valence electrons can move about in the metal space lattice (Chapter 5) to form what is sometimes described as an electron “cloud” or “gas.” The electrostatic attraction between the electron cloud and the positive ions in the lattice provides the force that bonds the metal atoms together as a solid.

The free electrons give the metal its characteristically high thermal and electrical conductivity. These electrons absorb light energy, so that all metals are opaque to transmitted light. The metallic bonds are also responsible for the ability of metals to deform plastically. The free electrons can move through the lattice, whereas their plastic deformability is associated with slip along crystal planes. During slip deformation, electrons easily regroup to retain the cohesive nature of the metal.

Combination of Primary Bonds

Although we can describe the three primary bonds separately, it is also possible to find more than one type of primary bond existing in one material. Consider calcium sulfate (CaSO4), the main ingredient of gypsum products (Chapter 9), as an example (Figure 2-3). In the sulfate ion (SO42−) the sulfur and oxygen atoms are held together covalently but they are short of two electrons. Calcium has two electrons in the outer orbit, which are easily removed and transferred to the SO4. The result is a Ca2+ ion with attraction for an SO42− ion.

Secondary Bonds

In contrast with primary bonds, secondary bonds do not share electrons. Instead, charge variations among atomic groups of the molecule induce dipole forces that attract adjacent molecules or parts of a large molecule.

van der Waals Forces

These van der Waals forces of attraction arise from dipole attractions (Figure 2-4). In the case of polar molecules, dipoles are induced by an unequal sharing of electrons (Figure 2-4, A). In the case of nonpolar molecules, random movement of electrons within the molecule creates fluctuating dipoles (Figure 2-4, B). Dipoles generated within these molecules will attract other similar dipoles. Such interatomic forces are quite weak compared with the primary bonds.

Hydrogen Bond

The hydrogen bond is a special case of dipole attraction of polar compounds. It can be understood by studying a water molecule (Figure 2-5). Attached to the oxygen atom are two hydrogen atoms. These bonds are covalent. As a consequence, the protons of the hydrogen atoms pointing away from the oxygen atom are not shielded efficiently by the electrons. They become positively charged. On the opposite side of the water molecule, the electrons that fill the outer shell of the oxygen provide a negative charge. The positive hydrogen nucleus is attracted to the unshared electrons of neighboring water molecules. This type of bond is called a hydrogen bridge. Polarity of this nature is important in accounting for the intermolecular reactions in many organic compounds—for example, the sorption of water by synthetic dental resins.

Atomic Arrangement

All materials we use consist of trillions of atoms. As described earlier, they are attracted to each other and retain a particular physical appearance. The question is in which configuration they are held together. In 1665, Robert Hooke (1635–1703) explained crystal shapes in terms of the packing of their component parts, like stacking musket balls in piles. This is an exact model of the atomic structure of many familiar metals, with each ball representing an atom.

In the solid state, atoms combine in a manner that ensures minimal internal energy. For example, sodium and chlorine share one electron at the atomic scale. In the solid state, like grains of salt, they do not exist in individual pairs; in fact, each sodium ion is attracted to six chlorine ions and vice versa (Figure 2-6). They form a regularly spaced configuration (long-range repetitive space lattice) known as a crystal. A space lattice can be defined as any arrangement of atoms in space in which every atom is situated similarly to every other atom.

There are structures where regularly spaced configurations do not occur in the solid state. For example, the molecules of some of the waxes used by a dentist or laboratory technician are distributed at random when solidified. This noncrystalline formation is also known as an amorphous structure.

Crystalline Structure

There are 14 possible lattice types. The type of space lattice is defined by the length of each of three unit cell edges (called the axes) and the angles between the edges. The simplest and most regular lattice is a cubic, as shown in Figure 2-7, A; it is characterized by axes that are all of equal length and meet at 90-degree angles, representing the smallest repetitive volume of a crystal, which is called a unit cell. Each sphere represents the positions of the atoms. Their positions are located at the points of intersection of three planes, each plane (surface of the cube) being perpendicular to the other two planes. These planes are often referred to as crystal planes. However, the simple cubic arrangement shown in Figure 2-7, A, is hypothetical, as it leaves enough space to fit additional atoms per unit cell. Most crystalline lattices of atoms also contain sites of missing atoms. Each missing atom site is called a vacancy.

Most metals used in dentistry belong to the cubic system. For example, iron at room temperature has an atom at each corner of the cube and another atom at the body center of the cube (Figure 2-7, B). This crystal form is called a body-centered cubic cell. Copper, on the other hand, has additional atoms at the center of each face of the unit cell but none at the center of the cube. This form is called a face-centered cubic cell (Figure 2-7, C).

Other types of space lattices of dental interest are shown in Figure 2-8. The hexagonal close-packed arrangement (Figure 2-8, G) observed in titanium, zinc, and zirconium has become an important crystalline structure in dentistry. Note that each unit cell consists of three layers of atoms.

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Jan 1, 2015 | Posted by in Dental Materials | Comments Off on 2: Structure of Matter and Principles of Adhesion

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