Outline
Interaction Between Two Atoms
Types of Bonding
Atomic Arrangement
Metals
Ceramics
Polymers
Composite Materials
Adhesion and Bonding
Bonding to Tooth Structure
Key Terms
Adhesion An attraction between two contacting surfaces promoted by the interfacial force of attraction between the molecules or atoms of two different species; adhesion may occur as chemical adhesion, mechanical adhesion (structural interlocking), or a combination of both.
Adhesive and adherend Substance that promotes adhesion of one substance or material to another. An adherend is a material substrate bonded to another material by means of an adhesive.
Alloy and alloy system An alloy is a crystalline solid with metallic properties that is composed of two or more elements, at least one of which is a metal, and all of which are mutually soluble in the molten state. Alloy systems are all possible alloyed combinations of two or more elements, at least one of which is a metal. For example, the binary gold-silver system includes all possible alloys of gold and silver, varying from 100% gold and 0% silver to 100% silver and 0% gold.
Bonding The action of joining objects or particles together by means of adhesive or force of attraction.
Ceramic Solid-phase compounds of metallic and nonmetallic elements.
Cohesion Bonding between molecules or atoms of the same species.
Composite A material made of two or more constituent components with significantly different physical or chemical properties that, when combined, produce a material with characteristics different from the individual components.
Contact angle Angle of intersection between a liquid and a surface of a solid that is measured from the solid surface through the liquid to the liquid/vapor tangent line originating at the terminus of the liquid/solid interface; used as a measure of wettability, whereby no wetting occurs at a contact angle of 180°, and complete wetting occurs at an angle of 0°.
Copolymer Polymer made of two or more types of monomers. A random copolymer is when there is no sequential order of monomer types along the polymer chain. A block copolymer is when groups of each type of monomer appear in the same polymer chain. A graft or branched copolymer is when a sequence of one type of mer unit is grafted onto the backbone of a second type of monomer to form a branched structure.
Crosslink A difunctional or multifunctional monomer that forms a link between two growing polymer chains during polymerization that results in a three-dimensional interconnected polymer network.
Curing Chemical reaction in which low-molecular-weight monomers are converted into higher-molecular-weight materials through polymerization to attain desired properties.
Elastic recovery The shape of objects changes under an applied force. When the force is removed, the object will recover to its original shape instantly if it is an elastic solid, or the recovery process occurs over a period of time if it is a viscoelastic material. The greater the viscous nature of a material, the less complete the recovery.
Glass-transition temperature (Tg) (1) The temperature above which a sharp increase in the thermal expansion coefficient occurs, indicating increased molecular mobility. (2) The temperature at which macromolecule molecular motion begins to force the polymer chains apart. Thus polymeric materials soften when heated above this temperature.
Grain and grain boundary A grain is a single crystal in the microstructure of a metal. Metals and alloys are solids consisting of many individual grains, and the interface between adjacent grains is the grain boundary.
Metal (1) An element or alloy whose atoms readily lose electrons to form positively charged ions. (2) A metallic material composed of one or more elements that is opaque, ductile, relatively malleable, a good conductor of electricity, a good thermal conductor, and usually lustrous when light is reflected from its polished surface.
Microstructure Structural features of a surface, including grains, grain boundaries, phases, and defects, such as porosity, revealed by microscopic imaging of the chemically or electrolytically etched surface of a flat, polished specimen.
Micromechanical bonding Retention associated with an adhesive penetrating a roughened adherend surface.
Monomer and mer A monomer is a chemical compound that is capable of reacting to form a polymer; each monomer becomes a me r of a polymer.
Phase and phase diagram (constitution diagram) A graph of equilibrium phases and solubility limits for an alloy system as a function of composition and temperature.
Polymer Large molecules (a.k.a. macromolecules ) formed by the union of many simple repeating units (mers); the chemical process is known as polymerization.
Primary bond Bonding between atoms that include ionic, covalent, and metallic bonds.
Resin or synthetic resin Blend of monomers and/or macromolecules with other components, which form a material with a set of useful properties.
Solid solution (metallic) A solid crystalline phase containing two or more elements, at least one of which is a metal and whose atoms share the same crystal lattice.
Strain Magnitude of deformation (stretching, compression, or shear) occurring in response to an applied force.
Stress The perpendicularly directed force, exerted as pressure or tension, on a material that causes the object to deform (strain), measured in units of force per unit area.
Surface energy The excess energy of attraction that the surface of a material (liquid and solid) has compared with the bulk of the material because molecules or atoms at the surface are not surrounded by their fellows as those in the bulk. It can be treated as the work (energy) required to build an area of a particular surface and has the unit of mN/m 2 .
Surface tension The tendency of fluid surfaces to contract to the smallest surface area by an inward force caused by the imbalance of mutual attraction between molecules on the liquid surface. The very inward force also keeps the molecules of the surface constantly under tension. The cohesive force keeping molecules bonded on the surface is known as the surface tension of the liquid, with the unit of mN/m.
Thermoplastic polymer Macromolecule material made of linear and/or branched chains that softens when heated above the glass-transition temperature (T g ); it can then be molded to a new shape and cooled below the T g to retain the new configuration.
Thermoset polymer Polymeric material that retains a crosslinked structure during polymerization; the acquired structure results in a rigid polymer that, upon heating, does not soften enough to be molded to a new configuration.
van der Waals forces Short-range force of physical attraction that promotes adhesion between molecules of liquids or molecular crystals. Also known as secondary bond.
Wetting and wetting agent Wetting is the ability of a liquid to maintain contact with a solid surface; it reflects the intermolecular interactions when the two are brought in intimate contact. A wetting agent is a surface-active substance that can be applied to a solid substrate to reduce the surface tension of the liquid to be placed on the solid; the purpose is to promote wetting or adhesion.
All matter is composed of indivisible particles called atoms, the smallest size of particles having properties identical to those of bulk material. The commonly accepted structure of atoms consists of three subatomic particles, a negatively charged electron, a positively charged proton, and an electrically neutral neutron, as depicted in the electron cloud model of a nitrogen atom ( Figure 2-1 ). The nucleus contains a mix of protons and neutrons, except for the hydrogen atoms, where there are no neutrons. Neutrons and protons are held together by nuclear force that is stronger than the repulsion forces between protons. Nuclear force only acts at an extremely short distance (≅1 × 10 −15 m) and does not affect the electromagnetic attraction. The electrons of an atom exist in different clouds at the various energy levels (shells). An atom becomes a negative ion when it gains an electron(s) or a positive ion when it loses an electron(s). The number of electrons in the shells of the cloud dictates the property of the matter made of that group of atoms. These electrons are also called valence electrons.
When the state of material (vapor, liquid, and solid) changes, what happens between atoms or molecules that make up the material?
Two or more atoms can bond together and form an electrically neutral entity called a molecule. The attraction between atoms within a molecule is strong, whereas the attractions between molecules are weaker. These attractions 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 vapor form. Kinetic energy is what allows molecules of known mass to maintain a constant in motion. 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 liquid form. Further cooling yields a solid called ice, where the kinetic energy is so low that the molecules are immobilized by attraction forces among them.
The transformation between vapor, liquid, and solid is called change of state, which is an example of how the attraction between molecules and the environment affects the behavior of a material as we observe it. The focus of this chapter will be on how the atomic and molecular attractions shape the four classes of materials used in dentistry, namely, metals , ceramics , polymers, and composites , and affect their behavior in their respective applications. Finally, we will discuss how we can use these bonding mechanisms to our advantage.
Interaction Between Two Atoms
We can treat each atom as a discrete entity with definite boundaries and volume established by the electromagnetic fields of the electrons. When two atoms interact with each other in space, they feel an attraction force from the interaction of the opposite-charged particles of either atom and an expulsion force from the interaction of the similar-charged particles of either atom. At a large distance, the interaction between atoms is negligible ( Figure 2-2, A ). In Figure 2-2, A, the repulsion and attraction forces are plotted as a function of the interatomic spacing, which is the distance between the centers of the cores of the atoms. When the atoms come close to each other, both forces increase as the distance between the atoms decreases. Initially, the force of attraction increases faster than the force of repulsion. The repulsion force is a short-range force compared with the attractive force that becomes noticed at a closer range. Later, the force of repulsion increases much more than the force of attraction as the atoms get closer. The resultant force, which is the sum of attraction and repulsion forces, increases initially as attraction force when the two atoms get close. As the repulsion force increases, the resultant attractive force, after reaching a peak, starts to decline. It then reaches a state of equilibrium when the valence electron shells of the two atoms keep the atoms from getting closer. At the state of equilibrium, there is no net resultant force ( Figure 2-2, A ). In other words, the atoms have reached a stable state. Before the atoms reach equilibrium, we can treat the resultant force as the force needed to keep them apart at that distance. After equilibrium, the negative resultant force means that external force is needed to push the atoms closer. There is a great deal of repulsion to overcome. Depending on the number of valence electrons and species of atoms involved, there are three types of interaction between atoms when equilibrium is reached.
Bond Distance
Regardless of the type of structure in the solid state, there is a limiting factor that prevents the atoms or molecules from approaching each other too closely. The position at which both attraction and repulsion forces are equal in magnitude (but opposite in direction) is considered the equilibrium position of the atoms ( Figure 2-2, A ). The interatomic distance at equilibrium represents the sum of the radii of the two adjacent atoms.
Bond Energy
Because the conditions of equilibrium are usually described in terms of energy rather than interatomic forces, the relationships in Figure 2-2, A can be logically explained in terms of interatomic energy, which is also known as bond energy or potential energy . Energy is defined as the product of force and distance. Integration of the interatomic force (dashed line in Figure 2-2, A ) over the interatomic distance yields the interatomic energy (solid line in Figure 2-2, B ). In contrast with the resultant force, the bond energy can be treated as the energy needed to keep two atoms apart. Initially, the bond energy decreases gradually as two atoms come closer together. As the resultant attractive force passes the peak and begins to decline rapidly, the bond energy also decreases steeply ( Figure 2-2, B ). The bond energy reaches a minimum when the resultant force becomes zero. Thereafter, the energy increases rapidly because the resultant force becomes repulsive and increases rapidly with little change in interatomic distance. The minimum energy corresponds to the condition of equilibrium and defines the equilibrium interatomic distance.
Influence of Interatomic Bond
The atoms are in a constant state of vibration, and the average amplitude is dependent on the temperature; the higher the temperature, the greater is the amplitude and, consequently, the greater is the kinetic energy. At a certain temperature, the minimal energy required to maintain equilibrium is denoted by the bottom of the trough in Figure 2-2, B . As the temperature increases, the amplitude of the atomic (or molecular) vibration increases. It follows also that the mean interatomic distance increases ( Figure 2-3 ), as does the internal energy. The overall effect is the phenomenon known as thermal expansion.
As the temperature increases from T 0 to T 5 in Figure 2-3, A, the mean increase in interatomic distance is less with the deeper energy trough ( Figure 2-3, A ) than that in the shallower energy trough ( Figure 2-3, B ). This means that the linear coefficient of thermal expansion (α) of materials with similar atomic or molecular structures tends to be inversely proportional to the melting temperature. If the temperature continues to increase, the increase in interatomic distance will result in a change of state, where solid melts to a liquid, and liquid subsequently vaporizes to gas. For a solid with lower potential energy (i.e., a deeper trough depth) ( Figure 2-3, A ), greater amounts of kinetic energy are required to achieve melting and boiling, which corresponds to higher melting and boiling temperatures.
As shown in Figure 2-2, A, the net force on the atoms at the equilibrium distance is zero, but small displacements, as in an increased bond distance, result in rapidly increasing forces to maintain the equilibrium distance. The magnitude of this force needed to maintain the displacement is commonly referred to as the elastic modulus or stiffness of the material, and this property will be discussed in detail in Chapter 4, Stress-Strain Properties . The stiffness or elastic modulus of the material is proportional to the rate of change of the force with a change in displacement that is measured by the slope of the net-force curve drawn from the equilibrium ( Figure 2-2, B ). A greater slope of the force curve implies a narrower, deeper trough in the energy curve ( Figure 2-3, A ). Hence, a high melting point is usually accompanied by a greater stiffness.
The preceding principles represent generalities, and exceptions do occur. Nevertheless, they allow one to estimate the influence of temperature on the properties of most of the dental materials discussed in subsequent chapters.
Types of Bonding
The preceding discussion establishes the role of bonding in the change of state and interatomic bond energy of the thermal and mechanical properties. When the atoms are in the equilibrium state with the neighboring atoms, they establish different types of bonds according to the valence of their electrons. The electron 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 primary bonds between atoms. Intermolecular bonds, on the other hand, rely on dipole forces induced by nonuniform distribution of electrons within the molecule. These are often weaker and are 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 re-form after breakage determine the physical properties of a material. Primary atomic bonds ( Figure 2-4 ), 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 2-4, 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 and phosphate-based cements.
Covalent Bonds
A covalent bond is a chemical bond that involves the sharing of electron pairs between atoms. Figure 2-4, B shows that the fluorine molecule shares a pair of electrons. 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, H 2 , 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 dental resins, where they link to form the backbone structure of hydrocarbon chains.
Metallic Bonds
The third type of primary atomic interaction is the metallic bond ( Figure 2-4, C ). It is found in elements with one, two, or three valence electrons. 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 to form what is sometimes described as an electron “cloud” or “gas.” The electromagnetic 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 all metals are opaque to transmitted light. The metallic bonds are also responsible for the ability of plastic deformation of metals. Plastic deformation means a material can be reshaped by force without fracture. 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 (CaSO 4 ), the main ingredient of gypsum products ( Chapter 13, Gypsum Products ), as an example. In the sulfate ion (SO 4 2– ), the sulfur and oxygen atoms are held together covalently but are short two electrons. Calcium has two electrons in the outer orbit, which are easily removed and transferred to the SO 4 . The result is a Ca 2+ ion with attraction for an SO 4 2– ion.
Secondary Bonds
In contrast with primary bonds, secondary bonds do not share electrons between molecules. Instead, the asymmetrical distribution of electrons within each molecule induces dipole forces that attract molecules together.
Van der Waals Forces
Van der Waals forces of attraction arise from dipole attractions ( Figure 2-5 ). In the case of polar molecules, dipoles are induced by an unequal sharing of electrons ( Figure 2-5, A ). In the case of nonpolar molecules, the random movement of electrons within the molecule creates fluctuating dipoles ( Figure 2-5, 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 2-6 ). Attached to the oxygen atom are two hydrogen atoms. These bonds are covalent. Therefore 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 physical appearance. The question is in which configuration they are held together. In 1665, Robert Hooke 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 and the most efficient packing of atoms. 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 2-7 ). 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 the three unit cell edges ( a -, b -, and c -axes) and the angles (α, β, and γ) between the edges. The simplest and most regular lattice is a cubic, as shown in Figure 2-8, A; it is characterized by axes that are all equal length and meet at 90° 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 arrangements shown in Figure 2-8, A, is hypothetical because it leaves enough space to fit additional atoms per unit cell.
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 (BCC) 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 (FCC) cell ( Figure 2-7, C ). At first glance of Figure 2-7, B, one may think NaCl retains a simple cubic crystal if we consider that sodium and chlorine occupy the same lattice. That is incorrect. Sodium atoms occupy an FCC lattice, and chlorine atoms occupy their own FCC lattice. The two FCC lattices interpenetrate to form the final lattice. Other types of space lattices of dental interest are shown in Figure 2-9 . The hexagonal close-packed arrangement ( Figure 2-9, 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, and the base, being hexagonal, has a 120° edge.
All metallic-based dental materials are crystalline. Some pure ceramics, such as alumina and zirconia core materials, are entirely crystalline.
Noncrystalline Structure
Glass is a typical noncrystalline solid of SiO 2 because its atoms tend to be arranged in nonrepeating units ( Figure 2-10 ). The ordered arrangement of the glass is locally interspersed with a considerable number of disordered units. Because this arrangement is also typical of liquids, such solids are sometimes called supercooled liquids. Because of the complexity of the physical configuration of polymer chains, the molecules of resins are not favored to arrange in orderly, repeating patterns. Therefore polymer-based materials used in dentistry are usually noncrystalline.
The structural arrangements of the noncrystalline solids do not represent such low internal energies as do crystalline arrangements of the same atoms and molecules. Noncrystalline solids do not have a definite melting temperature but, rather, gradually soften as the temperature is raised. The temperature at which there is an abrupt increase in the thermal expansion coefficient, indicating increased molecular mobility, is called the glass-transition temperature (T g ) ( Figure 2-11 ); it is characteristic of the glassy structure. Below T g , the material loses its fluid characteristics and gains significant resistance to shear deformation. When set, synthetic dental resins are examples of materials that often have glassy structures, with T g greater than body temperature.
Many dental materials often consist of a noncrystalline glassy matrix and crystalline inclusions (filler phase). Crystalline inclusions provide desired properties, including color, opacity, increased thermal expansion coefficients, and in some dental ceramics, increased radiopacity. The filler phase of resin-based composites ( Chapter 5, Fillers ), on the other hand, can be crystalline, such as quartz particles or noncrystalline glass spheres.
Diffusion
When we place a drop of ink in a bowl of water, we observe the spread of the ink in the water. It will eventually disperse throughout the entire body of the water. This process is called diffusion. The same process also occurs within solid materials but at a substantially slower rate. An understanding of diffusion in a solid requires two new concepts.
First, the atoms in a space lattice, as previously described, are constantly in vibration about their centers. However, the atoms in the material do not all possess the same level of energy. Rather, there is a distribution of atoms with energy that varies from very low to high, with the average energy at equilibrium. If the energy possessed by an atom exceeds the bond energy, it can move to another position in the lattice.
Second, there are a finite number of missing atoms (called vacancies ) within a solid formed during solidification. A noncrystalline structure, because of short-range order, also contributes some space. Both conditions represent pathways through which diffusion can occur. Atoms change position in pure, single-element solids even under equilibrium conditions; this process is known as self-diffusion . As with any diffusion process, the atoms or molecules diffuse in the solid state to reach an equilibrium state. Just as ink disperses uniformly in water, a concentration of atoms in a solid metal can also be redistributed through the diffusion process.
Diffusion may also occur in the other direction to produce a concentration of atoms in a solid. For example, if the sugar in water becomes supersaturated, the molecules of sugar diffuse toward each other, and the sugar crystallizes out of solution. In the same manner, a solid copper-silver alloy, to be discussed in Phase Diagram, with a higher copper concentration may cause supersaturation of copper in silver, which forces the diffusion of copper atoms to increase the concentration of copper locally, causing them to precipitate.
Why are mercury and gallium of interest as components of direct restorative materials?
The diffusion of elements in most crystalline solids at room temperature is very low. Yet at temperatures that are a few hundred degrees higher, the bond energy between atoms decreases, thus allowing rapid atomic diffusion. For the same reason, the diffusion in a crystalline solid of a lower melting point, the rate of diffusion is greater. Diffusion in a noncrystalline material may occur at a more rapid rate and often may be evident at room or body temperature. The disordered structure enables the molecules or atoms to diffuse more rapidly with less activation energy. Both mercury and gallium are liquid at room temperature because of their melting points at –38.36 °C (–7.05 °F) and 29.78 °C (85.60 °F), respectively. When either liquid metal is mixed with a suitable metal alloy, atoms in the alloy dissolve and diffuse rapidly within the liquid metal at intraoral temperature. The result is a new solid metal compound. This process has been used in dentistry for making metallic direct restorative materials ( Chapter 8, Amalgam Alloy ).
Can you describe metals by their appearance and usage?
Metals
Metals, by their appearance and usage, have several characteristics. A clean, polished metallic surface exhibits a luster that cannot be duplicated by other classes of solids. Most metals emit a metallic sound (“ring”) when they are struck by another metal. They are generally excellent thermal and electrical conductors. Metals are more resistant to deformation or rupture from external forces and denser than nonmetallic structures. Most metals are also far more ductile and malleable than nonmetals, which are generally brittle. The significance of strength, ductility, malleability, and brittleness will be discussed in the chapter on mechanical properties ( Chapter 4 ). Only a few metals are resistant to tarnish and corrosion in the air at room temperature.
Pure metals, in common with other elements, can be identified by their specific melting and boiling points and by their basic physical and chemical properties. Some relevant properties for metals of dental interest are listed in Table 2-1 . Information about the atom size, crystal structure, and chemical properties of these elements is provided in Table 2-2 . Note that several important elements have more than one crystal structure (polymorphic forms), depending on the temperature, and have multiple valences.
| Element | Symbol | Atomic Weight | Melting Point (°C) | Density (g/cm 3 ) | Thermal Expansion Coefficient (10 −6 /K) |
|---|---|---|---|---|---|
| Aluminum | Al | 26.98 | 660 | 2.70 | 23.6 |
| Carbon | C | 12.01 | 630 | 2.22 | 6.0 |
| Chromium | Cr | 52.00 | 271 | 7.19 | 6.2 |
| Cobalt | Co | 58.93 | 1495 | 8.85 | 13.8 |
| Copper | Cu | 63.54 | 1083 | 8.96 | 16.5 |
| Gold | Au | 196.97 | 1063 | 19.32 | 14.2 |
| Indium | In | 114.82 | 156 | 7.31 | 33.0 |
| Iridium | Ir | 192.20 | 2454 | 22.50 | 6.8 |
| Iron | Fe | 55.85 | 1527 | 7.87 | 12.3 |
| Magnesium | Mg | 24.31 | 650 | 1.74 | 25.2 |
| Mercury | Hg | 200.59 | −39 | 13.55 | 40.0 |
| Molybdenum | Mo | 95.94 | 2619 | 10.22 | 4.9 |
| Nickel | Ni | 58.71 | 1453 | 8.90 | 13.3 |
| Palladium | Pd | 106.40 | 1552 | 12.02 | 11.8 |
| Platinum | Pt | 195.09 | 1769 | 21.45 | 8.9 |
| Rhodium | Rh | 102.91 | 1966 | 12.44 | 8.3 |
| Silicon | Si | 28.09 | 1410 | 2.33 | 7.3 |
| Silver | Ag | 107.87 | 961 | 10.49 | 19.7 |
| Tin | Sn | 118.69 | 232 | 7.30 | 23.0 |
| Titanium | Ti | 47.90 | 1668 | 4.51 | 8.5 |
| Tungsten | W | 183.85 | 3410 | 19.30 | 4.6 |
| Zinc | Zn | 65.37 | 420 | 7.133 | 0.4 |
| Element | Symbol | Atom Radius (nm) | Crystal Structure | Oxidation States | Electronegativity |
|---|---|---|---|---|---|
| Aluminum | Al | 0.118 | Face-centered cubic | +3 | 1.61 |
| Carbon | C | 0.067 | Hexagonal (graphite)Diamond cubic (diamond) | +4 | 2.55 |
| Chromium | Cr | 0.166 | Body-centered cubic | +2, +3 | 1.66 |
| Cobalt | Co | 0.152 | Hexagonal close-packed (≤450 °C)Face-centered cubic (>450 °C) | +2 | 1.88 |
| Copper | Cu | 0.145 | Face-centered cubic | +1, +2 | 1.9 |
| Gallium | Ga | 0.136 | Orthorhombic | +3 | 1.81 |
| Gold | Au | 0.174 | Face-centered cubic | +1, +3 | 2.54 |
| Iron | Fe | 0.156 | Body-centered cubic (≤911 °C)Face-centered cubic (>911 °C) | +2, +3 | 1.83 |
| Mercury | Hg | 0.171 | (Liquid above –39 °C) | +2 | 2 |
| Molybdenum | Mo | 0.190 | Body-centered cubic | +6 | 2.16 |
| Nickel | Ni | 0.149 | Face-centered cubic | +2, +3 | 1.91 |
| Platinum | Pt | 0.177 | Face-centered cubic | +2, +4 | 2.28 |
| Palladium | Pd | 0.169 | Face-centered cubic | +2, +4 | 2.2 |
| Silver | Ag | 0.165 | Face-centered cubic | +1 | 1.93 |
| Tin | Sn | 0.145 | Body-centered tetragonal | +4 | 1.96 |
| Titanium | Ti | 0.176 | Hexagonal close-packed (≤880 °C)Body-centered cubic (>880 °C) | +2, +3, +4 | 1.54 |
| Zinc | Zn | 0.142 | Hexagonal close-packed | +2 | 1.65 |
Of the 118 elements currently listed in the periodic table ( Figure 2-12 ), about 88 can be classified as metals. One common feature in the atomic structure of metal elements is that the outermost electrons (valence electrons) around the neutral metallic atoms are easily given up. This property constitutes the basis of the metallic bond, where a “cloud” of valence electrons becomes the “glue” for the positively charged ionic cores made of the nuclei and the balance of the bound electrons. The bond is nondirectional, allowing the ionic cores to break interatomic bonds with facility and establish new bonds with other ionic cores at different positions. This property enables metals to undergo distortion before fracturing (properties of ductility and malleability) under mechanical stresses. It should be noted that many metals of importance for dental alloys are transition elements that have incomplete inner electron shells.
Why are pure metals not often used for dental applications?
Pure metals have limited uses in dental and engineering applications because they are too soft to resist deformation from external forces. It has been found that mixtures of two or more metallic elements or mixtures of one or more metals and nonmetallic elements often result in improved properties beyond what the individual elements can provide. This combination of elements is called an alloy. An alloy system refers to all possible combinations (weight percentage or atomic percentage) of elements in the system. Alloys are generally prepared by fusion of the elements above their melting points. In this book, the term “metal” is used all-inclusively to include alloys and pure metals. If the phenomenon discussed does not apply to both alloys and pure metals, a distinction will be made. The mechanisms for property improvement, and the effects on physical and chemical properties of alloys, will be discussed in some detail in Chapter 9 .
To make a metal prosthesis, an exact wax or plastic pattern of the prosthesis is prepared initially. Using a dental investment, an expanded mold is prepared from the pattern, into which the molten alloy is filled under pressure ( Chapter 14, Investing Procedure ). When the alloy solidifies, it shrinks to the intended dimensions. The mechanical properties of the final cast prosthesis can often be manipulated by regulating the cooling procedure. In this chapter segment, we will first discuss how a solid emerges from the molten (liquid) state and the resultant microstructure of pure metals. We then will discuss the criteria for forming different types of alloys, the interpretation of phase diagrams, and refinement of the properties of alloys by controlled heating and cooling.
Why do dental alloys begin freezing by heterogeneous nucleation rather than by homogeneous nucleation?
Solidification and Microstructure of Metals
If a pure metal is melted and cooled to room temperature under very clean laboratory conditions, a graph of its temperature-versus-time behavior would look similar to that in Figure 2-13 . The temperature decreases steadily from point A to point B′. An increase in temperature then occurs from point B′ to point B, at which time the temperature remains constant until the time indicated at point C is reached. Subsequently, the temperature of the metal decreases steadily to room temperature (point D). The temperature, T f , indicated by the “plateau” portion of the curve at point BC, is the solidification temperature of the pure metal. This is also the melting point or fusion temperature.
During the initial cooling, the molten metal remains in the liquid state to point B′, which is below the solidification temperature; this process is termed supercooling or undercooling. During the supercooling stage, crystallization of the pure metal begins. Once the crystals begin to form, the release of latent heat from liquid changing to solid causes the temperature to rise to T f , where it remains until crystallization is completed at point C. Supercooling of pure metals occurs only in clean, inert containers under circumstances in which heterogeneous nucleation of metal crystals is not possible.
Nucleus Formation
Solidification of a metal requires nuclei of crystallization. The liquid state can be imagined as one of a multitude of random atoms surrounding numerous unstable atomic aggregates or clusters that are attempting to form crystal nuclei. These temporary nuclei are called embryos. On approaching the solidification temperature, these embryos increase in number and size, but they are still unstable and tend to dissolve into the liquid matrix. However, once the supercooled process is completed, there is a tendency for some of these embryos to survive and thus become solidification nuclei. This method of nucleus formation is called homogeneous nucleation, and the formation of nuclei in the molten metal is a random process with equal probability of occurring at any point in the melt.
Another method of forming stable embryos is for the atoms in the molten metal to contact some surface, particles in the melt, or particles in the container. Such a process is known as heterogeneous nucleation because a foreign body “seeded” the nucleus. The distribution of these foreign bodies is not necessarily random. Under dental casting conditions, heterogeneous nucleation occurs at the mold wall (investment) or in the molten alloy at impurity particles or particles of special grain-refining elements.
It should be noted that supercooling is not necessary for heterogeneous nucleation. In fact, heterogeneous nucleation should account for most of the nucleation that occurs during dental casting. The uniformity of distribution of heterogeneously formed nuclei is difficult, in principle, to control because the distribution of the different types of “seeds” is not likely to be uniform.
Why might the rates of cooling and solidification affect the grain size of a dental casting?
Liquid-to-Solid Transformation of Cast Metals
As a liquid metal cools and solidifies, crystals are formed by diffusion of atoms from the molten metal to existing nuclei of crystallization. The crystals extend along certain favored crystallographic directions and form dendrites. The heat released by the solidifying metal (in extending the arms of dendrites) reduces supercooling around the arm, which impedes solidification in affected regions and results in highly elongated extension of the arms. Subsequently, secondary extension arms of the dendrites branch off the primary arms, and tertiary arms off the secondary arms, resulting in a three-dimensional dendritic structure. Dendrites can also grow perpendicular to the walls of the mold cavity toward the center of the casting. Numerous dendritic crystals form and grow until they eventually become large enough to impinge upon each other. Solidification is completed when the liquid metal in the interdendritic spaces between the dendrite arms freezes to yield crystals. Each crystal within the solidified metal is called a grain. Thus a grain is a microscopic single crystal that has a different orientation from that of the adjacent grains. The interface between adjacent grains is called a grain boundary. The solidified metal is also called polycrystalline.
The grain boundaries in a solidified metal are revealed when examining the microstructure of a polished surface that has been etched. When a highly polished metal surface is subjected to a chemical or an electrolytic etching solution, the atoms near grain boundaries are preferentially removed, creating microscopic grooves that scatter light because these atoms have higher energy compared with atoms in the interiors of grains. An as-cast dendritic pattern may be present, like the Pd-based alloy shown in Figure 2-14 , or a dendritic microstructure may not exist, such as for the high-palladium alloy having an equiaxed grain structure shown in Figure 2-15 . The term equiaxed means that the three dimensions of each grain are similar, in contrast to the elongated shape and size of the dendrites in Figure 2-14 .
Figure 2-16 is a schematic illustration of the coalescence of individual grains to form an equiaxed grain structure. Solidification starts from isolated nuclei in the molten metal, and these crystals gradually grow by the clustering of atoms and extend toward each other. When the adjacent crystals are eventually in contact, their growth stops, as shown in Figure 2-16, E. The final polished and etched microstructure depicting grain boundaries is shown in Figure 2-16, F. The grain boundary regions are the last portions of the molten metal to solidify and can serve as sites for precipitate formation and sinks for impurity atoms.
How does the grain size affect the properties of cast dental alloys?
Grain Refinement and Grain Size
In general, the smaller the grain size of the metal, the better are its physical properties. For example, the yield strength of many metals has been found to vary inversely with the square root of the grain size. Consequently, obtaining a small grain size during casting is an advantage.
There are two variables that can reduce the grain size during solidification: number of nuclei of crystallization and rate of crystallization. A greater number of nuclei will yield a greater number of grains, and thus the size of each grain is reduced. As pointed out previously, the formation of nuclei of crystallization can be increased by the amount of supercooling and the rate of cooling. If the crystals grow more rapidly than the formation of nuclei of crystallization, the grains will be larger than if the reverse condition prevails. Conversely, if the formation of nuclei occurs faster than the growth of crystals, a small grain size will be obtained. Consequently, slow cooling results in large grains. In other words, the more rapidly the liquid state can be changed to the solid state, the smaller or finer the grains will be. Small concentrations of high-melting-point metals, such as iridium (Ir), ruthenium (Ru), or rhenium (Re), are generally added as grain-refining elements to noble metal casting alloys for dental prostheses.
