Solids, Liquids and Gases
The three states of matter--solid, liquid, and gas--differ primarily in two respects: a) the distance between the ions or molecules, and b) the extent to which the ions or molecules move. In the solid and liquid states, the ions or molecules are very close, whereas in the gaseous state, these particles are separated by relatively large distances. In the solid state, the ions or molecules do not translate; that is, they move around within the rigid form that constrains them. These particles do, however, vibrate about their positions. In the liquid and solid states, the particles are free to translate.
At this point you are probably wondering why some substances exist in the solid state, whereas others exist in the liquid or gaseous states at room temperature. We are familiar with the three states in which water exists--as ice up to its melting point of 0 °C, as a liquid between 0 and 100 °C, and as a gas (steam) above 100 °C. We also know that any substance can be vaporized if the temperature is high enough. Hence, it appears that there are two factors--the nature of the substance and the temperature--that determine in what state it exists.
Let's deal first with the nature of the substance. There four basic types of solids: a) ionic, b) molecular, c) covalent network, and d) metallic. As we know from our previous discussion of bonding, ionic compounds consist of ions. These ions are packed very efficiently to make the best use of the available space and to maximize the number of ions of opposite charge that surround a particular ion. Figure 43 shows a photo of the sodium chloride lattice (a lattice is just a regular, systematic arrangement of particles). Part (b) of this figure focuses in on a smaller part of the lattice so that we can see a sodium ion (the silver sphere in the center) surrounded by chloride ions.
Figure 43. The NaCl Structure.
[When you examine models of this type, ignore the metal rods that link the spheres; they exist simply to support the entire structure.] If you examine this arrangement carefully, you will find that the sodium ion is surrounded by six nearest neighboring chloride ions. This is called the coordination number. The coordination number of each chloride ion is also six. Figure 44 shows a model of zinc sulfide. A careful examination of a zinc ion (silver) shows that for this compound the coordination number is four.
Figure 44. The ZnS Structure.
Thus, some aspect of zinc sulfide makes it more favorable for the zinc to be surrounded by only four anions (it is believed that the relative size of the cation and anion determine the coordination number). The ions in all ionic compounds, regardless of structure and coordination number, are held together by electrostatic forces. These are the very same forces that produce the static electricity that you experience after shuffling along a carpet on a dry day, and they are the very same forces that we have discussed previously. Because these forces are very strong, it takes a lot of energy to melt ionic compounds. That is, ionic compounds will only convert from the solid state to the liquid state at high temperatures. The melting point of NaCl is 804 °C; the melting point of calcium carbonate (the main ingredient in limestone) is 1339 °C.
Molecular solids contain molecules. These molecules are also efficiently packed as shown by the model of solid carbon tetrachloride in Figure 45. The forces between molecules, called intermolecular forces, are much weaker than the electrostatic ionic forces. Consequently, the melting points of molecular solids are much lower than that of ionic solids.
Figure 45. The CCl4 Structure.
Carbon tetrachloride, for example, melts at -23 °C. Water, crucial to life but frequently taken for granted, is a very unusual molecular substance. Compounds similar to H2O, such as H2S, NH3, CH4, and HF (you might want to look at the periodic table to find out why we consider these to be similar to H2O), are gases at room temperature. Water, on the other hand, is a liquid and boils at 100 °C. It also has greater viscosity, higher density, and greater surface tension than similar compounds. All of these characteristics are a result of unusually strong intermolecular forces between water molecules. These forces are called hydrogen bonding, and they occur only when a molecule contains a hydrogen connected to a very electronegative atom (fluorine is the most electronegative atom, oxygen is the second most electronegative atom). These special forces are not as strong as the electrostatic ionic forces, but they have a characteristic that makes them unique among the repertoire of forces at the disposal of molecules--direction. The hydrogen bonding in water is shown in Figure 46 by dotted lines. These dotted lines always go from a hydrogen of one molecule to a lone pair of electrons on an atom of a neighboring molecule.
Figure 46. Hydrogen-bonding in water. Notice that hydrogens are directed toward lone pairs.
Incidentally, it is this hydrogen bonding that is responsible for the fact that ice has a lower density than water (imagine the consequences for aquatic life if ice sank to the bottom of lakes). Most solids are more dense than their liquid states. Hydrogen bonding is also believed to be responsible for the unique structures of proteins. For example, the double-helix formed by many types of proteins (Figure 47) is due to hydrogen bonding between the strands.
Figure 47. Double helical structure of proteins results from hydrogen bonding.
The third type of solid is the covalent network, probably best exemplified by the diamond structure shown in Figure 48. In diamond, which consists solely of carbon atoms, each atom is covalently bonded to four other atoms. Each of these four atoms is covalently bonded to four others, and on and on. In a perfect diamond crystal, this gigantic network is one molecule. Another elemental form of carbon (carbon is a good example of allotropism, the existence of more than one form or structure for the element), graphite, contains a different type of covalent network. In graphite (Figure 49) there are layers of carbon atoms attached to one another within a plane. A close look at Figure 49 will show that each plane is essentially composed of attached benzene rings. Because covalent bonds are strong, these covalent network materials have very high melting points; graphite withstands heat up to 2800 °C. Incidentally, it is the layer structure of graphite and the consequent movement of layers past one another that give graphite its lubricating characteristics. Before leaving the elemental forms of carbon, we should mention the exciting discovery within the last decade of a third elemental form of carbon. Fullerene, named for Buckminister Fuller, an architect known for interesting geodesic forms, was found in soot. It consists of molecules that contain 60 atoms of carbon linked in the form shown in Figure 50. Notice that this structure is identical to the seams of a soccer ball, and represents an economical stable structure that arranges 60 sites within a sphere.
Figure 48. The covalent network structure of diamond.
Figure 49. The covalent network layer structure of graphite.
Figure 50. Fullerene, another allotrope of carbon.
The fourth type of solid, the metallic solid, is the elemental form for most of the elements. Indeed, almost all of the s-fillers, d-fillers, and f-fillers have metallic elemental forms. Metals are characterized by their shiny luster, malleability and ductility, ability to conduct electricity, and tendency to form cations. The atoms within a metal are usually packed in the most economical fashion with one layer that looks like that shown in Figure 51, followed by identical "close-packed" layers positioned on top of one another. The bonding in metals is not well-understood, and the forces ("metallic forces") range from weak to very strong. Hence, the melting points of solids range from -39 °C for mercury, a liquid at room temperature, to 1535 °C for iron.
Figure 51. Close-packing of atoms in a metal.
Make a drawing of the seven spheres close packed in three rows. Then place three spheres in a close packed arrangement around the middle sphere.
Now it is time to inquire about the effect of temperature. How does temperature affect the state of matter? Presumably, in order to cause a solid to convert to a liquid, or a liquid to a gas, the ions, atoms, or molecules that occupy the lattice positions of the solid must be given sufficient energy to overcome the forces that hold them in their lattice positions. The stronger the forces holding the particles in the lattice, the greater the heat that must be supplied in order to disrupt these forces. In the case of a liquid, the effect of temperature is somewhat more easily understood. Suppose we have a molecular substance in the liquid state. Figure 52 shows these molecules, represented as ellipses moving about, or translating, throughout the liquid.
Figure 52. Molecules translating and rotating in liquid state.
The molecules also rotate as they move, and some move fasster than others. The velocity with which they move determines their kinetic energy through the equation
KE = 1/2 mv2
The distribution of kinetic energy of the molecules (this is called a Boltzmann distribution) is shown in Figure 53. [This is an important and interesting graph. Notice that a) there are no molecules that have zero kinetic energy, and b) the curve does not ever return to zero again, meaning that a few molecules have exceedingly large kinetic energies. This is called approaching zero asymptotically.
Figure 53. The Boltzmann distribution of molecular velocities.
Make a drawing of this curve using a bar graph. Is a bar graph as useful or as accurate as the xy plot in Figure 53?.
Let us assume that the energy required to wrench a molecule from the "grasp" of the other molecules on the surface of liquid is Ew. Now let us redraw Figure 53 so that it will contain Ew. This curve (Figure 54) shows that all of the molecules with an energy greater than Ew will escape from the surface and enter the gaseous state. If we now increase the temperature of the liquid, the distribution of kinetic energy among the molecules changes. The distributions at two temperatures, T1 and T2, are shown in Figure 55 along with Ew. Notice that at the higher temperature, T2, a greater percentage of molecules have sufficient kinetic energy to enter the gaseous state.
Figure 54. Fraction of molecules with energy Ew necessary to escape the liquid.
Figure 55. Distribution of molecular speeds at two temperatures.