It’s a classic prank: Fill a saltshaker with sugar and wait for the meal to take an unexpected turn as your dining companions wonder why their chicken is oddly sweet. Or go the other direction and you can really disrupt someone’s morning when they take the first salty sip of coffee. These pranks work well because salt and sugar are almost indistinguishable by the naked eye: Both are crystalline solids with similar structures. Nonetheless, they have very different flavors, and they behave differently too. For example, you can pass an electrical current through salt water and light a light bulb (you might have done this experiment yourself); but you can’t do this with sugar water. Differences arise from the different properties of the two crystals, including the atoms that compose them and the actual structure of the crystal itself. In this module we will explore different types of solids and discuss how their structures relate to their behavior.
From ancient Greece until the birth of modern chemistry in the 17th century, people may have been confused about what made salt and sugar so different. Without today’s tools to identify the components in the crystals and their structures, the two would have looked as similar to them as they do to our naked eye today (see Figures 1 and 2). As scientists began identifying and characterizing elements in the 17th and 18th centuries, they would have been able to determine that salt is made of sodium and chlorine, while sugar consists of carbon, hydrogen, and oxygen, but they would probably still have wondered how such combinations of completely different elements lead to such similar-looking crystals.
It wasn’t until the early 1900s that scientists were first able to look inside crystals, when German scientist Max von Laue and English father and son scientists William Bragg and Lawrence Bragg developed a method that uses X-rays to determine the microscopic structures of crystalline solids. In fact, salt was the first solid investigated by this method, called X-ray crystallography, which revealed a regular lattice of sodium and chlorine atoms. Applying X-ray crystallography to sugar reveals a similar but not identical well-ordered crystal (Figure 3). Similarities in their crystal structure account for similarities in crystal appearance; however, the different types of atoms that make up each crystal and the different arrangements of the atoms account for the differences in behavior between the two solids. X-ray crystallography has also become a critical tool in modern biology research, helping to reveal the double helix structure of DNA in the 1950s (see our DNA II: The Structure of DNA module) and the structure of many simple and complex biological systems since that time.
Now that researchers can see this level of detail through X-ray crystallography and other methods, they can understand why some solids behave the way they do. And they can also use their understanding of the relationship between structure and behavior to design new and useful materials.
What is a solid?
You may not think of salt and sugar as solids because when you see them in the kitchen they are such small particles. But each of these particles is as much a solid as a wooden table, a glass window, or a gold piece of jewelry. A solid is a collection of atoms or molecules that are held together so that, under constant conditions, they maintain a defined shape and size. Solids, of course, are not necessarily permanent. Solid ice can melt to form liquid water at room temperature, and extremely high temperatures can be used to melt solid iron so it can be shaped into a skillet, for example. Once that skillet is formed and cools back to room temperature, though, its shape and size will not change on its own, as opposed to molten metal, which can be made to drip and change shape by gravity and molds. The same is true for ice cubes that are kept in the freezer: Once they are formed, their size and shape doesn’t change. Solids have constant shape and size because they are formed when the attractive forces between individual atoms or molecules are greater than the energy causing them to move apart. In other words, the atoms or molecules don’t have enough energy to move and are stuck together in whatever shape they were in when they lost the energy to separate. (See our States of Matter module for more about how solids differ from other states of matter.)
Salt and sugar are both crystalline solids. The other main category of solids is called amorphous. While crystalline solids are well ordered at the atomic level, with each atom or molecule inhabiting a specific point on a lattice, amorphous solids are disordered at an atomic level, with the atoms or molecules held together in a completely random formation. Consider a game of checkers. A board carefully set up with a checker in each square is analogous to a crystalline solid, while an amorphous solid could be represented as a checker pieces randomly scattered across the board.
Quartz and glass are atomic-level examples of these two categories of solids. Quartz is a crystalline solid containing a high silicate (SiO2) content. If we were to examine the structure of quartz, we could see that the silicate subunits are arranged very precisely (see Figure 4). Glass, on the other hand, is an amorphous solid. Although its typical smooth, transparent appearance may make it seem like it must have a neat, organized microscopic structure, the opposite is true: The silicate units are unevenly scattered throughout the solid in a completely disordered fashion.
Like quartz, glass has a very high silicate (SiO2) content. (See our Defining Minerals and The Silicate Minerals modules for more about silicates and quartz.) The crucial difference between crystalline and amorphous solids is not what they are made of, but how they are made, and more precisely how their structures are arranged. Quartz forms on a very slow, geological timescale so the atoms have time to achieve a highly ordered crystal structure, in which the atoms optimize the attractive forces and minimize the repulsive forces between them and which is therefore energetically favorable. Glass, on the other hand, is made by melting sand (among other methods) and letting it cool very quickly, “freezing” the atoms in place, resulting in a disordered amorphous solids. Amorphous solids are often formed when atoms and molecules are frozen in place before they have a chance to reach the crystalline arrangement, which would otherwise be the preferred structure because it is energetically favored.
One important consequence of the irregular structure of amorphous solids is that they don’t always behave consistently or uniformly. For example, they may melt over a wide range of temperatures, in contrast to a crystalline solid’s very precise melting point. Returning to the glass versus quartz example, the most prevalent type of glass, called soda lime glass, can melt anywhere between 550°C and 1450°C, while cristobalite, a quartz polymorph, melts precisely at 1713°C. In addition, amorphous solids break unpredictably and produce fragments with irregular, often curved surfaces, while crystalline solids break along specific planes and at specific angles defined by the crystal’s geometry. (See our Defining Minerals module for more about how a crystal’s external appearance reflects the regular arrangement of its atoms.)
As an amorphous solid, glass has a precise melting point.
Crystal structure determines a lot more about a solid than simply how it breaks. Structure is directly related to a number of important properties, including, for example, conductivity and density, among others. To explain these relationships, we first need to introduce the four main types of crystalline solids – molecular, network, ionic, and metallic – which are each described below.
Individual molecules are composed of atoms held together by strong covalent bonds (see our Chemical Bonding module for more about covalent bonding). To form molecular solids, these molecules are then arranged in a specific pattern and held together by relatively weak intermolecular forces. Examples include ice (H2O(s) – s here stands for "solid") and table sugar (sucrose, C12H22O11). The individual water and sugar molecules each exist as their own independent entities that interact with their neighbors in specific ways to create an ordered crystalline solid. (See Figure 5).
In network solids, on the other hand, there are no individually defined molecules. A continuous network of covalent bonds holds together all the atoms. For example, carbon can form two different network solids: diamond and graphite. These materials are made up of only carbon atoms that are arranged in two different ways. Diamond is a three-dimensional crystal that is the hardest known natural material in the world. In contrast, graphite is a two-dimensional network solid. The carbon atoms essentially form flat sheets, which are relatively slippery and can slide past each other. While these two materials are made of the same very simple component – just carbon atoms – their appearance and behavior are completely different because of the different types of bonding in the solids. (See our Defining Minerals module for more about diamond and graphite.) This ability of a single element to form multiple solids is called allotropy.
Network solids can also incorporate multiple elements. For example, consider quartz, the second most abundant material in the earth’s crust. The chemical formula for quartz is SiO2, but this formula indicates the ratio of silicon to oxygen and is not meant to imply that there are distinct SiO2 molecules present. Each silicon atom is bonded to four different oxygen atoms and each oxygen atom is bonded to two different silicon atoms, creating a large network of covalent bonds, as shown in Figure 6. (See our Defining Minerals and The Silicate Minerals modules for more about quartz.)
Ionic solids are similar to network solids in one way: There are no distinct molecules. But instead of atoms held together by covalent bonds, ionic solids are composed of positively and negatively charged ions held together by ionic bonds. (See our Chemical Bonding module for more about ionic bonding.) Table salt (sodium chloride, NaCl) is a common ionic solid, as is just about anything that is called a “salt.” Simple salts usually consist of one metal ion and one non-metal ion. In the case of sodium chloride, sodium is the metal and chloride is the non-metal. Salts can also consist of more complex ions, such as ammonium sulfate, whose components ammonium (NH4+) and sulfate (SO42-), are individually held together by covalent bonds and attracted to each other through ionic bonds.
Finally, metallic solids are a type all their own. Although we are discussing them last here, about three quarters of the known elements are metals. You can read more about these metallic elements in The Periodic Table of Elements module. Here we will focus on how these elements behave as metallic solids.
Metal atoms are held together by metallic bonds, in which the atoms pack together and the outer electrons can easily move around within the solid (Figure 7). Metallic bonds are nondirectional, meaning that metal atoms can remain bonded while they roll against each other as long as some parts of their surfaces are in contact. These unique properties of metallic bonds are largely responsible for some of the valuable behavior of metals, including their conductivity and malleability, which we discuss in the next section.
All crystalline solids are held together by covalent bonds.
Properties of solids
As described in the previous section, crystalline solids can vary in their atomic compositions, bonding, and structure. Together, these attributes determine how the different solids behave under different conditions. Solids have many different properties, including conductivity, malleability, density, hardness, and optical transmission, to name a few. We will discuss just a handful of these properties to illustrate some of the ways that atomic and molecular structure drives function.
Electrical and thermal conductivity
As you read this lesson on your computer, you’re probably not thinking about the wires your computer uses to get the electrical power it needs to run. Those wires are made of metal, probably copper, because metals generally have good electrical conductivity. Electricity is essentially a flow of electrons from one place to another, and in metallic bonds the outer electrons are relatively free to move between adjacent atoms. This electron mobility means it is easy for an electrical current to move from one end of a piece of metal to the other. When an electron is introduced at one end of a piece of wire by an electric current, this causes electrons to move from one to another metal atom continuously down the wire, allowing the current to flow. In other solids, though, the electrons are engaged in the covalent or ionic bonds and therefore are not able to conduct electricity, or do so only poorly. Materials that do not conduct electricity are called electrical insulators.
Heat, or thermal, conductivity is closely related to electrical conductivity. Just as metals are good electrical conductors, you probably know from experience that they’re good at conducting heat too. (That’s why most kitchen pots, pans, and baking sheets are metal, so they can absorb the heat from the stove or oven and pass it on to the food that’s being cooked.) To understand how this works, consider that temperature is a measurement of how much molecules are moving (see our States of Matter and Temperature modules). For a solid to conduct heat, the movement of one molecule or atom needs to be easily transferrable to its neighbor. The non-directional nature of the metallic bond makes this type of transfer relatively easy, so metals conduct heat well. In a network solid, on the other hand, where the bonds are more rigid and the angles between the atoms are strictly defined, such transfer is more difficult. Such solids would be expected to have low heat conductivity and would be called heat insulators.
Graphite is an interesting exception to this trend. Because of the specific energy and orientation of the typical bonds in graphite sheets, they are relatively good at conducting heat and electricity. You may have heard about carbon nanotubes, which are similar to graphite sheets but exist in the form of tubes (Figure 8). These tubes can conduct electricity and heat from one end to the other and are being tested for many possible applications, including in electrical circuits, solar cells, and textiles.
Metal conducts heat and electricity well because the bonds between atoms are
Malleability and ductility
Two additional properties, malleability and ductility, follow trends similar to those for electrical and thermal conductivity. Malleability describes the ability to hammer a solid into a sheet without breaking it, and ductility refers to whether a solid can be stretched to form a wire. As you may have guessed, metals tend to be both malleable and ductile, largely due to the non-directionality of metallic bonds. In contrast, covalent and ionic bonds, which are directional and require specific geometries resulting in fixed three-dimensional lattice structures, make many other types of solids brittle so they break under force.
Metallic malleability and ductility are a crucial reason that metals are so useful. Their electrical conductivity would be much less useful if it weren’t possible to stretch them into wires that could then be bent and shaped at room temperature for an incredible array of applications. They also create some drawbacks though. Metal jewelry can be crushed and deformed in the bottom of a purse, or a metal figurine can be dented if it’s dropped. Manufacturers must consider all the properties of the materials they plan to work with to find the best option for each application.
Another way to deform a solid is to melt it. A solid’s melting point depends on the strength of the interactions between its components: Stronger interactions mean a higher melting point. For molecular solids, melting means breaking the weak intermolecular forces (the forces between different molecules), not the strong covalent bonds that hold the individual molecules together, so a compound like sugar can be easily melted on your stovetop. For network solids (held together by covalent bonds), ionic solids (held together by ionic bonds), and metallic solids (held together by metallic bonds), though, the melting temperature depends on the strength of the specific bonds in each solid. Some metals have relatively low melting points, like mercury, which is actually a liquid at room temperature (its melting point is -38°C), while others, such as tungsten, melt only at extremely high temperatures (tungsten’s melting point is 3,422°C). Among network solids, a type of quartz called tridymite melts at 1,670°C while graphite melts at 4,489°C, and among ionic solids, sodium chloride melts at 801°C while lithium bromide melts at 552°C. Ionic bonds tend to be weaker than covalent and metallic bonds, which is why the melting points for these salts are somewhat lower than most of the other example melting points included here.
Melting is one way of changing a solid’s shape. Another approach is dissolving the solid into some type of liquid, in this case referred to as a solvent. The extent to which a solid dissolves in a particular solvent is called its solubility. Solids can be dissolved into a variety of types of solvents, but for now we will focus on solubility in water.
Dissolving a solid requires breaking different types of bonds for different types of solids. Dissolving a metal requires breaking metallic bonds, and dissolving a network solid requires breaking covalent bonds. Both of these types of bonds are very strong and hard to break. Therefore, metals and network solids are generally not soluble in water. (Diamond rings probably wouldn’t be as valuable if the band and the stone dissolved in the shower.) In contrast, dissolving a molecular solid requires breaking only weak intermolecular forces, not the covalent bonds that actually hold the individual molecules together. Therefore, molecular solids are relatively soluble, as you might have been able to guess given how we use sugar in so many drinks.
Finally, to dissolve ionic solids, the ionic bonds between the atoms or molecules must be broken, which water does particularly well. Each atom or molecule within an ionic solid carries a charge, and water molecules also carry a charge due to polarity (see our Water: Properties and Behavior module for more information). As a result, the negative charges within water are attracted to the positively charged ions, and the positive charges within water are attracted to the negative ions. This allows the water molecules to dissolve ionic solids by separating the parts, essentially trading the favorable ionic interaction in the solid crystal with favorable ionic interactions between the individual ions and the water molecules. Therefore, most salts are relatively water-soluble.
Both salt and sugar are quite soluble in water, but because of the differences between ionic solids (salt) and molecular solids (sugar), salt water behaves differently than sugar water (remember the light bulb experiment from the previous section). When salt dissolves in water, the positively (Na+) and negatively (Cl-) charged ions that compose the solid separate, creating a liquid solution of charged particles. These charged particles can pick up electrons and transfer them across the solution, effectively conducting electricity. When salts such as ammonium sulfate dissolve, the ionic bonds between the ions break, but the covalent bonds holding the individual complex ions together remain intact. By comparison, when sugar dissolves, each individual sucrose molecule separates from its neighbors but the sucrose molecules themselves remain intact and without charge, so they don’t conduct electricity.
Dissolving a molecular solid requires breaking
Density, defined as the amount of mass that exists in a certain volume (see our Density module for more information), is another important property that depends on the solid’s structure and composition. It’s important to note that although we described the different types of crystal solids as having certain structural characteristics, there is significant variation within each type as well. For example, metallic solids do not all share a similar arrangement of atoms. The atoms and molecules that make up crystals can pack in many different ways, which affects density (Figure 9). Imagine a jar of neatly ordered marbles, with each dimple between marbles in one row filled with a marble in the row above. This closely packed arrangement leads to a very high density. Gold takes on approximately this type of packing, resulting in its high density of 19.3 grams per cubic centimeter. Now imagine another jar where the marbles are still neatly ordered, but each marble is stacked directly on top of another instead of in the dimple. This type of packing leaves a lot more empty space in the jar because those dimples aren’t filled, so if the jar is the same size as the first jar, it can’t hold as many marbles and is less dense. Lithium, which is the least dense metal at 0.534 grams per cubic centimeter, is an example of this type of packing.
Another important variable is size. Bigger marbles can’t pack as closely as smaller marbles, even if they are in the same arrangement, so contents of the jar will be less dense. However, if you are allowed to use marbles of different sizes, you might be able to fit small marbles in the holes left between big marbles, which could lead to an even higher density than you would get from just the small marbles alone. This principle is particularly relevant for ionic solids, which are made up of two different ions that are usually different sizes. Lithium bromide, for example, is denser than potassium chloride. The size difference between lithium and bromide is greater than the size difference between potassium and chloride, so the lithium and bromide ions leave less empty space when they pack together than the potassium and chloride do, resulting in a higher density.
While the properties of solids may at first appear trivial, the unique characteristics of different solids influence almost every aspect of daily life in more ways than you may think. Fine watches and, increasingly, other electronic devices use sapphire crystals instead of glass because the strong network bonding makes sapphire incredibly hard (in fact, it is the third hardest substance known) and scratch-resistant. The peculiar molecular structure of ice results in its being less dense than liquid water, and it can be argued that without this property life on Earth would have never have come into existence. On a less existential level, it means that we can go ice skating on frozen ponds in the winter even if it hasn’t frozen all the way through. Developing new solid materials with specific properties, such as electrical semiconductors and superconductors, is an active area of research with many potential applications. But solids aren’t the only substances with useful and entertaining properties, as we will see in the next modules on liquids and gases.
Solids are formed when the forces holding atoms or molecules together are stronger than the energy moving them apart. This module shows how the structure and composition of various solids determine their properties, including conductivity, solubility, density, and melting point. The module distinguishes the two main categories of solids: crystalline and amorphous. It then describes the four types of crystalline solids: molecular, network, ionic, and metallic. A look at different solids makes clear how atomic and molecular structure drives function.
A solid is a collection of atoms or molecules that are held together so that, under constant conditions, they maintain a defined shape and size.
There are two main categories of solids: crystalline and amorphous. Crystalline solids are well ordered at the atomic level, and amorphous solids are disordered.
There are four different types of crystalline solids: molecular solids, network solids, ionic solids, and metallic solids. A solid's atomic-level structure and composition determine many of its macroscopic properties, including, for example, electrical and heat conductivity, density, and solubility.