This results from the structural nonequivalence of the molecules in the amorphous solid. Some forces are weaker than others, and when an amorphous material is heated, the weakest intermolecular attractions break first.
As the temperature is increased further, the stronger attractions are broken. Thus amorphous materials soften over a range of temperatures. Carbon is an essential element in our world. The unique properties of carbon atoms allow the existence of carbon-based life forms such as ourselves. Carbon forms a huge variety of substances that we use on a daily basis, including those shown in Figure 7.
You may be familiar with diamond and graphite, the two most common allotropes of carbon. Allotropes are different structural forms of the same element. Diamond is one of the hardest-known substances, whereas graphite is soft enough to be used as pencil lead. These very different properties stem from the different arrangements of the carbon atoms in the different allotropes.
You may be less familiar with a recently discovered form of carbon: graphene. Graphene was first isolated in by using tape to peel off thinner and thinner layers from graphite. It is essentially a single sheet one atom thick of graphite. Graphene, illustrated in Figure 8 , is not only strong and lightweight, but it is also an excellent conductor of electricity and heat. These properties may prove very useful in a wide range of applications, such as vastly improved computer chips and circuits, better batteries and solar cells, and stronger and lighter structural materials.
In a crystalline solid, the atoms, ions, or molecules are arranged in a definite repeating pattern, but occasional defects may occur in the pattern. Several types of defects are known, as illustrated in Figure 9. Vacancies are defects that occur when positions that should contain atoms or ions are vacant.
Less commonly, some atoms or ions in a crystal may occupy positions, called interstitial sites , located between the regular positions for atoms. Other distortions are found in impure crystals, as, for example, when the cations, anions, or molecules of the impurity are too large to fit into the regular positions without distorting the structure. Trace amounts of impurities are sometimes added to a crystal a process known as doping in order to create defects in the structure that yield desirable changes in its properties.
For example, silicon crystals are doped with varying amounts of different elements to yield suitable electrical properties for their use in the manufacture of semiconductors and computer chips. Some substances form crystalline solids consisting of particles in a very organized structure; others form amorphous noncrystalline solids with an internal structure that is not ordered.
The main types of crystalline solids are ionic solids, metallic solids, covalent network solids, and molecular solids. The properties of the different kinds of crystalline solids are due to the types of particles of which they consist, the arrangements of the particles, and the strengths of the attractions between them.
Because their particles experience identical attractions, crystalline solids have distinct melting temperatures; the particles in amorphous solids experience a range of interactions, so they soften gradually and melt over a range of temperatures. Some crystalline solids have defects in the definite repeating pattern of their particles. These defects which include vacancies, atoms or ions not in the regular positions, and impurities change physical properties such as electrical conductivity, which is exploited in the silicon crystals used to manufacture computer chips.
Ice has a crystalline structure stabilized by hydrogen bonding. That is a different process. This endothermic reaction gives rise to the other definition of lattice energy: the energy that must be expended to break up an ionic solid into gaseous ions. Lattice energy, while due mainly to coulombic attraction between each ion and its nearest neighbors six in the case of NaCl is really the sum of all the interactions within the crystal.
Lattice energies cannot be measured directly, but they can be estimated from the energies of other processes.
The most energetically stable arrangement of solids made up of identical molecular units are generally those in which there is a minimum of empty space. These are known as close-packed structures, and there are several kinds of them. In ionic solids of even the simplest stoichiometry, the positive and negative ions usually differ so much in size that packing is often much less efficient.
This may cause the solid to assume lattice geometries that differ from the one illustrated above for NaCl. Consider the structure of cesium chloride, CsCl. The CsCl lattice therefore assumes a different arrangement. CsCl structure : In CsCl, metal ions are shifted into the center of each cubic element of the Cl—-ion lattice. Each cesium ion has eight nearest-neighbor chloride ions, while each chloride ion is also surrounded by eight cesium ions in 8,8 coordination. The two kinds of lattice arrangements exemplified by NaCl and CsCl are found in a large number of other ionic solids, and these names are used generically to describe the structures of these other compounds.
There are many other fundamental lattice arrangements not all cubic , but the two described here are sufficient to illustrate the point that the radius ratio the ratio of the radii of the positive to the negative ion plays an important role in the structures of simple ionic solids. Atoms in covalent solids are covalently bonded with their neighbors, creating, in effect, one giant molecule.
A covalent bond is a chemical bond that involves the sharing of pairs of electrons between atoms. This sharing results in a stable balance of attractive and repulsive forces between those atoms. Covalent solids are a class of extended-lattice compounds in which each atom is covalently bonded to its nearest neighbors. This means that the entire crystal is, in effect, one giant molecule. The extraordinarily strong binding forces that join all adjacent atoms account for the extreme hardness of these solids.
They cannot be broken or abraded without breaking a large number of covalent chemical bonds. When heated to very high temperatures, these solids usually decompose into their elements. Another property of covalent network solids is poor electrical conductivity, since there are no delocalized electrons.
When molten, unlike ionic compounds, the substance is still unable to conduct electricity, since the macromolecule consists of uncharged atoms rather than ions. This is also contrary to most forms of metallic bonds. Graphite is an allotrope of carbon. This grants graphite electrical conductivity. Its melting point is high, due to the large amount of energy required to rearrange the covalent bonds. It is also quite hard because of the strong covalent bonding throughout the lattice.
However, because of the planar bonding arrangements between the carbon atoms, the layers in graphite can be easily displaced, so the substance is malleable. Graphite is generally insoluble in any solvent due to the difficulty of solvating a very large molecule. Diamond and Graphite: Two Allotropes of Carbon : These two allotropes of carbon are covalent network solids which differ in the bonding geometry of the carbon atoms. In diamond, the bonding occurs in the tetrahedral geometry, while in graphite the carbons bond with each other in the trigonal planar arrangement.
This difference accounts for the drastically different appearance and properties of these two forms of carbon. Diamond is also an allotrope of carbon. The diamond unit cell is face-centered cubic and contains eight carbon atoms. Boron nitride BN is similar to carbon because it exists as a diamond-like cubic polymorph as well as in a hexagonal form similar to graphite. Hexagonal boron nitride : Hexagonal boron nitride, a two-dimensional material, is similar in structure to graphite.
Cubic boron nitride is the second-hardest material after diamond, and it is used in industrial abrasives and cutting tools. Cubic boron nitride : Cubic boron nitride adopts a crystal structure, which can be constructed by replacing every two carbon atoms in diamond with one boron atom and one nitrogen atom. Cubic boron nitride is the second-hardest material, after diamond.
Recent interest in boron nitride has centered on its carbon-like ability to form nanotubes and related nanostructures. Silicon carbide SiC is also known as carborundum. Its structure is very much like that of diamond, with every other carbon replaced by silicon.
Silicon carbide exists in about crystalline forms. It is used mostly in its synthetic form because it is extremely rare in nature. It is found in a certain type of meteorite that is thought to originate outside of our solar system. Structurally, silicon carbide is very complex; at least 70 crystalline forms have been identified. Its extreme hardness and ease of synthesis have led to a diversity of applications — in cutting tools and abrasives, high-temperature semiconductors and other high-temperature applications, the manufacturing of specialty steels and jewelry, and many more.
Tungsten carbide WC is probably the most widely encountered covalent solid, owing to its use in carbide cutting tools and as the material used to make the rotating balls in ball-point pens. In many of its applications, it is embedded in a softer matrix of cobalt or coated with titanium compounds. Silicon Carbide : Silicon carbide is an extremely rare mineral, and in nature is is mostly found in a certain type of meteorite.
Recall that a molecule is defined as a discrete aggregate of atoms bound together sufficiently tightly by directed covalent forces to allow it to retain its individuality when the substance is dissolved, melted, or vaporized. The two words italicized in the preceding sentence are important. Covalent bonding implies that the forces acting between atoms within the molecule intra molecular are much stronger than those acting between molecules inter molecular , The directional property of covalent bonding gives each molecule a distinctive shape which affects a number of its properties.
Liquids and solids composed of molecules are held together by van der Waals or intermolecular forces, and many of their properties reflect this weak binding. Molecular solids tend to be soft or deformable, have low melting points, and are often sufficiently volatile to evaporate directly into the gas phase. This latter property often gives such solids a distinctive odor. Thus, many corresponding substances are either liquid water or gaseous oxygen at room temperature.
Molecular solids also have relatively low density and hardness. The elements involved are light, and the intermolecular bonds are relatively long and are therefore weak. Because of the charge neutrality of the constituent molecules, and because of the long distance between them, molecular solids are electrical insulators. Because dispersion forces and the other van der Waals forces increase with the number of atoms, large molecules are generally less volatile, and have higher melting points than smaller ones.
Also, as one moves down a column in the periodic table, the outer electrons are more loosely bound to the nucleus, increasing the polarisability of the atom, and thus its propensity to van der Waals-type interactions. However, among the solids we eat, three in particular are, or are produced from, rocks. Yes, rocks! The first one is NaCl, or common salt. Salt is the only solid that we ingest that is actually mined as a rock hence the term rock salt ; it really is a rock.
Salt preserves food, a function that was much more important before the days of modern food preparation and storage. The fact that saltiness is one of the major tastes the tongue can detect suggests a strong evolutionary link between ingesting salt and survival. There is some concern today that there is too much salt in the diet; it is estimated that the average person consumes at least three times as much salt daily than is necessary for proper bodily function.
However, we do not mine these substances directly from the ground; we mine trona, whose chemical formula is Na 3 H CO 3 2. Another process, called the Solvay process, is also used to make Na 2 CO 3.
Either way, we get these two products from the ground i. NaHCO 3 is also known as baking soda, which is used in many baked goods. Na 2 CO 3 is used in foods to regulate the acid balance.
It is also used in laundry where it is called washing soda to interact with other ions in water that tend to reduce detergent efficiency. What is the difference between a crystalline solid and an amorphous solid? What two properties do solids have in common? What two properties of solids can vary? Explain how the bonding in an ionic solid explains some of the properties of these solids.
Explain how the bonding in a molecular solid explains some of the properties of these solids. Explain how the bonding in a covalent network solid explains some of the properties of these solids. Explain how the bonding in a metallic solid explains some of the properties of these solids. Which type s of solid conduct s electricity in their solid state? In their liquid state? At the atomic level, a crystalline solid has a regular arrangement of atoms, whereas an amorphous solid has a random arrangement of atoms.
The oppositely charged ions are very strongly held together, so ionic crystals have high melting points. Ionic crystals are also brittle because any distortion of the crystal moves same-charged ions closer to each other, so they repel.
The covalent network solid is essentially one molecule, making it very hard and giving it a very high melting point. Previous Section. Table of Contents. Next Section. Describe the six different types of solids. Example 5 Predict the type of crystal exhibited by each solid. Silver is a metal, so it would exist as a metallic solid in the solid state.
CO 2 is a covalently bonded molecular compound.
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