For more than 200 years, chemists have divided compounds into two categories. Those that were isolated from plants or animals were called organic, while those extracted from ores and minerals were inorganic. Organic chemistry is often defined as the chemistry of car-bon. But that definition would include calcium carbonate (CaCO3) and graphite, which more closely resemble inorganic compounds. We will therefore define organic chemistry as the study of compounds such as formic acid (HCO2H), methane (CH4), and vitamin C (C6H8O6) that contain both carbon and hydrogen. This section focuses on inorganic car-bon compounds.
Elemental Forms of Carbon: Graphite, Diamond, Coke, and Carbon Black Carbon occurs as a variety of allotropes. There are two crystalline forms—diamond and graphite—and a number of amorphous (noncrystalline) forms, such as charcoal, coke, and carbon black.
References to the characteristic hardness of diamond (from the Greek adamas, “in-vincible”) date back at least 2600 years. It was not until 1797, however, that Smithson Tennant was able to show that diamonds consist solely of carbon. The properties of diamond are remarkable. It is among the least volatile substances known (MP 3550°C, BP 4827°C), it is the hardest substance known, and it expands less on heating than any other material.
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The properties of diamond are a logical consequence of its structure. Carbon, with four valence electrons, forms covalent bonds to four neighboring carbon atoms arranged toward the corners of a tetrahedron, as shown in Figure N.24. Each of the sp3-hybridized atoms is then bound to four other carbon atoms, which form bonds to four other carbon atoms, and so on. As a result, a perfect diamond can be thought of as a single giant molecule. The strength of the individual COC bonds and their arrangement in space give rise to the un-usual properties of diamond.
FIGURE N.24 The simplest repeating unit in diamond.
In some ways, the properties of graphite are like those of diamond. Both compounds boil at 4827°C, for example. But graphite is also very different from diamond. Diamond (3.514 g/cm3) is significantly more dense than graphite (2.26 g/cm3). Whereas diamond is the hardest substance known, graphite is one of the softest. Diamond is an excellent insu-lator, with little or no tendency to carry an electric current. Graphite is such a good con-ductor of electricity that graphite electrodes are used in electrical cells.
The physical properties of graphite can be understood from the structure of the solid shown in Figure N.25. Graphite consists of extended planes of sp2-hybridized carbon atoms in which each carbon is tightly bound to three other carbon atoms. (It takes 477 kJ to break a mole of the bonds within the planes.) The strong bonds between carbon atoms within each plane explain the exceptionally high melting point and boiling point of graphite. The bonds between planes of carbon atoms, however, are relatively weak. (The force of attrac-tion between planes is only 17 kJ/mol.) Because the bonds between planes are weak, it is easy to deform the solid by allowing one plane of atoms to move relative to another. As a result, graphite is soft enough to be used in pencils and as a lubricant in motor oil.
The characteristic properties of graphite and diamond might lead you to expect that diamond would be more stable than graphite. This isn’t what is observed experimentally.
Graphite at 25°C and 1 atm pressure is slightly more stable than diamond. At very high
FIGURE N.25 Portion of the structure of extended planes of carbon atoms found in graphite.
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temperatures and pressures, however, diamond becomes more stable than graphite. In 1955 General Electric developed a process to make industrial-grade diamonds by treating graphite with a metal catalyst at temperatures of 2000 to 3000 K and pressures above 125,000 atm. Roughly 40% of industrial-quality diamonds are now synthetic. Although gem-quality diamonds can be synthesized, the costs involved are prohibitive.
Both diamond and graphite occur as regularly packed crystals. Other forms of carbon are amorphous—they lack a regular structure. Charcoal, carbon black, and coke are all amorphous forms of carbon. Charcoal results from heating wood in the absence of oxygen.
To make carbon black, natural gas or other carbon compounds are burned in a limited amount of air to give a thick, black smoke that contains extremely small particles of car-bon, which can be collected when the gas is cooled and passed through an electrostatic pre-cipitator. Coke is a more regularly structured material, closer in structure to graphite than either charcoal or carbon black, which is made from coal.
Carbides: Covalent, Ionic, and Interstitial
Carbon reacts with less electronegative elements at high temperatures to form compounds known as carbides. When carbon reacts with an element of similar size and electronega-tivity, a covalent carbide is produced. Silicon carbide, for example, is made by treating sil-icon dioxide from quartz with an excess of carbon in an electric furnace at 2300 K.
SiO2(s) 3 C(s) 88n SiC(s) 2 CO(g)
Covalent carbides have properties similar to those of diamond. Both SiC and diamond are inert to chemical reactions, except at very high temperatures; both have very high melt-ing points; and both are among the hardest substances known. SiC was first synthesized by Edward Acheson in 1891. Shortly thereafter, Acheson founded the Carborundum Com-pany to market the material. Then, as now, materials in this class are most commonly used as abrasives.
Compounds that contain carbon and one of the more active metals are called ionic carbides.
CaO(s) 3 C(s) 88n CaC2(s) CO(g)
It is useful to think about ionic carbides as if they contained negatively charged carbon ions: [Ca2][C22] or [Al3]4[C4]3. This model is useful because it explains why these car-bides burst into flame when added to water. Ionic carcar-bides such as Al4C3that formally con-tain the C4 ion react with water to form methane, which is ignited by the heat given off in the reaction.
C4 4 H2O 88n CH4 4 OH
The ionic carbides such as CaC2 that formally contain the C22 ion react with water to form acetylene, which is ignited by the heat of reaction.
C22 2 H2O 88n C2H2 2 OH
At one time, miners’ lamps were fueled by the combustion of acetylene prepared from the reaction of calcium carbide with water.
The difference between covalent carbides and ionic carbides can be understood by adding compounds such as SiC, Al4C3, and CaC2 to the bond type triangle introduced in
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Chapter 5. When those compounds are added to Figure 5.10, for example, we find that SiC falls well into the region expected for covalent compounds. CaC2, on the other hand, is clearly an ionic compound. Al4C3falls on the borderline between ionic and covalent, which is consistent with the fact that the compound is hard—as one would expect for a covalent carbide—and yet reacts with water to form methane—as might be expected for an ionic carbide.
Interstitial carbides, such as tungsten carbide (WC), form when carbon combines with a metal that has an intermediate electronegativity and a relatively large atomic ra-dius. In these compounds, the carbon atoms pack in the holes (interstices) between planes of metal atoms. The interstitial carbides, which include TiC, ZrC, and MoC, re-tain the properties of metals. They act as alloys, rather than as either salts or covalent compounds.
The Oxides of Carbon
Although the different forms of carbon are essentially inert at room temperature, they combine with oxygen at high temperatures to produce a mixture of carbon monoxide and carbon dioxide.
2 C(s) O2(g) 88n 2 CO(g) H° 221.05 kJ/molrxn
C(s) O2(g) 88n CO2(g) H° 393.51 kJ/molrxn
CO can also be obtained by reacting red-hot carbon with steam.
C(s) H2O(g) 88n CO(g) H2(g)
Because the mixture of gases is formed by the reaction of charcoal or coke with water it is often referred to as water gas. It is also known as town gas because it was once made by towns and cities for use as a fuel. Water gas, or town gas, was a common fuel for both home and industrial use before natural gas became readily available. The H2burns to form wa-ter, and the CO is oxidized to CO2. Eventually, as our supply of natural gas is depleted, it will become economical to replace natural gas with other fuels, such as water gas, that can be produced from our abundant supply of coal.
Both CO and CO2are colorless gases. CO boils at 191.5°C, and CO2sublimes (passes directly from the solid to the gaseous state) at 78.5°C. Although CO has no odor or taste, CO2has a faint, pungent odor and a distinctly acidic taste. Both are dangerous substances but at very different levels of exposure. Air contaminated with as little as 0.002 gram of CO per liter can be fatal because CO binds tightly to the hemoglobin that carries oxygen through the blood. CO2is not lethal until the concentration in the air approaches 15%. At that point, it has replaced so much oxygen that a person who attempts to breathe the at-mosphere suffocates. The danger of CO2 poisoning is magnified by the fact that CO2 is roughly 1.5 times more dense than the air in our atmosphere. Thus, CO2can accumulate at the bottom of tanks or wells.
CO2 in the Atmosphere
Carbon dioxide influences the temperature of the atmosphere by a phenomenon known as the greenhouse effect. The glass walls and ceilings of a greenhouse absorb some of the lower energy, longer wavelength radiation from sunlight thereby inevitably raising the tem-perature inside the building. CO2in the atmosphere does exactly the same thing, it absorbs
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low energy, long wavelength radiation from the Sun that would otherwise be reflected back from the surface of the planet. Thus, CO2 in the atmosphere traps heat. Although there are other factors at work, it is worth noting that Venus, whose atmosphere contains a great deal of CO2, has a surface temperature of roughly 400°C, whereas Mars, with little or no atmosphere, has a surface temperature of 50°C.
There are many sources of CO2in the atmosphere. Over geologic time scales, the largest source has been volcanos. Within the twentieth century, the combustion of petroleum, coal, and natural gas has made a significant contribution to atmospheric levels of CO2(see Fig-ure N.8). Between 1958 and 1978, the average level of CO2 in the atmosphere increased by 6%, from 315.8 to 334.6 ppm.
At one time, the amount of CO2 released to the atmosphere wasn’t a matter for con-cern because natural processes that removed CO2from the atmosphere could compensate for the CO2that entered the atmosphere. The vast majority of the CO2 liberated by vol-canic action, for example, was captured by calcium oxide or magnesium oxide to form cal-cium carbonate or magnesium carbonate.
CaO(s) CO2(g) 88n CaCO3(s) MgO(s) CO2(g) 88n MgCO3(s)
CaCO3is found as limestone or marble, or mixed with MgCO3as dolomite. The amount of CO2in deposits of carbonate minerals is at least several thousand times larger than the amount in the atmosphere.
CO2also dissolves, to some extent, in water.
H2O
CO2(g) 88n CO2(aq) It then reacts with water to form carbonic acid, H2CO3.
CO2(aq) H2O(l) 88n H2CO3(aq)
As a result of the reactions, the sea contains about 60 times more CO2 than the atmo-sphere.
Can the sea absorb more CO2from the atmosphere, or is it near its level of saturation?
Is the rate at which the sea absorbs CO2greater than the rate at which we are adding it to the atmosphere? The observed increase in the concentration of CO2 in recent years sug-gests pessimistic answers to those two questions. A gradual warming of the earth’s atmo-sphere could result from continued increases in CO2levels, with adverse effects on the cli-mate and therefore the agriculture of at least the northern hemisphere.
The Chemistry of Carbonates: CO32 and HCO3
Eggshells are almost pure calcium carbonate. CaCO3 can also be found in the shells of many marine organisms and in both limestone and marble. The fact that none of those substances dissolves in water suggests that CaCO3is normally insoluble in water. Calcium carbonate will dissolve in water saturated with CO2, however, because carbonated water (or carbonic acid) reacts with calcium carbonate to form calcium bicarbonate, which is sol-uble in water.
CaCO3(s) H2CO3(aq) 88n Ca2(aq) 2 HCO3(aq)
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When water rich in carbon dioxide flows through limestone formations, part of the limestone dissolves. If the CO2 escapes from the water, or if some of the water evapo-rates, solid CaCO3 is redeposited. When this happens as water runs across the roof of a cavern, stalactites, which hang from the roof of the cave, are formed. If the water drops before the carbonate reprecipitates, stalagmites, which grow from the floor of the cave, are formed.
The chemistry of carbon dioxide dissolved in water is the basis of the soft drink in-dustry. The first artificially carbonated beverages were introduced in Europe at the end of the nineteenth century. Carbonated soft drinks today consist of carbonated water, a sweet-ening agent (sugar, saccharin, or aspartame), an acid to impart a sour or tart taste, flavor-ing agents, colorflavor-ing agents, and preservatives. As much as 3.5 liters of gaseous CO2 dis-solve in a liter of soft drink to provide the characteristic “bite” associated with carbonated beverages.
Carbonate chemistry plays an important role in other parts of the food industry as well.
Baking soda, or bicarbonate of soda, is sodium bicarbonate, NaHCO3, a weak base, which is used to neutralize the acidity of other ingredients in a recipe. Baking powder is a mix-ture of baking soda and a weak acid, such as tartaric acid or calcium hydrogen phosphate (CaHPO4). When mixed with water, the acid reacts with the HCO3ion to form CO2gas, which causes the dough or batter to rise.
HCO3(aq) H(aq) 88n H2CO3(aq) 88n H2O(l) CO2(g)
Before commercial baking powders were available, cooks obtained the same effect by mix-ing roughly a teaspoon of bakmix-ing soda with a cup of sour milk or buttermilk. The acids that give sour milk and buttermilk their characteristic taste also react with the bicarbon-ate ion to give CO2.
Allotrope Anhydrous Carbide Diamagnetic Dimer
Disproportionation reaction
Haber process Halide
Halogen Hydride
Ostwald process Oxidizing agent
Oxyacid Oxyanion Paramagnetic Peroxide Rare gases Reducing agent
PROBLEMS
Metals, Nonmetals, and Semimetals
1. List the elements that are nonmetals. Describe where these elements are found in the periodic table.
2. Explain why semimetals (such as B, Si, Ge, As, Sb, Te, Po, and At) exist and describe some of their physical properties.
3. Which member of each of the following pairs of elements is more nonmetallic?
(a) As or Bi (b) As or Se (c) As or S (d) As or Ge (e) As or P