CHEMISTRY THE CENTRAL SCIENCE
22 CHEMISTRY OF THE NONMETALS
22.8 THE OTHER GROUP 5A ELEMENTS: P, As, Sb, AN D Bi
Of the other group 5A elements—phosphorus, arsenic, antimony, and bismuth— phosphorus has a central role in several aspects of biochemistry and environmental chemistry.
General Characteristics of the Group 5A Elements
The group 5A elements have the outer-shell electron configuration ns2np3, with n ranging from 2 to 6. A noble-gas configuration is achieved by adding three electrons to form the –3 oxidation state. Ionic compounds containing X 3– ions are not common, however. More commonly, the group 5A element acquires an octet of electrons via covalent bonding and oxidation numbers ranging from –3 to +5.
Because of its lower electronegativity, phosphorus is found more frequently in positive oxidation states than is nitrogen. Furthermore, compounds in which phosphorus has the +5 oxidation state are not as strongly oxidizing as the corresponding compounds of nitrogen. Compounds in which phosphorus has a –3 oxidation state are much stronger reducing agents than are the corresponding nitrogen compounds.
Some properties of the group 5A elements are listed in TABLE 22.7. The general pattern is similar to what we saw with other groups: Size and metallic character increase as atomic number increases in the group.
TABLE 22.7 • Properties of the Group 5A Elements
The variation in properties among group 5A elements is more striking than that seen in groups 6A and 7A. Nitrogen at the one extreme exists as a gaseous diatomic molecule, clearly nonmetallic. At the other extreme, bismuth is a reddish white, metallic-looking substance that has most of the characteristics of a metal.
The values listed for X—X single-bond enthalpies are not reliable because it is difficult to obtain such data from thermochemical experiments. However, there is no doubt about the general trend: a low value for the N— N single bond, an increase at phosphorus, and then a gradual decline to arsenic and antimony. From observations of the elements in the gas phase, it is possible to estimate the X ≡ X triple-bond enthalpies. Here we see a trend that is different from that for the X — X single bond. Nitrogen forms a much stronger triple bond than the other elements, and there is a steady decline in the triple-bond enthalpy down through the group. These data help us to appreciate why nitrogen alone of the group 5A elements exists as a diatomic molecule in its stable state at 25 °C. All the other elements exist in structural forms with single bonds between the atoms.
Occurrence, Isolation, and Properties of Phosphorus
Phosphorus occurs mainly in the form of phosphate minerals. The principal source of phosphorus is phosphate rock, which contains phosphate principally as Ca3(PO4)2. The element is produced commercially by reduction of calcium phosphate with carbon in the presence of SiO2:
The phosphorus produced in this fashion is the allotrope known as white phosphorus. This form distills from the reaction mixture as the reaction proceeds.
Phosphorus exists in several allotropic forms. White phosphorus consists of P4 tetrahedra (FIGURE 22.27). The bond angles in this molecule, 60°, are unusually small, so there is much strain in the bonding, which is consistent with the high reactivity of white phosphorus. This allotrope bursts spontaneously into flames if exposed to air. When heated in the absence of air to about 400 °C, white phosphorus is converted to a more stable allotrope known as red phosphorus, which does not ignite on contact with air. Red phosphorus is also considerably less poisonous than the white form. We will denote elemental phosphorus as simply P(s).
FIGURE 22.27 White and red phosphorus. Despite the fact that both contain nothing but phosphorus atoms, these two forms of phosphorus differ greatly in reactivity. The white allotrope, which reacts violently with oxygen, must be stored under water so that it is not exposed to air. The much less reactive red form does not need to be stored this way.
Phosphorus forms a wide range of compounds with the halogens, the most important of which are the trihalides and pentahalides. Phosphorus trichloride (PCl3) is commercially the most significant of these compounds and is used to prepare a wide variety of products, including soaps, detergents, plastics, and insecticides.
Phosphorus chlorides, bromides, and iodides can be made by direct oxidation of elemental phosphorus with the elemental halogen. PCl3, for example, which is a liquid at room temperature, is made by passing a stream of dry chlorine gas over white or red phosphorus:
If excess chlorine gas is present, an equilibrium is established between PCl3 and PCl5.
The phosphorus halides hydrolyze on contact with water. The reactions occur readily, and most of the phosphorus halides fume in air because of reaction with water vapor. In the presence of excess water the products are the corresponding phosphorus oxyacid and hydrogen halide:
Oxy Compounds of Phosphorus
Probably the most significant phosphorus compounds are those in which the element is combined with oxygen. Phosphorus(III) oxide (P4O6) is obtained by allowing white phosphorus to oxidize in a limited supply of oxygen. When oxidation takes place in the presence of excess oxygen, phosphorus(V) oxide (P4O10) forms. This compound is also readily formed by oxidation of P4O6. These two oxides represent the two most common oxidation states for phosphorus, +3 and +5. The structural relationship between P4O6 and P4O10 is shown in FIGURE 22.28. Notice the resemblance these molecules have to the P4 molecule (Figure 22.27); all three substances have a P4 core.
How do the electron domains about P in P4O6 differ from those about P in P4O10?
FIGURE 22.28 Structures of P4O6 (top) and P4O10 (bottom).
SAMPLE EXERCISE 22.8 Calculating a Standard Enthalpy Change
The reactive chemicals on the tip of a “strike anywhere” match are usually P4S3 and an oxidizing agent such as KClO3. When the match is struck on a rough surface, the heat generated by the friction ignites the P4S3, and the oxidizing agent brings about rapid combustion. The products of the combustion of P4S3 are P4O10 and SO2. Calculate the standard enthalpy change for the combustion of P4S3 in air, given the following standard enthalpies of formation: P4S3 (–154.4 kJ/mol), P4O10 (–2940.1 kJ/mol), SO2 (–296.9 kJ/mol).
Analyze We are given the reactants (P4S3 and O2 from air) and the products (P4O10 and SO2) for a reaction, together with their standard enthalpies of formation, and asked to calculate the standard enthalpy change for the reaction.
Plan We first need a balanced chemical equation for the reaction. The enthalpy change for the reaction is then equal to the enthalpies of formation of products minus those of reactants (Equation 5.31). We also need to recall that the standard enthalpy of formation of any element in its standard state is zero. Thus,
Solve The balanced chemical equation for the combustion is
Thus, we can write
Comment The reaction is strongly exothermic, making it evident why P4S3 is used on match tips.
Write the balanced equation for the reaction of P4O10 with water, and calculate ΔH° for this reaction using data from Appendix C.
Phosphorus(V) oxide is the anhydride of phosphoric acid (H3PO4), a weak triprotic acid. In fact, P4O10 has a very high affinity for water and is consequently used as a drying agent. Phosphorus(III) oxide is the anhydride of phosphorous acid (H3PO3), a weak diprotic acid (FIGURE 22.29).*
FIGURE 22.29 Structures of H3PO4 (top) and H3PO3 (bottom).
One characteristic of phosphoric and phosphorous acids is their tendency to undergo condensation reactions when heated. (Section 12.8) For example, two H3PO4 molecules are joined by the elimination of one H2O molecule to form H4P2O7:
Phosphoric acid and its salts find their most important uses in detergents and fertilizers. The phosphates in detergents are often in the form of sodium tripolyphosphate (Na5P3O10).
The phosphate ions “soften” water by binding their oxygen groups to the metal ions that contribute to the hardness of water. This keeps the metal ions from interfering with the action of the detergent. The phosphates also keep the pH above 7 and thus prevent the detergent molecules from becoming protonated.
Most mined phosphate rock is converted to fertilizers. The Ca3(PO4)2 in phosphate rock is insoluble (Ksp = 2.0 × 10–29). It is converted to a soluble form for use in fertilizers by treatment with sulfuric or phosphoric acid. The reaction with phosphoric acid yields Ca(H2PO4)2:
Although the solubility of Ca(H2PO4)2 allows it to be assimilated by plants, it also allows it to be washed from the soil and into bodies of water, thereby contributing to water pollution. (Section 18.4)
Phosphorus compounds are important in biological systems. The element occurs in phosphate groups in RNA and DNA, the molecules responsible for the control of protein biosynthesis and transmission of genetic information. It also occurs in adenosine triphosphate (ATP), which stores energy in biological cells and has the structure
The P — O — P bond of the end phosphate group is broken by hydrolysis with water, forming adenosine diphosphate (ADP):
This reaction releases 33 kJ of energy under standard conditions, but in the living cell, the Gibbs free energy change for the reaction is closer to –57 kJ/mol. The concentration of ATP inside a living cell is in the range of 1–10 mM, which means a typical human metabolizes her or his body mass of ATP in one day! ATP is continually made from ADP and continually converted back to ADP, releasing energy that can be harnessed by other cellular reactions.
CHEMISTRY AND LIFE
ARSENIC IN DRINKING WATER
“Arsenic,” meaning its oxides, has been known as a poison for centuries. The current Environmental Protection Agency (EPA) standard for arsenic in public water supplies is 10 ppb (equivalent to 10μg/L). Most regions of the United States tend to have low to moderate (2–10 ppb) groundwater arsenic levels (FIGURE 22.30). The western region tends to have higher levels, coming mainly from natural geological sources in the area. Estimates, for example, indicate that 35% of water-supply wells in Arizona have arsenic concentrations above 10 ppb.
FIGURE 22.30 Geographic distribution of arsenic in groundwater.
The problem of arsenic in drinking water in the United States is dwarfed by the problem in other parts of the world—especially in Bangladesh, where the problem is tragic. Historically, surface water sources in that country have been contaminated with microorganisms, causing significant health problems, including one of the highest infant mortality rates in the world. During the 1970s, international agencies, headed by the United Nations Children's Fund (UNICEF), began investing millions of dollars of aid money in Bangladesh for wells to provide “clean” drinking water. Unfortunately, no one tested the well water for the presence of arsenic; the problem was not discovered until the 1980s. The result has been the biggest outbreak of mass poisoning in history. Up to half of the country's estimated 10 million wells have arsenic concentrations above 50 ppb.
In water the most common forms of arsenic are the arsenate ion and its protonated hydrogen anions (AsO43–, HAsO42–, and H2AsO4–) and the arsenite ion and its protonated forms (AsO33–, HAsO32–, H2AsO3–, and H3AsO3). These species are collectively referred to by the oxidation number of the arsenic as arsenic(V) and arsenic(III), respectively. Arsenic(V) is more prevalent in oxygen-rich (aerobic) surface waters, whereas arsenic(III) is more likely to occur in oxygen-poor (anaerobic) groundwaters. In the pH range from 4 to 10, the arsenic(V) is present primarily as HAsO42– and H2AsO4–, and the arsenic(III) is present primarily as the neutral acid H3AsO3.
One of the challenges in determining the health effects of arsenic in drinking waters is the different chemistry of arsenic(V) and arsenic(III), as well as the different concentrations required for physiological responses in different individuals. In Bangladesh, skin lesions were the first sign of the arsenic problem. Statistical studies correlating arsenic levels with the occurrence of disease indicate a lung and bladder cancer risk arising from even low levels of arsenic.
The current technologies for removing arsenic perform most effectively when treating arsenic in the form of arsenic(V), so water treatment strategies require preoxidation of the drinking water. Once in the form of arsenic(V), there are a number of possible removal strategies. For example, Fe2(SO4)3 could be added to precipitate FeAsO4, which is then removed by filtration.
RELATED EXERCISE: 22.104.