Group VA

Submitted by ChemPRIME Staff on Thu, 12/16/2010 - 14:34

Although all the elements in this groupThose elements that comprise a single column of the periodic table. Also called family. form compounds in which their oxidation stateA formally defined charge that an atom in a compound or ion would have if the compound or ion consisted entirely of monatomic ions. Based on a Lewis diagram, the charge that an atom would have if all electrons in bonds were assigned to the more electronegative atom or divided equally between atoms of the same electronegativity. of +5 equals the group number, their other properties vary considerably. Nitrogen is clearly nonmetallic and consists of diatomicContaining two atoms per molecule. triply-bonded N2 molecules.

Phosphorus, also a nonmetal, exists as tetrahedral P4 molecules(Fig. 1a) in the vaporThe gaseous state of a substance that typically exists as a liquid or solid; a gas at a temperature near or below the boiling point of the corresponding liquid. and the white allotropic form of the solidA state of matter having a specific shape and volume and in which the particles do not readily change their relative positions.. On standing, white phosphorus slowly changes to the red allotropeOne of two or more different structural forms for an element that exist in the same physical state at the same temperature and pressure., whose structure is shown in Fig. 1c. The most stable form of the element is black phosphorus, which has a layer structure (Fig. 1b). Black phosphorus can be made by heating the white form with a mercury catalystA substance that increases the rate of a chemical reaction but that undergoes no net change during the reaction. for 8 days at 220 to 370°C.

Figure 1 Allotropic forms of phosphorus. (a) P4 molecule of phosphorus vapor and solid white phosphorus; (b) layer structure of black phosphorus; (c) chain (polymeric) structure of red phosphorus. Note that in all structures each phosphorus atom forms three bonds.


Arsenic is a semimetal and consists of As4 molecules in the gasA state of matter in which a substance occupies the full volume of its container and changes shape to match the shape of the container. In a gas the distance between particles is much greater than the diameters of the particles themselves; hence the distances between particles can change as necessary so that the matter uniformly occupies its container. phase. When As4(g) is condensed to a solid, three allotropes may form. The most stable of these is metallic arsenic, in which each arsenic atom has three nearest neighbors, with three more arsenic atomsThe smallest particle of an element that can be involved in chemical combination with another element; an atom consists of protons and neutrons in a tiny, very dense nucleus, surrounded by electrons, which occupy most of its volume. somewhat farther away.

Antimony, also a semimetal, has two allotropes, the more stable one being metallic, like arsenic. In the case of bismuth, only the metallic form occurs. Note that for all the group VA elements the 8 – N rule is followed. The number of bonds or nearest neighbors for each atom is 8 minus the group number in Image:N-Ntriple bond.jpg, P4, and even in the metallic forms of As, Sb, and Bi. The table summarizes the atomic properties of the group VA elements. Overall, the trends are what we would expect, based on our experience with previous groups. These elements exhibit a much wider variety of oxidationThat part of a chemical reaction in which a reactant loses electrons; simultaneous reduction of a reactant must occur. states, however, especially in the case of nitrogen. This element forms compounds in which it has every possible oxidation numberA formally defined charge that an atom in a compound or ion would have if the compound or ion consisted entirely of monatomic ions. Based on a Lewis diagram, the charge that an atom would have if all electrons in bonds were assigned to the more electronegative atom or divided equally between atoms of the same electronegativity. from –3 (the group number minus 8) to +5 (the group number). As in previous groups, the oxidation state in which the ns2 pair of electrons is not used for bonding becomes more prominent toward the bottom of the periodic tableA chart showing the symbols of the elements arranged in order by atomic number and having chemically related elements appearing in columns.. There are a few compounds, Bi(NO3)3, for example, in which discrete Bi3+ ions are present.


Properties of the Group VA Elements.

Element Symbol Electron

Configuration

Usual

Oxidation State

Radius/pm


Covalent Ionic (Charge)


Nitrogen N [He]2s22p3 +5, +3, -3 70 (3-)171


Phosphorus P [Ne]3s23p3 +5, +3, -3 110 -


Arsenic As [Ar]4s23d104p3 +5, +3 121 -


Antimony Sb [Kr]5s24d105p3 +5, +3 141 -


Bismuth Bi [Xe]6s24f145d106p3 +5, +3 146 (3+)108



Symbol Ionization Energy/MJ mol–1 Density/

g cm–3

Electro-

negativity

Melting

Point (in °C)

First Second Third Fourth Fifth
N 1.407 2.862 4.585 7.482 9.452 1.25×10-3 3.0 -210


P 1.018 1.909 2.918 4.963 6.280 1.82 2.1 44
As 0.953 1.804 2.742 4.843 6.049 5.72 2.0 817
Sb 0.840 1.601 2.450 4.271 5.403 6.69 1.9 631
Bi 0.710 1.616 2.472 4.380 5.417 9.80 1.9 271


Chemical Reactions and Compounds

The most important compounds of the group VA elements are those of nitrogen and phosphorus. Both elements are essential to all living organisms, and both are progressively removed from soil when plants are cultivated and crops harvested. According to Liebig’s law of the minimum, an insufficient supply of either element can limit plant growth and reduce crop yields, and so these elements are important components of fertilizer. More recently both elements have been implicated in several kinds of pollutionThe contamination of the air, water, and earth by personal, industrial, and farm waste. problems. As we discuss the properties of nitrogen and phosphorus compounds, their effects on food production and environmental degradation will also be discussed.


Nitrogen..The importance of nitrogen fertilizer was first recognized over a century ago. By the late 1800s the only major ore of nitrogen, Chile saltpeter, NaNO3, was being mined in Chile and shipped to Europe for application to agricultural land, but the supply was obviously limited. Most nitrogen at the earth’s surface is in the form of N2(g), which makes up 78 percent of the atmosphereA unit of pressure equal to 101.325 kPa or 760 mmHg; abbreviated atm. Also, the mixture of gases surrounding the earth. by volume (or by amount of substanceA material that is either an element or that has a fixed ratio of elements in its chemical formula.). Therefore chemists began to look for ways of obtaining nitrogen compounds directly from the atmosphere. Any process which does this is called nitrogen fixation.

Nitrogen fixation can occur naturally when an electrical discharge (lightning) heats air to a high temperatureA physical property that indicates whether one object can transfer thermal energy to another object.. The following reaction occurs:


N2(g) + O2(g) → 2NO(g)      (1)


The nitrogen monoxide (nitric oxide) formed can react further at ordinary temperatures, producing the brown gas, nitrogen dioxide:


2NO(g) + O2(g) → 2NO2(g)      (2)


The Lewis diagrams for these and other important nitrogen compounds are shown in Fig. 2.


Figure 2 Lewis diagrams for important nitrogen compounds.

From the figure you can see that both NO and NO2 have an odd number of electrons and violate the octet ruleThe generalization that atoms tend to gain or lose electrons until there is an octet in the outermost electron shell.. In such a case it is common for two molecules to combine (dimerize) by pairing their odd electrons. In the case of NO2, dimerization occurs below room temperature, producing colorless dinitrogen tetroxide:


2NO2(g) \rightleftharpoons N2O4(g)


At room temperature, however, the NO2 and N2O4 are in equilibriumA state in which no net change is occurring, that is, in which the concentrations of reactants and products remain constant; chemical equilibrium is characterized by forward and reverse reactions occurring at the same rate., as evidenced by the brown color of the mixtureA combination of two or more substances in which the substances retain their chemical identity.. NO dimerizes only at very low temperatures in the solid state.

The first industrial nitrogen fixation was done by mimicking nature.

Reaction (1) was carried out in a plant near Niagara Falls, where hydroelectric generation provided inexpensive power to support an electric arc. NO was further oxidized to NO2 which was dissolved in H2O to convert it to nitric acidIn Arrhenius theory, a substance that produces hydrogen ions (hydronium ions) in aqueous solution. In Bronsted-Lowry theory, a hydrogen-ion (proton) donor. In Lewis theory, a species that accepts a pair of electrons to form a covalent bond., HNO3:


3NO2(g) + H2O(l) → 2H+(aq) + 2NO3(aq) + NO(g)      (3)


Note that NO2 is not the acid anhydrideA substance that reacts with water to form an acid or a base. A substance from which water has been removed. of HNO3. This reaction involves, dis-proportionation of NO2 (which contains N in the +4 oxidation state) to form HNO3 in the + 5 state) and NO (N in the + 2 state). The NO can be recycled by reoxidizing it to NO2, and so it was not wasted. The HNO3 produced in Eq. (12.5) was neutralized with NaOH to make a substitute for Chile saltpeter:


NaOH(aq) + HNO3(aq) → NaNO3(aq) + H2O(l)


Fixation of nitrogen by the electric-arc process used a great deal of energy and was rather expensive. Other methods were designed to replace it, and the most successful of these is the Haber process, which is the major one used today. Nitrogen is reacted with hydrogen at a high temperature and extremely high pressureForce per unit area; in gases arising from the force exerted by collisions of gas molecules with the wall of the container. over a catalyst consisting of iron and aluminum oxide:

N2(g) + 3H2(g) Image:Haber Process Conditions.jpg 2NH3(g)


The ammonia produced by the Haber process is used directly as a fertilizer. It can be liquefied under pressure and injected through special nozzles about a foot under the soil surface. This prevents loss of gaseous ammonia which would otherwise irritate the nose, throat, and lungs of anyone near a fertilized field. You are probably familiar with the odor of ammonia since it is the most common weak baseAn base that ionizes only partially in a given solvent. encountered in the chemical laboratory.

Prior to the recent development of underground injection techniques, most ammonia was converted to ammonium nitrate for fertilizer use:


NH3(aq) + HNO3(aq) → NH4NO3(aq)      (4)


Except for ammonia, ammonium nitrate contains a greater massA measure of the force required to impart unit acceleration to an object; mass is proportional to chemical amount, which represents the quantity of matter in an object. fraction of nitrogen than any other compound of comparable cost. Ammonium nitrate manufacture requires that half the ammonia produced in the Haber process be converted to nitric acid. The first step is oxidation of ammonia over a catalyst of platinum metal:


4NH3(g) + 5O2(g) \xrightarrow[\text{800}{}^\circ \text{C}]{\text{Pt}} 4NO(g) + 6H2O(g)


This is called the Ostwald process. It is followed by Eqs. (2) and (3), yielding nitric acid, which can be combined with ammonia [Eq. (4)].

Nitric acid and nitrates have commercial applications other than fertilizer production. Because NO3 is a strong oxidizing agentA chemical species that accepts electrons in order to oxidize another species. In the process the oxidizing agent is itself reduced., it reacts vigorously with substances whose elements are in low oxidation states. One example of this is black powder, which consists of charcoal (carbon), sulfur, an potassium nitrate, KNO3 (saltpeter or nitre). During the American revolution, for example, both armies had numerous persons whose job was to find caves in which the relatively solubleAble to dissolve in a solvent to a significant extent. KNO3 had been deposited as water evaporated. A second example is nitroglycerin which contains carbon and hydrogen in low oxidation states as well as nitrate. Still another example of an explosive nitrate is NH4NO3, which contains nitrogen in its maximum and minimum oxidation state. NH4NO3 decomposes as follows:


NH4NO3(s) \xrightarrow[\text{shock}]{\text{heat or}} N2O(g) + 2H2O(g)      ΔHm = –37 kJ mol–1


The reaction is exothermicDescribes a process in which energy is transferred to the surroundings as a result of a temperature difference. and produces 3 mol of gaseous products for every mole of solid reactantA substance consumed by a chemical reaction.. This causes a tremendous increase in pressure, and, if the reaction is rapid, an explosion.

The compound dinitrogen monoxide (nitrous oxide or laughing gas), produced by decomposition of NH4NO3, is a third important oxide of nitrogen (in addition to NO and NO2). N2O is produced during microbial decomposition of organicRefers to the branch of chemistry that studies compounds containing carbon, usually in combination with hydrogen and other elements such as O, N, S, and P. Certain small ions and compounds containing carbon (such as carbonate ions and carbon dioxide) are not considered to be organic, but rather are classed as inorganic. matterAnything that occupies space and has mass; contrasted with energy. containing nitrogen. Because it is quite unreactive, it is the second most-concentratedIncreased the concentration of a mixture or solution (verb). Having a large concentration (adjective). nitrogen-containing substance in the atmosphere (after N2). It is used commercially as an anesthetic, is mildly intoxicating, and is poisonous in large doses.

The other two important oxides of nitrogen, NO and NO2, play a major role in an air-pollution problem known as photochemical smog (or Los Angeles smog). NO is formed by Eq. (1) in automobile engines and other high-temperature combustionVigorous combination of a material with oxygen gas, usually resulting in a flame. processes. At normal temperatures NO is oxidized to NO2 [Eq. (2)]. Both these oxides are [[CoreChem:Free Radicals|free radicals]] and are rather reactive. Moreover, brown-colored NO2 absorbs sunlight, and the energy of the absorbed photons breaks a nitrogen-oxygen bond:


Image:NO2 decomposition.jpg


The oxygen atoms produced are highly reactive. They combine with hydrocarbonA compound containing only the elements carbon and hydrogen. molecules (from evaporated or unburned gasoline) to form aldehydes, ketones, and a number of other compounds which form an almost fog-like cloud and irritate the eyes, throat, and lungs. Photochemical smog is especially bad in cities like Los Angeles and Denver which have lots of sunshine and automobile traffic, but its effects have been observed in every large city in the United States.


Phosphorus    As in the case of carbon and silicon, there are major differences between the chemistries of nitrogen and phosphorus. The concentrations of phosphorus compounds in the earth’s atmosphere are so small as to be negligible, but phosphorus is more abundant than nitrogen in the solid crust. Here it is found as phosphate rock, which is mainly hydroxyapatite, Ca10(PO4)6(OH)2, or fluorapatite, Ca10(PO4)6F2. (These are the same substances involved in the discussion of dental decay in the section on group IVA elements.)

Phosphate rock is quite insolubleUnable to dissolve appreciably in a solvent., and hence its phosphate ions cannot be assimilated by plants. Production of phosphate fertilizer requires treatment of apatite acid. This protonates the PO43 – ions, converting them to H2PO4, whose calcium salt is much more soluble:


Ca10(PO4)6(OH)2 + 7H2SO4 + H2O → 3Ca(H2PO4)2•H2O + 7CaSO4


Ca10(PO4)6(OH)2 + 14H3PO4 → 10Ca(H2PO4)2 + 2H2O


The compound Ca(H2PO4)•H2O is known as superphosphate, and Ca(H2PO4)2 is called triple superphosphate.

The phosphoric acid, H3PO4, used to make triple superphosphate is also obtained from phosphate rock. The first step is a reductionThat part of a chemical reaction in which a reactant gains electrons; simultaneous oxidation of a reactant must occur. with carbon (coke) and silicon dioxide in an electric furnace:


2Ca10(PO4)6(OH)2 + 18SiO2 + 30C → 3P4 + 30CO + 2Ca(OH)2 + 18CaSiO3


The phosphorus obtained this way is then oxidized to phosphorus pentoxide:


P4(s) + 5O2(g) → P4O10(s)


(The name phosphorus pentoxide for P4O10 comes from the empirical formulaThe chemical formula of a substance written using the smallest possible integer subscripts that reflect the elemental composition. P2O5 of this compound.) Phosphorus pentoxide is the acid anhydride of phosphoric acid:.


P4O6(s) + 6H2O(l) → 4H3PO3(aq)


Although not a very strong acidAn acid that ionizes completely in a particular solvent., phosphoric acid is triprotic. Therefore, 1 mol of this acid can transfer 3 mol of protons to a strong baseA base that dissociates completely or ionizes completely in a particular solvent.. There is another oxide of phosphorus, P4O6, which involves the + 3 oxidation state, corresponding to use of the 3p3, but not the 3s2, electrons for bonding. P4O6 is the acid anhydride of phosphorous acid, H3PO3:


P4O6(s) + 6H2O(l) → 4H3PO3(aq)


Phosphorous acid is quite weak, and, contrary to what its formula might suggest, can only donate two protons. This is apparently because its Lewis structureA representation of the structure of a covalent molecule that uses dots to indicate for each atom how many electrons ae shared with other atoms (form bonds) and how many are unshared; also called electron dot structure. is


Image:Phosphoric Acid.jpg


Only the two protons bonded to highly electronegative oxygen atoms are expected to be acidic.

Another major commercial use of phosphates is in laundry detergents. The problem of precipitationThe formation of a solid within a solution, often by the combination of cations and anions to form an insoluble ionic compound. of soapA salt of a fatty acid produced by the saponification of fat. by hard-water ions such as Ca2+ was mentioned in the section on alkaline earth metals. This can be prevented, and the cleaning power of syntheticDescribes a substance that has been manufactured‐one that has not been created by natural processes. detergents can be improved, by adding phosphates. The compound usually used is sodium tripolyphosphate, whose anionA negatively charged ion. An ion that is attracted toward the anode in an electrolytic cell. is a condensation polymerA polymer formed by condensation reactions between monomers and the growing polymer chain; in a condensation polymerization a small molecule each time a monomer becomes bonded to the polymer chain. of hydrogen phosphate and dihydrogen phosphate ions:


Image:Tripolyphosphate.jpg


The tripolyphosphate ion has numerous O atoms whose lone pairs of electrons can form coordinate covalent bonds to metal ions like Ca2+:


Image:Coordinate covalent bonds.jpg


The Ca2+ ions are effectively removed from solutionA mixture of one or more substances dissolved in a solvent to give a homogeneous mixture. (they are said to be sequestered) because they are bonded to the tripolyphosphate ion. Consequently Ca2+(aq) is not available to precipitate soap or detergent molecules.

The use of phosphates in detergents is responsible in part for an environmental problem known as accelerated eutrophication, or premature, aging of bodies of water. Over a periodThose elements from a single row of the periodic table. of many thousands of years, a lake or other body of water will slowly accumulate essential nutrient elements such as nitrogen or phosphorus because their compounds dissolve in streams that feed the lake. As the water becomes richer in nutrients, more plants and microorganisms can grow. Some of the organic matter which remains when these organisms die precipitates to the bottom of the lake and is not decomposed. Eventually the lake fills up with debris, becoming a swamp, and finally dry land.

This process of eutrophication can be greatly accelerated by human input of nutrients such as nitrogen or phosphorus fertilizers, or phosphates from detergents. Since reduction in the use of detergent phosphates would appear to have the least negative effects—people’s clothes might not look as clean—many have suggested that prohibiting or limiting phosphate content is the way to solve the problem. Many states have passed laws implementing such limitations or bans.