THE FIRST SUSTAINED USE OF CHEMICALS AS AGENTS OF WARFARE

The talk and rhetoric of the late 19th century should have prepared the countries involved in World War I for chemical warfare. However, that was not case (Smart, 1997). World War I clearly demonstrated the deadly and destructive nature of chemicals in modern warfare. Both alliances in the war experimented with novel forms of warfare, to include chemical weapons, and followed the lead of their advisory (Hay, 2000). It is little wonder this war is known as the ‘‘chemist’s war’’ (Fitzgerald, 2008). Initially, the French used gas grenades with little effect and were followed by the German use of shells filled with tear gas (Joy, 1997). The Germans, capitalizing on their robust chemical industry, produced shells filled with dianisidine chlorosulfate (Smart, 1997). These shells were used in October of 1914 against the British at Neuve-Chapelle but had little effect. In the winter of 1914–15, the Germans fired 150 mm howitzer shells filled with xylyl bromide (Smart, 1997). The xylyl bromide shells were fired on both the eastern and western fronts with disappointing effects. Despite the inauspicious start of chemical warfare on both fronts, efforts were continued to develop new uses. It would soon be evident that chemical warfare would be devastating on the battlefield (Coleman, 2005; Tucker, 2006). Fritz Haber, a German scientist who later won the Nobel prize in Chemistry, had proposed the possibility of releasing chlorine gas from cylinders (Joy, 1997). Chemical warfare was attractive to Germans for two reasons: the shortage of German artillery shells and the ability to defeat the enemy trench system (Smart, 1997). After consideration and debate, the Germans released chlorine in April 1915 at Ypres, Belgium (Coleman, 2005). The German military was not prepared for the tremendous operational advantage the chlorine release provided. It did not take long for the British and French forces to respond in kind to the German offensive (Vedder and Walton, 1925; Joy, 1997; Smart, 1997; Coleman, 2005). In the fall of 1915, a British officer, William Livens, introduced a modified mortar (Figure 2.1) that could project gas-filled shells of chlorine or phosgene, the two agents of choice at that time (Joy, 1997). Both chlorine and phosgene caused extreme respiratory problems to those soldiers who were exposed (Vedder and Walton, 1925; Joy, 1997; Smart, 1997; Coleman, 2005; Hurst et al., 2007) (Figure 2.2). 

As the USA entered the war in the spring of 1917, an obvious concern of the military command was the effect of chemical warfare on standard operations. Chemistry departments at universities were tasked with investigating and developing novel chemical agents (Joy, 1997). Protective equipment (Figure 2.3) and basic studies of the biological effects of chemical agents were assigned to the US Army Medical Department (Joy, 1997). In the fall of 1917, the Army began to build an industrial base for producing 

chemical agents at Edgewood Arsenal, Maryland (Joy, 1997). As the effects of chlorine and phosgene became diminished by the advent of gas masks (Figure 2.4), the Germans turned to dichlorethyl sulfide (mustard) at Ypres against the British (Joy, 1997). As opposed to the gases, mustard remained persistent in the area and contact avoidance was the major concern (Joy, 1997). It is worth noting that almost 100 years after it was first used on the battlefield, mustard still has no effective treatment and research continues for effective therapeutics (Babin and Ricketts et al., 2000; Baskin and Prabhaharan, 2000; Casillas and Kiser, 2000; Hay, 2000; Schlager and Hart, 2000; Hurst et al., 2007; Romano et al., 2008). It has been estimated that there were over one million chemical casualties (Figure 2.5) of World War I with almost 8% being fatal (Joy, 1997). The Russians on the eastern front had a higher percentage of fatalities when compared with other countries in the war, primarily due to the later introduction of a protective mask (Joy, 1997). The relatively low mortality rate of chemical casualties in World War I demonstrated the most insidious aspect of their use, the medical and logistical burden it placed on the affected army. The eventual Allied victory brought a temporary end to chemical warfare. In 1919, the Treaty of Versailles prohibited the Germans from productio and use of chemical weapons.

 

 

Chemical properties of zinc, copper, vanadium, chromium, molybdenum, and cobalt

The chemical properties of the other essential transition elements simplify their transport properties. For zinc there is only the +2 oxidation state, and the hydrolysis of this ion is not a limiting feature of its solubility or transport. Zinc is an essential element for both animals and plants. 8,9,20,21 In general, metal ion uptake into the roots of plants is an extremely complex phenomenon. A crosssectional diagram of a root is shown in Figure 1.6. It is said that both diffusion

and mass flow of the soil solution are of significance in the movement of metal ions to roots. Chelation and surface adsorption, which -are pH dependent, also affect the availability of nutrient metal ions. Acid soil conditions in general retard uptake of essential divalent metal ions but increase the availability (sometimes with toxic results) of manganese, iron, and aluminum, all of which are normally of very limited availability because of hydrolysis of the trivalent ions. Vanadium is often taken up as vanadate, in a pathway parallel to phosphate. However, its oxidation state within organisms seems to be highly variable. Unusually high concentrations of vanadium occur in certain ascidians (the specific transport behavior of which will be dealt with later). The workers who first characterized the vanadium-containing compound of the tunicate, Ascidia nigra, coined the name tunichrome.  The characterization of the compound as a dicatecholate has been reported. 

 

Quite a different chemical environment is found in the vanadium-containing material isolated from the mushroom Amanita muscaria. Bayer and Kneifel, who named and first described amavadine,24 also suggested the structure shown in Figure 1. 7. 25 Recently the preparation, proof of ligand structure, and (by implication) proof of the complex structure shown in Figure 1.7 have been established.  Although the exact role of the vanadium complex in the mushroom

 

 

remains unclear, the fact that it is a vanadyl complex is now certain, although it may take a different oxidation state in vivo.
 

The role of chromium in biology remains even more mysterious. In human beings the isolation of "glucose tolerance factor" and the discovery that it contains chromium goes back some time. This has been well reviewed by Mertz, who has played a major role in discovering what is known about this elusive and apparently quite labile compound.27 It is well established that chromium is taken up as chromic ion, predominantly via foodstuffs, such as unrefined sugar, which presumably contain complexes of chromium, perhaps involving sugar hydroxyl groups. Although generally little chromium is taken up when it is administered as inorganic salts, such as chromic chloride, glucose tolerance in many adults and elderly people has been reported to be improved after supplementation with 150-250 mg of chromium per day in the form of chromic chloride. Similar results have been found in malnourished children in some studies in Third World countries. Studies using radioactively labeled chromium have shown that, although inorganic salts of chromium are relatively unavailable to mammals,

brewer's yeast can convert the chromium into a usable form; so l:irewer's yeast is today the principal source in the isolation of glucose tolerance factor and has been used as a diet supplement.
 

Although chromium is essential in milligram amounts for human beings as the trivalent ion, as chromate it is quite toxic and a recognized carcinogen. 30 The uptake-reduction model for chromate carcinogenicity as suggested by Connett and Wetterhahn is shown in Figure 1.8. Chromate is mutagenic in bacterial and mammalian cell systems, and it has been hypothesized that the difference between chromium in the +6 and +3 oxidation states is explained by the' 'uptake- reduction" model. Chromium(III), like the ferric ion discussed above, is readily hydrolyzed at neutral pH and extremely insoluble. Unlike Fe 3+ , it undergoes extremely slow ligand exchange. For both reasons, transport of chromium( III) into cells can be expected to be extremely slow unless it is present as specific complexes; for example, chromium(III) transport into bacterial cells has been reported to be rapid when iron is replaced by chromium in the siderophore iron-uptake mediators. However, chromate readily crosses cell membranes and enters cells, much as sulfate does. Because of its high oxidizing power, chromate can undergo reduction inside organelles to give chromium(m), which binds to small molecules, protein, and DNA, damaging these cellular components.

In marked contrast to its congener, molybdenum is very different from chromium in both its role in biology and its transport behavior, again because of fundamental differences in oxidation and coordination chemistry properties. In contrast to chromium, the higher oxidation states of molybdenum dominate its chemistry, and molybdate is a relatively poor oxidant. Molybdenum is an essential element in many enzymes, including xanthine oxidase, aldehyde reductase, and nitrate reductase. 19 The range of oxidation states and coordination geometries of molybdenum makes its bioinorganic chemistry particularly interesting and challenging.
 

The chemistry of iron storage and transport is dominated by high concentrations, redox chemistry (and production of toxic-acting oxygen species), hydrolysis (pKa is about 3, far below physiological pH), and insolubility. High-affinity chelators or proteins are required for transport of iron and high-capacity sequestering protein for storage. By comparison to iron, storage and transport of the other metals are simple. Zinc, copper, vanadium, chromium, manganese, and molybdenum appear to be transported as simple salts or loosely bound protein complexes. In vanadium or molybdenum, the stable anion, vanadate or molybdate, appears to dominate transport. Little is known about biological storage of any metal except iron, which is stored in ferritin. However, zinc and copper are bound to metallothionein in a fonn that may participate in storage.