Environmental Impact and Health Effects of the Alkali Metals
Lithium
Like all alkali metals, lithium reacts easily in water and does not occur freely in nature due to its activity. Lithium is a moderately abundant element and its present in the Earth’s crust in 65 ppm (parts per million). Lithium is easily adsorbed by plants and the amount of lithium in plants varies widely. While the lithium surface becomes coated with a mixture of lithium hydroxide, lithium carbonate, and lithium nitride (Li3N), lithium hydroxide represents a potentially significant hazard because it is extremely corrosive. Special attention should be given to water organisms.
Many reactions of lithium may cause fire or explosion when exposed to Lithium. It gives off irritating or toxic fumes (or gases) in a fire. There is risk of fire and explosion on contact with combustible substances and water. On inhalation it gives burning sensation, coughs, labored breathing, shortness of breath, and sore throat. However, the symptoms may be delayed. When it comes in contact of the skin, the skin becomes red. On ingestion, there are chances of abdominal cramps, abdominal pain, burning sensation, nausea, shock or collapse, vomiting, and weakness. The substance is corrosive to the eyes, the skin and the respiratory tract. Inhalation of the substance may cause lung oedema. The symptoms of lung oedema are often not manifested until a few hours have passed and they are aggravated by physical effort.
Rest and medical observation is, therefore, essential in all cases of probable actions of lithium. Immediate administration of an appropriate spray, by a doctor or a person authorized by him/her, should be considered. The substance can be absorbed into the body by inhalation of its aerosol and by ingestion. Inhalation due to evaporation at 20°C is negligible; a harmful concentration of airborne particles can, however, be reached quickly when dispersed. Heating the element may cause violent combustion or explosion. The substance may spontaneously ignite on contact with air when finely dispersed. Upon heating, toxic fumes are formed. It reacts violently with strong oxidants, acids and many compounds causing fire and explosion hazards. It also reacts violently with water, forming highly flammable hydrogen gas and corrosive fumes of lithium hydroxide.
Sodium is the most abundant element in the Earth’s crust and is found in nature only in the combined state ranking sixth or seventh in order of abundance of the elements. It occurs in the ocean and in salt lakes as sodium chloride, NaCl, and less often as sodium carbonate, Na2CO3, and sodium sulfate, Na2SO4. Sodium is the second most abundant element after chlorine (as chloride ions) dissolved in seawater. The most important sodium salts found in nature are sodium chloride (halite or rock salt), sodium carbonate (trona or soda), sodium borate (borax), sodium nitrate and sodium sulfate. Sodium salts are found in all water bodies. A huge amount of this salt is extracted mainly from salt deposits by pumping water down bore holes to dissolve it and pumping up brine. The sun and many other stars shine with visible light in which the yellow component dominates and this is given out by sodium atoms in a high-energy state. Sodium’s powdered form is highly explosive in water and a poison combined and uncombined with many other elements. This chemical is not mobile in solid form, although it absorbs moisture very easily. Once liquid, sodium hydroxide leaches rapidly into the soil
Sodium salts are important ingredients of many foodstuffs (for instance common salt) as besides imparting salty taste, it is necessary for humans to maintain the balance of the physical fluids system. Sodium is also required for nerve and muscle functioning. However, too much sodium can damage our kidneys and increases the chances of high blood pressure. The amount of sodium a person consumes each day varies from individual to individual and from culture to culture. Some people get as little as 2 g/day, some as much as 20 grams. Sodium is essential, but controversially surrounds the amount required. Contact of sodium with water, including perspiration causes the formation of sodium hydroxide fumes, which are highly irritating to skin, eyes, nose and throat. This may cause sneezing and coughing. Very severe exposures may result in difficult breathing, coughing and chemical bronchitis. Contact to the skin may cause itching, tingling, thermal and caustic burns and permanent damage. Its contact with eyes may result in permanent damage and loss of sight.
Most potassium occurs in the Earth’s crust as minerals, such as feldspars and clays. Potassium is leached from these by weathering, which explains why there is quite a lot of this element in the sea (0.75 g/liter). Minerals mined for their potassium are pinkish and sylvite, carnallite and alunite. The main mining area used to be Germany, which had a monopoly of potassium before the First World War. Today most potassium minerals come from Canada, USA and Chile. The world production of potassium ores is about 50 million tonnes, and reserves are vast. Potassium is a key plant element. Although it is soluble in water, little is lost from undisturbed soils because as it is released from dead plants and animal excrements, it quickly become strongly bound to clay particles, and it is retained ready to be readsorbed by the roots of other plants.
Potassium can be found in vegetables, fruit, potatoes, meat, bread, milk and nuts. It plays an important role in the physical fluid system of humans and assists nerve functions. Potassium, as the ion K+, concentrate inside cells, and 95% of the body’s potassium is so located. When our kidneys are somehow malfunctioning, an accumulation of potassium will take place. This can lead to disturbing heartbeats. Potassium can affect us when breathed in. Inhalation of dust or mists can irritate the eyes, nose, and throat, lungs with sneezing, coughing and sore throat. Higher exposures may cause a build up of fluid in the lungs, this can cause death. Skin and eye contact can cause severe burns leading to permanent damage.
Rubidium is a widely distributed element, ranking about 16th in order of abundance of the elements in Earth’s crust. The relative abundance of rubidium has been reassessed in recent years and it is now suspected of being more plentiful than previously calculated. It is not found in large deposits but occurs in small amounts in certain mineral waters and in many minerals usually associated with other alkali metals. It is also found in small quantities in tea, coffee, tobacco, and other plants, and trace quantities of the element may be required by living organisms. Rubidium is used in making certain catalysts. The rate of radioactive decay of the rubidium-87 can be used in geologic age determination. It is very like potassium and there are no environments where it is seen as a threat. No minerals of rubidium are known, but rubidium is present in significant amounts in other minerals such as lepodite (1.5%), pollucite and carnallite. It is also present in traces in trace amounts in other minerals such as zinnwaldite and leucite. The amount of rubidium produced every year is small, and what demand there is can be met from a stock of a mixed carbonate by-product that is collected during the extraction of lithium from lepodite. The little rubidium that is produced is used for research purposes only, these is no incentive to seek commercial outlets for the material.
Rubidium has no known biological role but has a slight stimulatory effect on metabolism, probably because it is like potassium. The two elements are found together in minerals and soils, although potassium is much more abundant than rubidium. Plant will adsorb rubidium quite quickly. When stresses by deficiency of potassium some plants, such as sugar beet, will respond to the addition of rubidium. In this way rubidium enters the food chain and so contributes to a daily intake of between 1 and 5 mg. No negative environmental effects have been reported.
It is moderately toxic by ingestion. If rubidium ignites, it will cause thermal burns. Rubidium readily reacts with skin moisture to form rubidium hydroxide, which causes chemical burns of eyes and skin. Signs and symptoms of overexposure to this element are skin and eye burns, failure to gain weight, ataxia, hyper irritation, skin ulcers, and extreme nervousness. Medical condition is aggravated by exposure to heart patients due to potassium imbalance. In case of exposure to eyes immediately flush with running water for 15 minutes while holding eyelid. Obtain medical attention immediately. In case of skin exposure remove material and flush with soap and water. Remove contaminated clothing. Get medical attention promptly. In case of inhalation move to fresh air immediately. If irritation persists, get medical attention. In case of ingestion do not induce vomiting. Rather try to get medical attention immediately.
Although cesium is much less abundant than the other alkali metals, it is still more common than elements many important elements. Few cesium mineral are know, pollucite is the main: they are silicate magmas cooled from granites. Cesium ranks about 46th in natural abundance among the elements in crustal rocks. The natural source yielding the greatest quantity of cesium is the rare mineral Pollux (or pollucite). Ores of this mineral found on the island of Elba contain 34 percent of cesium oxide; American ores of Pollux, found in Maine and South Dakota, contain 13 percent of the oxide. It is extracted by separating the cesium compound from the mineral, transforming the compound thus obtained into the cyanide, and electrolysis of the fused cyanide. Cesium can also be obtained by heating its hydroxides or carbonates with magnesium or aluminum and by heating its chlorides with calcium. Commercial cesium usually contain such elements with which it usually occurs in minerals and which resembles it so closely that no effort is made to separate them.
Cesium occurs naturally in the environment mainly from erosion and weathering of rocks and minerals. It is also released into the air, water and soil through mining and milling of ores. Radioactive isotopes of cesium may be released into the air by nuclear power plants and during nuclear accidents and nuclear weapons testing. The radioactive isotopes can only be decreased in concentration through radioactive decay. Non-radioactive cesium can either be destroyed when it enters the environment or react with other compounds into very specific molecules. Both radioactive and stable cesium act the same way within the bodies of humans and animals chemically. Cesium in air can travel long distances before settling on earth. In water and soils most cesium compounds are very water-soluble. In soils, however, cesium does not rinse out into the groundwater. It remains within the top layers of soils as it strongly bonds to soil particles and as a result it is not readily available for uptake through plant roots. Radioactive cesium does have a chance of entering plants by falling on leaves. Animals that are exposed to very high doses of cesium show changes in behavior, such as increased or decreased activity.
Humans may be exposed to cesium by breathing, drinking or eating. In air the levels of cesium are generally low, but radioactive cesium has been detected at some level in surface water and in many types of foods. The amount of cesium in foods and drinks depends upon the emission of radioactive cesium through nuclear power plants, mainly through accidents. These accidents have not occurred since the Chernobyl disaster in 1986. People that work in the nuclear power industry may be exposed to higher levels of cesium, but many precautionary measurements can be taken to prevent this. It is not very likely that people experience health effects that can be related to cesium itself. When contact with radioactive cesium occurs, which is highly unlikely, a person can experience cell damage due to radiation of the cesium particles. Due to this, effects such as nausea, vomiting, diarrhoea and bleeding may occur. When the exposure lasts a long time people may even lose consciousness. Coma or even death may than follow. How serious the effects are depends upon the resistance of individual persons and the duration of exposure and the concentration a person is exposed to.
Francium
Francium is the heaviest of the alkali metals and the most electropositive or the least electronegative of all the known elements. All its isotopes are radioactive and short-lived. The element is extremely rare, though its atoms have been detected in uranium ores. It is because of its extreme rarity that its chemical and physical properties are not known. It has been studied by radiochemical techniques, which show that it’s most stable state is the ion Fr+. Francium is. Francium is the second rarest element in the crust, after astatine.
No use has been found for what little francium can be produced. No doubt Francium occurs naturally to a very limited extent in uranium minerals. Nevertheless it has been estimated that there might be from 340 to 550 grams of francium in the earth’s crust at any one time. Due to its extremely short half-life, there’s no reason for considering the effects of francium in the environment. As it is so unstable, any amount formed would decompose to other elements so quickly that there’s no reason to study its effects on human health.
About the Author
Dr.Badruddin Khan teaches Chemistry in the University of kashmir, srinagar, India.
The Pulse 3.33: Celiac Disease
