Chemical element of potassium

Educational institution

Summary

Subject: Potassium

Content

1. Introduction

2. Occurrence

3. History of the free element

4. Production

5. Isotopes

6. Properties

6.1 Physical

6.2 Chemical

7. Potassium cations in the body

7.1 Biochemical function

7.2 Membrane polarization

7.3 Filtration and excretion

7.4 Potassium in the diet and by supplement

7.4.1 Adequate intake

7.4.2 Optimal intake

7.4.3 Medical supplementation and disease

8. Applications

8.1 Biological applications

8.2 Food applications

8.3 Industrial applications

9. Precautions

10. Conclusion

11. References

1. Introduction

Potassium (pronounced /p??t?si?m/ po-TAS-ee-?m) is the chemical element with the symbol K (Latin: kalium, from Arabic: ЗбЮубънуе? al-qalyah “plant ashes”, cf. Alkali from the same root), atomic number 19, and atomic mass 39.0983. Potassium was first isolated from potash. Elemental potassium is a soft silvery-white metallic alkali metal that oxidizes rapidly in air and is very reactive with water, generating sufficient heat to ignite the evolved hydrogen.

Potassium in nature occurs only as ionic salt. As such, it is found dissolved in seawater, and as part of many minerals. Potassium ion is necessary for the function of all living cells, and is thus present in all plant and animal tissues. It is found in especially high concentrations in plant cells, and in a mixed diet, it is most highly concentrated in fruits.

In many respects, potassium and sodium are chemically similar, although they have very different functions in organisms in general, and in animal cells in particular.

Figure. 0. Silvery white

2. Occurrence

Elemental potassium does not occur in nature because it reacts violently with water. As various compounds, potassium makes up about 1.5% of the weight of the Earth's crust and is the seventh most abundant element. As it is very electropositive and highly reactive potassium metal is difficult to obtain from its minerals.[1]

Figure. 1. Potassium in feldspar

3. History of the free element

Elemental potassium was not known in Roman times, and its names are not Classical Latin but rather neo-Latin. The name kalium was taken from the word "alkali", which came from Arabic al qalоy = "the calcined ashes". The name potassium was made from the word "potash", which is English, and originally meant an alkali extracted in a pot from the ash of burnt wood or tree leaves. The structure of potash was not then known,[when?] but is now understood to be mostly potassium carbonate. By heating, the carbonate could be freed of carbon dioxide, leaving "caustic potash", so called because it caused chemical burns in contact with human tissue.

Potassium metal was discovered in 1807 in England by Sir Humphry Davy, who derived it from caustic potash (KOH), by the use of electrolysis of the molten salt with the newly discovered voltaic pile. Before the 18th century, no distinction was made between potassium and sodium. Potassium was the first metal that was isolated by electrolysis.[2] Davy extracted sodium by a similar technique, demonstrating the elements to be different.[3]

4. Production

Pure potassium metal may be isolated by electrolysis of its hydroxide in a process that has changed little since Davy.[1] Thermal methods also are employed in potassium production, using potassium chloride

Potassium salts such as carnallite, langbeinite, polyhalite, and sylvite form extensive deposits in ancient lake and seabeds, making extraction of potassium salts in these environments commercially viable. The principal source of potassium, potash, is mined in Saskatchewan, California, Germany, New Mexico, Utah, and in other places around the world. It is also found abundantly in the Dead Sea. Three thousand feet below the surface of Saskatchewan are large deposits of potash which are important sources of this element and its salts, with several large mines in operation since the 1960s. Saskatchewan pioneered the use of freezing of wet sands (the Blairmore formation) in order to drive mine shafts through them. The main mining company is the Potash Corporation of Saskatchewan. The oceans are another source of potassium, but the quantity present in a given volume of seawater is relatively low compared with sodium.[4][5]

5. Isotopes

There are 24 known isotopes of potassium. Three isotopes occur naturally: 39K (93.3%), 40K (0.0117%) and 41K (6.7%). Naturally occurring 40K decays to stable 40Ar (11.2% of decays) by electron capture or positron emission, or decays to stable 40Ca (88.8% of decays) by beta decay; 40K has a half-life of 1.250?109 years. The decay of 40K to 40Ar enables a commonly used method for dating rocks. The conventional K-Ar dating method depends on the assumption that the rocks contained no argon at the time of formation and that all the subsequent radiogenic argon (i.e., 40Ar) was quantitatively retained. Minerals are dated by measurement of the concentration of potassium and the amount of radiogenic 40Ar that has accumulated. The minerals that are best suited for dating include biotite, muscovite, plutonic/high grade metamorphic hornblende, and volcanic feldspar; whole rock samples from volcanic flows and shallow instrusives can also be dated if they are unaltered.

Outside of dating, potassium isotopes have been used extensively as tracers in studies of weathering. They have also been used for nutrient cycling studies because potassium is a macronutrient required for life.

40K occurs in natural potassium (and thus in some commercial salt substitutes) in sufficient quantity that large bags of those substitutes can be used as a radioactive source for classroom demonstrations. In healthy animals and people, 40K represents the largest source of radioactivity, greater even than 14C. In a human body of 70 kg mass, about 4,400 nuclei of 40K decay per second.[6] The activity of natural potassium is 31 Bq/g.

6. Properties

6.1 Physical

Potassium is the second least dense metal; only lithium is less dense. It is a soft, low-melting solid that can easily be cut with a knife. Freshly cut potassium is silvery in appearance, but in air it begins to tarnish toward grey immediately.[1]

In a flame test, potassium and its compounds emit a pale violet color, which may be masked by the strong yellow emission of sodium if it is also present. Cobalt glass can be used to filter out the yellow sodium color.[7] Potassium concentration in solution is commonly determined by flame photometry, atomic absorption spectrophotometry, inductively coupled plasma, or ion selective electrodes.

Figure. 2. The flame-test color for potassium

6.2 Chemical

Potassium must be protected from air for storage to prevent disintegration of the metal from oxide and hydroxide corrosion. Often samples are maintained under a hydrocarbon medium which does not react with alkali metals, such as mineral oil or kerosene.

Like the other alkali metals, potassium reacts violently with water, producing hydrogen. The reaction is notably more violent than that of lithium or sodium with water, and is sufficiently exothermic that the evolved hydrogen gas ignites.

2 K(s) + 2 H2O(l) > H2(g) + 2 KOH(aq)

Because potassium reacts quickly with even traces of water, and its reaction products are nonvolatile, it is sometimes used alone, or as NaK (an alloy with sodium which is liquid at room temperature) to dry solvents prior to distillation. In this role, it serves as a potent desiccant.

Potassium hydroxide reacts strongly with carbon dioxide to produce potassium carbonate, and is used to remove traces of CO2 from air. Potassium compounds generally have excellent water solubility, due to the high hydration energy of the K+ ion. The potassium ion is colorless in water.

Methods of separating potassium by precipitation, sometimes used for gravimetric analysis, include the use of sodium tetraphenylborate, hexachloroplatinic acid, and sodium cobaltinitrite

7. Potassium cations in the body

7.1 Biochemical function

Potassium cations are important in neuron (brain and nerve) function, and in influencing osmotic balance between cells and the interstitial fluid, with their distribution mediated in all animals (but not in all plants) by the so-called Na+/K+-ATPase pump.[8]

Potassium may be detected by taste because it triggers three of the five types of taste sensations, according to concentration. Dilute solutions of potassium ion taste sweet (allowing moderate concentrations in milk and juices), while higher concentrations become increasingly bitter/alkaline, and finally also salty to the taste. The combined bitterness and saltiness of high potassium content solutions makes high-dose potassium supplementation by liquid drinks a palatability challenge.[9]

7.2 Membrane polarization

Potassium is also important in preventing muscle contraction and the sending of all nerve impulses in animals through action potentials. By nature of their electrostatic and chemical properties, K+ ions are larger than Na+ ions, and ion channels and pumps in cell membranes can distinguish between the two types of ions, actively pumping or passively allowing one of the two ions to pass, while blocking the other.[10]

A shortage of potassium in body fluids may cause a potentially fatal condition known as hypokalemia, typically resulting from diarrhea, increased diuresis and vomiting. Deficiency symptoms include muscle weakness, paralytic ileus, ECG abnormalities, decreased reflex response and in severe cases respiratory paralysis, alkalosis and cardiac arrhythmia.

7.3 Filtration and excretion

Potassium is an essential mineral micronutrient in human nutrition; it is the major cation (positive ion) inside animal cells, and it is thus important in maintaining fluid and electrolyte balance in the body. Sodium makes up most of the cations of blood plasma at a reference range of about 145 milliequivalents per liter (3.345 grams) and potassium makes up most of the cell fluid cations at about 150 milliequivalents per liter (4.8 grams). Plasma is filtered through the glomerulus of the kidneys in enormous amounts, about 180 liters per day.[11]

Thus 602 grams of sodium and 33 grams of potassium are filtered each day. All but the 1-10 grams of sodium and the 1-4 grams of potassium likely to be in the diet must be reabsorbed. Sodium must be reabsorbed in such a way as to keep the blood volume exactly right and the osmotic pressure correct; potassium must be reabsorbed in such a way as to keep serum concentration as close as possible to 4.8 milliequivalents (about 0.190 grams) per liter.[12]

Sodium pumps in the kidneys must always operate to conserve sodium. Potassium must sometimes be conserved also, but as the amount of potassium in the blood plasma is very small and the pool of potassium in the cells is about thirty times as large, the situation is not so critical for potassium. Since potassium is moved passively[13][14] in counter flow to sodium in response to an apparent (but not actual) Donnan equilibrium,[15] the urine can never sink below the concentration of potassium in serum except sometimes by actively excreting water at the end of the processing. Potassium is secreted twice and reabsorbed three times before the urine reaches the collecting tubules.[16]

At that point, it usually has about the same potassium concentration as plasma. If potassium were removed from the diet, there would remain a minimum obligatory kidney excretion of about 200 mg per day when the serum declines to 3.0-3.5 milliequivalents per liter in about one week,[17] and can never be cut off completely. Because it cannot be cut off completely, death will result when the whole body potassium declines to the vicinity of one-half full capacity. At the end of the processing, potassium is secreted one more time if the serum levels are too high.

The potassium moves passively through pores in the cell wall. When ions move through pumps there is a gate in the pumps on either side of the cell wall and only one gate can be open at once. As a result, 100 ions are forced through per second. Pores have only one gate, and there only one kind of ion can stream through, at 10 million to 100 million ions per second.[18] The pores require calcium in order to open[19] although it is thought that the calcium works in reverse by blocking at least one of the pores.[20] Carbonyl groups inside the pore on the amino acids mimics the water hydration that takes place in water solution[21] by the nature of the electrostatic charges on four carbonyl groups inside the pore.[22]

7.4 Potassium in the diet and by supplement

7.4.1 Adequate intake

A potassium intake sufficient to support life can generally be guaranteed by eating a variety of foods, especially plant foods. Clear cases of potassium deficiency (as defined by symptoms, signs and a below-normal blood level of the element) are rare in healthy individuals eating a balanced diet. Foods with high sources of potassium include orange juice, potatoes, bananas, avocados, tomatoes, broccoli, soybeans, brown rice, garlic and apricots, although it is also common in most fruits, vegetables and meats.[23]

7.4.2 Optimal intake

Epidemiological studies and studies in animals subject to hypertension indicate that diets high in potassium can reduce the risk of hypertension and possibly stroke (by a mechanism independent of blood pressure), and a potassium deficiency combined with an inadequate thiamine intake has produced heart disease in rats.[24] With these findings, the question of what is the intake of potassium consistent with optimal health, is debated. For example, the 2004 guidelines of the Institute of Medicine specify a DRI of 4,000 mg of potassium (100 mEq), though most Americans consume only half that amount per day, which would make them formally deficient as regards this particular recommendation.[25] Similarly, in the European Union, particularly in Germany and Italy, insufficient potassium intake is somewhat common.[26]

7.4.3 Medical supplementation and disease

Supplements of potassium in medicine are most widely used in conjunction with loop diuretics and thiazides, classes of diuretics which rid the body of sodium and water, but have the side effect of also causing potassium loss in urine. A variety of medical and non-medical supplements are available. Potassium salts such as potassium chloride may be dissolved in water, but the salty/bitter taste of high concentrations of potassium ion make palatable high concentration liquid supplements difficult to formulate.[9] Typical medical supplemental doses range from 10 milliequivalents (400 mg, about equal to a cup of milk or 6 oz. of orange juice) to 20 milliequivalents (800 mg) per dose. Potassium salts are also available in tablets or capsules, which for therapeutic purposes are formulated to allow potassium to leach slowly out of a matrix, as very high concentrations of potassium ion (which might occur next to a solid tablet of potassium chloride) can kill tissue, and cause injury to the gastric or intestinal mucosa. For this reason, non-prescription supplement potassium pills are limited by law in the U.S. to only 99 mg of potassium.

Individuals suffering from kidney diseases may suffer adverse health effects from consuming large quantities of dietary potassium. End stage renal failure patients undergoing therapy by renal dialysis must observe strict dietary limits on potassium intake, as the kidneys control potassium excretion, and buildup of blood concentrations of potassium (hyperkalemia) may trigger fatal cardiac arrhythmia.

8. Applications

About 93% of the world potassium production was consumed by the fertilizer industry.[5]

8.1 Biological applications

Potassium ions are an essential component of plant nutrition and are found in most soil types. Its primary use in agriculture, horticulture and hydroponic culture is as a fertilizer as the chloride (KCl), sulfate (K2SO4) or nitrate (KNO3).

In animal cells, potassium ions are vital to keeping cells alive (see Na-K pump).

In the form of potassium chloride, it is used to stop the heart, e.g. in cardiac surgery and in a solution used in executions by lethal injection.

Figure. 3. Potassium and magnesium sulfate fertilizer

8.2 Food applications

Potassium ion is a nutrient necessary for human life and health. Potassium chloride is used as a substitute for table salt by those seeking to reduce sodium intake so as to control hypertension. The USDA lists tomato paste, orange juice, beet greens, white beans, potatoes, bananas and many other good dietary sources of potassium, ranked according to potassium content per measure shown.[27]

Potassium sodium tartrate, or Rochelle salt (KNaC4H4O6) is the main constituent of baking powder. Potassium bromate (KBrO3) is a strong oxidiser, used as a flour improver (E924) to improve dough strength and rise height.

The sulfite compound, potassium bisulfite (KHSO3) is used as a food preservative, for example in wine and beer-making (but not in meats). It is also used to bleach textiles and straw, and in the tanning of leathers.

8.3 Industrial applications

Potassium vapor is used in several types of magnetometers. An alloy of sodium and potassium, NaK (usually pronounced "nack"), that is liquid at room temperature, is used as a heat-transfer medium. It can also be used as a desiccant for producing dry and air-free solvents.

Potassium metal reacts vigorously with all of the halogens to form the corresponding potassium halides, which are white, water-soluble salts with cubic crystal morphology. Potassium bromide (KBr), potassium iodide (KI) and potassium chloride (KCl) are used in photographic emulsion to make the corresponding photosensitive silver halides.

Potassium hydroxide KOH is a strong base, used in industry to neutralize strong and weak acids and thereby finding uses in pH control and in the manufacture of potassium salts. Potassium hydroxide is also used to saponify fats and oils and in hydrolysis reactions, for example of esters and in industrial cleaners.

Potassium nitrate KNO3 or saltpeter is obtained from natural sources such as guano and evaporites or manufactured by the Haber process and is the oxidant in gunpowder (black powder) and an important agricultural fertilizer. Potassium cyanide KCN is used industrially to dissolve copper and precious metals particularly silver and gold by forming complexes; applications include gold mining, electroplating and electroforming of these metals. It is also used in organic synthesis to make nitriles. Potassium carbonate K2CO3, also known as potash, is used in the manufacture of glass and soap and as a mild desiccant.

Potassium chromate (K2CrO4) is used in inks, dyes, and stains (bright yellowish-red colour), in explosives and fireworks, in safety matches, in the tanning of leather and in fly paper. Potassium fluorosilicate (K2SiF6) is used in specialized glasses, ceramics, and enamels. Potassium sodium tartrate, or Rochelle salt (KNaC4H4O6) is used in the silvering of mirrors.

The superoxide KO2 is an orange-colored solid used as a portable source of oxygen and as a carbon dioxide absorber. It is useful in portable respiration systems. It is widely used in submarines and spacecraft as it takes less volume than O2 (g).

4 KO2 + 2 CO2 > 2 K2CO3 + 3 O2

Potassium chlorate KClO3 is a strong oxidant, used in percussion caps and safety matches and in agriculture as a weedkiller. Glass may be treated with molten potassium nitrate KNO3 to make toughened glass, which is much stronger than regular glass.

9. Precautions

Potassium reacts very violently with water producing hydrogen gas which then usually catches fire. Potassium is usually kept under a hydrocarbon oil such as mineral oil or kerosene to stop the metal from reacting with water vapour present in the air. Unlike lithium and sodium, however, potassium should not be stored under oil indefinitely. If stored longer than 6 months to a year, dangerous shock-sensitive peroxides can form on the metal and under the lid of the container, which can detonate upon opening. It is recommended that potassium, rubidium or caesium not be stored for longer than three months unless stored in an inert (oxygen free) atmosphere, or under vacuum.[28]

As potassium reacts with water to produce highly flammable hydrogen gas, a potassium fire is only exacerbated by the addition of water, and only a few dry chemicals are effective for putting out such a fire (see the precaution section in sodium).

Potassium also produces potassium hydroxide (KOH) in the reaction with water. Potassium hydroxide is a strong alkali and so is a caustic hazard, causing burns.

Due to the highly reactive nature of potassium metal, it must be handled with great care, with full skin and eye protection being used and preferably an explosive resistant barrier between the user and the potassium.

10. Conclusion

As an individual representative of the periodic table of chemical elements Dmitry Ivanovich Mendeleyev, the element has unique chemical and physical properties

Element is of great economic importance and plays a major role in world culture

11. References

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a b "Potassium Without the Taste". http://www.foodnavigator.com/Science-Nutrition/Potassium-without-the-taste. Retrieved Feb 14, 2009.

a b c Mark Winter. "Potassium: Key Information". Webelements. http://www.webelements.com/webelements/elements/text/K/key.html.

a b Ober, Joyce A.. "Mineral Yearbook 2006:Potash". United States Geological Survey. http://minerals.usgs.gov/minerals/pubs/commodity/potash/myb1-2006-potas.pdf. Retrieved 2008-11-20.

Anne Marie Helmenstine. "Qualitative Analysis - Flame Tests". About.com. http://chemistry.about.com/library/weekly/aa110401a.htm.

Bennett CM, Brenner BM, Berliner RW (1968). "Micropuncture study of nephron function in the rhesus monkey". J. Clin. Invest. 47 (1): 203-216. doi:10.1172/JCI105710. PMID 16695942.

Campbell, Neil (1987). Biology. Menlo Park, Calif.: Benjamin/Cummings Pub. Co.. p. 795. ISBN 0-8053-1840-2.

Davy, Humphry (1808). "On some new Phenomena of Chemical Changes produced by Electricity, particularly the Decomposition of the fixed Alkalies, and the Exhibition of the new Substances, which constitute their Bases". Philosophical Transactions of the Royal Society of London 98: 1-45. doi:10.1098/rstl.1808.0001. http://books.google.com/books?id=Kg9GAAAAMAAJ.

Enghag, P. (2004). "11. Sodium and Potassium". Encyclopedia of the elements. Wiley-VCH Weinheim. ISBN 3527306668.

Kernan, Roderick P. (1980). Cell potassium (Transport in the life sciences). New York: Wiley. pp. 40, 48. ISBN 0471048062.

Lans HS, Stein IF, Meyer KA (1952). "The relation of serum potassium to erythrocyte potassium in normal subjects and patients with potassium deficiency". Am. J. Med. Sci. 223 (1): 65-74. doi:10.1097/00000441-195201000-00011. PMID 14902792.

Lockless SW, Zhou M, MacKinnon R.. "Structural and thermodynamic properties of selective ion binding in a K+ channel". Laboratory of Molecular Neurobiology and Biophysics, Rockefeller University. http://www.ncbi.nlm.nih.gov/pubmed/17472437. Retrieved 2008-03-08.

Ober, Joyce A.. "Mineral Commodity Summaries 2008:Potash". United States Geological Survey. http://minerals.usgs.gov/minerals/pubs/commodity/potash/mcs-2008-potas.pdf. Retrieved 2008-11-20.

Potts, W.T.W.; Parry, G. (1964). Osmotic and ionic regulation in animals. Pergamon Press.

Solomon AK (1962). "Pumps in the living cell". Scientific American 207: 100-8. PMID 13914986.

Wright FS (1977). "Sites and mechanisms of potassium transport along the renal tubule". Kidney Int. 11 (6): 415-32. doi:10.1038/ki.1977.60. PMID 875263.