Potassium

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19 argonpotassiumcalcium
Na

K

Rb
General
Name, Symbol, Number potassium, K, 19
Chemical series alkali metals
Group, Period, Block 14, s
Appearance silvery white
Standard atomic weight 39.0983(1) g·mol−1
Electron configuration [Ar] 4s1
Electrons per shell 2, 8, 8, 1
Physical properties
Phase solid
Density (near r.t.) 0.89 g·cm−3
Liquid density at m.p. 0.828 g·cm−3
Melting point 336.53 K
(63.38 °C, 146.08 °F)
Boiling point 1032 K
(759 °C, 1398 °F)
Atomic properties
Crystal structure cubic body centered
Oxidation states 1
(strongly basic oxide)
Electronegativity 0.82 (Pauling scale)
Ionization energies
(more)
1st: 418.8 kJ·mol−1
2nd: 3052 kJ·mol−1
3rd: 4420 kJ·mol−1
Atomic radius 220 pm
Atomic radius (calc.) 243 pm
Covalent radius 196 pm
Van der Waals radius 275 pm
Miscellaneous
Magnetic ordering paramagnetic
Thermal conductivity (300 K) 102.5 W·m−1·K−1
Thermal expansion (25 °C) 83.3 µm·m−1·K−1
Speed of sound (thin rod) (20 °C) 2000 m/s
Young's modulus 3.53 GPa
Shear modulus 1.3 GPa
Bulk modulus 3.1 GPa
Mohs hardness 0.4
Brinell hardness 0.363 MPa
CAS registry number 7440-09-7
Selected isotopes
Main article: Isotopes of potassium
iso NA half-life DM DE (MeV) DP
39K 93.26% K is stable with 20 neutrons
40K 0.012% 1.277×109 y β- 1.311 40Ca
ε 1.505 40Ar
β+ 1.505 40Ar
41K 6.73% K is stable with 22 neutrons
References
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Potassium (IPA: /pə(ʊ)ˈtasiəm/) is a chemical element. It has the symbol K (Arabic: al qaljaLatin: kalium) and atomic number 19. The name "potassium" comes from the word "potash", as potassium was first isolated from potash. Potassium is a soft silvery-white metallic alkali metal that occurs naturally bound to other elements in seawater and many minerals. It oxidizes rapidly in air and is very reactive, especially towards water. In many respects, potassium and sodium are chemically similar, although organisms in general, and animal cells in particular, treat them very differently.

Contents

[edit] Notable characteristics

The flame test
The flame test

Potassium is the second least dense metal after lithium. 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 gray visibly and immediately. 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 reducing medium such as 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.

2K(s) + 2H2O(l) → H2(g) + 2KOH(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 is important in nerve function and in influencing osmotic balance between cells and the interstitiual fluid.[1]

Potassium and its compounds emit a violet color in a flame. This fact is the basis of the flame test for the presence of potassium in a sample.

Potassium compounds generally have excellent water solubility, due to the high hydration energy of the K+ ion. The potassium ion is colorless in water.

Potassium may be detected by taste because it triggers all the types of tastebuds, 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.

Potassium concentration in solution is commonly determined by flame photometry, atomic absorption spectrophotometry, inductively coupled plasma, or ion selective electrodes. Methods of separating potassium by precipitation, sometimes used for gravimetric analysis, include the use of sodium tetraphenyl boron, dihydrogen hexachloroplatinate (IV) hexahydrate, and sodium cobaltinitrite.

[edit] Applications

Many potassium salts are very important, and include: potassium bromide, potassium carbonate, potassium chlorate, potassium chloride, potassium chromate, potassium cyanide, potassium dichromate, potassium iodide, potassium nitrate, potassium sulfate.

[edit] History

Potassium was discovered in 1807 by Sir Humphry Davy, who derived it from caustic potash (KOH). Potassium was the first metal that was isolated by electrolysis.

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 qalīy = "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.

[edit] Occurrence

Potassium in feldspar
Potassium in feldspar

Potassium makes up about 1.5% of the weight of the Earth's crust and is the seventh most abundant element in it. As it is very electropositive, potassium metal is difficult to obtain from its minerals. It is never found free in nature.[citation needed] Potassium salts such as carnallite, langbeinite, polyhalite, and sylvite are found in ancient lake and sea beds. These minerals form extensive deposits in these environments, making extracting potassium and its salts more economical. The principal source of potassium, potash, is mined in Saskatchewan, California, Germany, New Mexico, Utah, and in other places around the world. 3000 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. See 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.

Potassium can be isolated through electrolysis of its hydroxide in a process that has changed little since Davy. Thermal methods also are employed in potassium production, using potassium chloride.

See also potassium minerals.

[edit] Isotopes

Main article: isotopes of potassium

There are 24 known isotopes of potassium. Three isotopes occur naturally: 39K (93.3%), 40K (0.012%) and 41K (6.7%). Naturally occurring 40K decays to stable 40Ar (11.2%) by electron capture and by positron emission, and decays to stable 40Ca (88.8%) 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.[2]

[edit] Precautions

Peroxides (Yellow) and Ozonides (Red) on surface of potassium metal.
Peroxides (Yellow) and Ozonides (Red) on surface of potassium metal.

Solid potassium reacts violently with water, and should therefore be kept under a mineral oil such as kerosene and handled with care. Unlike lithium and sodium however, potassium cannot 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. [3]

The extremely alkaline potassium hydroxide (KOH) residue on the surface of potassium which has been exposed to moisture, is a caustic hazard. As with sodium metal, the "soapy" feel of potassium metal on skin is due to caustic breakdown of the fats in skin into crude soft potassium soap, and represents the beginning of an alkali burn. Potassium should obviously be handled with care, with full skin and eye protection.

Potassium fires are exacerbated by water, and only a few dry chemicals are effective for them. For a fire discussion which applies to alkali metals in general.

[edit] Potassium in nutrition and medicine

Potassium is an essential mineral macronutrient 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.

Potassium is also important in allowing muscle contraction and the sending of all nerve impulses in animals. See action potential for an explanation of the interplay of sodium and potassium in all excitable animal cells. Because of the interaction of the charge on a potassium ion and its surrounding water molecules, K+ ions are actually a little larger than Na+ ions, and ion channels and pumps in cell membranes can easily distinguish between the two types of ions, actively pumping or passively allowing one of the two ions to pass, while blocking the other.

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

Eating a variety of foods that contain potassium is the best way to get an adequate amount. Healthy individuals who eat a balanced diet rarely need supplements. Foods with high sources of potassium include broccoli, orange juice, potatoes, bananas, soybeans, avocados, apricots, pomegranates, parsnips and turnips, although many other fruits, vegetables, and meats contain potassium. Research has indicated that diets high in potassium can reduce the risk of hypertension.

The 2004 guidelines of the Institute of Medicine specify an RDA of 4,000mg of potassium. However, it is thought that most Americans consume only half that amount per day . Similarly, in the European Union, particularly in Germany and Italy, insufficient potassium intake is widespread ([1]).

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 supplements are available.

Some people with kidney disease are advised to avoid large quantities of dietary potassium. End stage renal failure patients undergoing therapy by renal dialysis must observe strict dietary limits on potassium intake, since the kidneys control potassium excretion, and buildup of blood concentrations of potassium may trigger fatal cardiac arrhythmia.

[edit] See also

 

[edit] References

  1. ^ Campbell, Neil (1987). Biology, 795. ISBN 0-8053-1840-2. 
  2. ^ background radiation - potassium-40 - γ radiation.
  3. ^ Thomas K. Wray. DANGER: PEROXIDIZABLE CHEMICALS. Environmental Health & Public Safety (North Carolina State University).

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