Potassium


Potassium is a chemical element with symbol K (from Neo-Latin kalium) and atomic number 19. It was first isolated from potash, the ashes of plants, from which its name
derives. In the periodic table, potassium is one of the alkali metals. All of the alkali metals have a single valence electron in the outer electron shell, which is easily removed to create an ion with a positive charge – a cation, which combines with anions to form salts. Potassium in nature occurs only in ionic salts. Elemental potassium is a soft silvery-white alkali metal that oxidizes rapidly in air and reacts vigorously with water, generating sufficient heat to ignite hydrogen emitted in the reaction, and burning with a lilac-colored flame. It is found dissolved in sea water (which is 0.04% potassium by weight[5][6]), and is part of many minerals.

Potassium is chemically very similar to sodium, the previous element in group 1 of the periodic table. They have a similar first ionization energy, which allows for each atom to give up its sole outer electron. That they are different elements that combine with the same anions to make similar salts was suspected in 1702,[7] and was proven in 1807 using electrolysis. Naturally occurring potassium is composed of three isotopes, of which 40
K is radioactive. Traces of 40
K are found in all potassium, and it is the most common radioisotope in the human body.

Potassium ions are vital for the functioning of all living cells. The transfer of potassium ions through nerve cell membranes is necessary for normal nerve transmission; potassium deficiency and excess can each result in numerous signs and symptoms, including an abnormal heart rhythm and various electrocardiographic (ECG) abnormalities. Fresh fruits and vegetables are good dietary sources of potassium. The body responds to the influx of dietary potassium, which raises serum potassium levels, with a shift of potassium from outside to inside cells and an increase in potassium excretion by the kidneys.

Most industrial applications of potassium exploit the high solubility in water of potassium compounds, such as potassium soaps. Heavy crop production rapidly depletes the soil of potassium, and this can be remedied with agricultural fertilizers containing potassium, accounting for 95% of global potassium chemical production.

Physical Properties


Potassium is the second least dense metal after lithium. It is a soft solid with a low melting point, and can be easily cut with a knife. Freshly cut potassium is silvery in appearance, but it begins to tarnish toward gray immediately on exposure to air.

 In a flame test, potassium and its compounds emit a lilac color with a peak emission wavelength of 766.5 nanometers.


Chemical  Properties




Neutral potassium atoms have 19 electrons, one more than the extremely stable configuration of the noble gas argon.

 Because of this and its low first ionization energy of 418.8 kJ/mol, the potassium atom is much more likely to lose the last electron and acquire a positive charge than to gain one and acquire a negative charge (though negatively charged alkalide K− ions are not impossible).

This process requires so little energy that potassium is readily oxidized by atmospheric oxygen. In contrast, the second ionization energy is very high (3052 kJ/mol), because removal of two electrons breaks the stable noble gas electronic configuration (the configuration of the inert argon).

Potassium therefore does not form compounds with the oxidation state of +2 or higher.

Potassium is an extremely active metal that reacts violently with oxygen in water and air. 


With oxygen it forms potassium peroxide, and with water potassium forms potassium hydroxide. 

The reaction of potassium with water is dangerous because of its violent exothermic character and the production of hydrogen gas.

Hydrogen reacts again with atmospheric oxygen, producing water, which reacts with the remaining potassium. This reaction requires only traces of water; because of this, potassium and the liquid sodium-potassium (NaK) alloy are potent desiccants that can be used to dry solvents prior to distillation.

Because of the sensitivity of potassium to water and air, reactions with other elements are possible only in an inert atmosphere such as argon gas using air-free techniques. Potassium does not react with most hydrocarbons such as mineral oil or kerosene.

 It readily dissolves in liquid ammonia, up to 480 g per 1000 g of ammonia at 0 °C. Depending on the concentration, the ammonia solutions are blue to yellow, and their electrical conductivity is similar to that of liquid metals. 

In a pure solution, potassium slowly reacts with ammonia to form KNH
2, but this reaction is accelerated by minute amounts of transition metal salts.

Because it can reduce the salts to the metal, potassium is often used as the reductant in the preparation of finely divided metals from their salts by the Rieke method.
For example, the preparation of magnesium by this method employs potassium as the reductant:

MgCl2 + 2K → Mg + 2KCl

Isotopes






There are 24 known isotopes of potassium, three of which occur naturally: 39K (93.3%), 40K (0.0117%), and 41K (6.7%). Naturally occurring 40K has a half-life of 1.250×109 years. It decays to stable 40 Ar by electron capture or positron emission (11.2%) or to stable 40 Ca by beta decay (88.8%).

The decay of 40K to 40Ar is the basis of a common 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 (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 best suited for dating include biotite, muscovite, metamorphic hornblende, and volcanic feldspar; whole rock samples from volcanic flows and shallow instrusives can also be dated if they are unaltered.
Apart from dating, potassium isotopes have been used as tracers in studies of weathering and 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. 40K is the radioisotope with the largest abundance in the body. In healthy animals and people, 40
K represents the largest source of radioactivity, greater even than 14
C. In a human body of 70 kg mass, about 4,400 nuclei of 40
K decay per second.The activity of natural potassium is 31 Bq/g.

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