potassium – calcium Na K Rb       Full table
General
Name, Symbol, Number	potassium, K, 19
Series	alkali metals
Group, Period, Block	1(IA), 4 , s
Density, Hardness	856 kg/m3, 0.4
Appearance	silvery white
Atomic properties
Atomic weight	39.0983 amu
Atomic radius (calc.)	220 (243) pm
Covalent radius	196 pm
van der Waals radius	275 pm
Electron configuration	[Ar]4s1
e- 's per energy level	2, 8, 8, 1
Oxidation states (Oxide)	1 (strong base)
Crystal structure	cubic body centered
Physical properties
State of matter	solid
Melting point	336.53 K (146.08 °F)
Boiling point	1032 K (1398 °F)
Molar volume	45.94 ×10-6 m3/mol
Heat of vaporization	79.87 kJ/mol
Heat of fusion	2.334 kJ/mol
Vapor pressure	1.06×10-4Pa at __ K
Speed of sound	2000 m/s at 293.15 K
Miscellaneous
Electronegativity	0.82 (Pauling scale)
Specific heat capacity	757 J/(kg*K)
Electrical conductivity	13.9 106/(m·ohm)
Thermal conductivity	102.4 W/(m*K)
1st ionization potential	418.8 kJ/mol
2nd ionization potential	3052 kJ/mol
3rd ionization potential	4420 kJ/mol
4th ionization potential	5877 kJ/mol
5th ionization potential	7975 kJ/mol
6th ionization potential	9590 kJ/mol
7th ionization potential	11343 kJ/mol
8th ionization potential	14944 kJ/mol
9th ionization potential	16963.7 kJ/mol
10th ionization potential	48610 kJ/mol
Most stable isotopes
iso NA half-life DM DE MeV DP 39K 93.26% K is stable with 20 neutrons 40K 0.012% 1.277 E9 y β- ε 1.311 1.505 40Ca 40Ar 41K 6.73% K is stable with 22 neutrons
SI units & STP are used except where noted.

Potassium is a chemical element in the periodic table that has the symbol K (L. kalium) and atomic number 19. This 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, is very reactive, especially in water, and resembles sodium chemically.

Contents

1 Notable characteristics 2 Applications 3 History 4 Occurrence 5 Isotopes 6 Biological Role 7 Precautions 8 References 9 External links

Notable characteristics

With a density less than that of water, potassium is the second lightest metal after lithium. It is a soft solid that easily is cut with a knife and is silvery in color on fresh surfaces. It oxidizes in air rapidly and must be stored in mineral oil for preservation.

Similar to other alkali metals, potassium reacts violently with water producing hydrogen. When in water, it may catch fire spontaneously. Its salts emit a violet color when exposed to a flame.

Applications

• Potassium oxide, best known as potash, is primarily used in fertilizer.
• Potassium nitrate is used in gunpowder.
• Potassium carbonate is used in glass manufacture.
• NaK an alloy of sodium and potassium is used as a heat-transfer medium.
• This element is an essential component needed in plant growth and is found in most soil types.
• In animal cells potassium ions are vital to keeping cells alive (see Na-K pump)
• Potassium chloride is used as a substitute for table salt and is also used to stop the heart, e.g. in cardiac surgery and in executions by lethal injection.

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

History

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

Occurrence

This element makes up about 2.4% of the weight of the Earth's crust and is the seventh most abundant element in it. Due to its insolubility , it is very difficult to obtain potassium from its minerals.

However, other minerals, 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 principle source of potassium, potash is mined in California, Germany, New Mexico, Utah, and in other places around the world. At 3000 ft below the surface of Saskatchewan are large deposits of potash which may become important sources of this element and its salts in the future.

The oceans are another source of potassium but the quantity present in a given volume of seawater is relatively low compared to 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. Potassium is almost never found unbound in nature. However, in living organisms K⁺ ions are important in the physiology of excitable cells.

Isotopes

There are seventeen isotopes of potassium known to exist. The non-synthetic form of potassium are composed of three isotopes: K-39 (93.3%), K-40 (0.01%) and K-41 (6.7%). Naturally occurring K-40 decays to stable Ar-40 (11.2%) by electron capture and by positron emission, and decays to stable Ca-40 (88.8%) by negatron emission; K-40 has a half-life of 1.250 × 10⁹ years.

The decay of K-40 to Ar-40 is commonly used as a 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., Ar-40) was quantitatively retained, i.e., closed system. Minerals are dated by measurement of the concentration of potassium, and the amount of radiogenic Ar-40 that has accumulated. The minerals that are best suited for dating include biotite, muscovite, and 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 be used for nutrient cycling studies because potassium is a macro-nutrient required for life.

K-40 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.

Biological Role

Potassium in the body exists as a monovalent positive ion (a cation), K+, concentrated by energy-requiring mechanisms primarily within cells (intracellular), where it comprises the cell's most abundant monovalent inorganic cation. The body regulates the K+ concentration in blood moderately closely, as substantial fluctuations can affect action potentials, causing heart and nervous problems. Many antibiotics, such as the one produced by the bacterium Bacillus brevis, affect cells by setting up positive ion gates, where the K+ and Na+ ions are permitted to cross the membrane, thus disrupting the action potential of the cell membrane.

The body maintains potassium ion concentration relatively low in blood plasma (usually 3.5 - 5.0 mmol/L), but much higher inside cells (about 100 mmol/L). Abnormally low blood levels, hypokalemia, and abnormally high levels, hyperkalemia, both can adversely affect the heart.

Precautions

Solid potassium reacts violently with water. This metal should therefore be kept under a mineral oil such as kerosene and handled with care.

References

• Los Alamos National Laboratory – Potassium http://periodic.lanl.gov/elements/19.html

External links

• WebElements.com – Potassium http://www.webelements.com/webelements/elements/text/K/index.html
• EnvironmentalChemistry.com – Potassium http://environmentalchemistry.com/yogi/periodic/K.html