Potassium: Properties, Reactivity, and Diverse Applications

Potassium, together with its close sibling, sodium, is among the most abundant elements found in the Earth’s crust, in seawater and in biological systems (Brown et al. 256).  However, potassium is very rarely found free or in pure form in nature.  This is due to its intense reactivity, which results to its being always found as an ion in nature, either in solid minerals or in aqueous solution (Silberberg 978).  The major ore of potassium is sylvite (chemical name potassium chloride), which is mined from bedded salt deposits

The pure form is obtained from the ore by electrolytic reduction of molten potassium chloride by liquid sodium (eq.

1):

KCl _(l)   +   Na _(l)                                           NaCl _(l)    +   K_(g)              (eq 1)

This reduction is done at 850°C, well above the boiling point of potassium, which is 757°C.  As the potassium gas is removed, the system shifts to produce more potassium.  The gas is then condensed and purified by fractional distillation (Silberberg 979).  It is then cast into sticks, which are of course, in solid form (Holtzclaw and Robinson 364).

Get quality help now

Bella Hamilton

Verified writer

Proficient in: Chemistry

5 (234)

“ Very organized ,I enjoyed and Loved every bit of our professional interaction ”

+84 relevant experts are online

Hire writer

            Potassium, along with sodium and lithium belong to the Group 1A elements or the Alkali Metals.  Like all elements in this group, potassium is a soft metallic solid with silvery metallic luster and high thermal and electrical conductivity (Brown et al. 257).

            As mentioned earlier, potassium is very reactive, due to the fact that alkali metals have very low ionization values.  This means that the outer s electron can be easily removed, which explains why potassium is normally found as an ion with a +1 charge (K⁺).  One of potassium’s vigorous reactions is with water.  All alkali metals react violently with water.  Such reactions are very exothermic and in many cases produce enough heat to ignite H₂, producing a fire or explosion (Brown et al. 257).

This reaction is almost always explosive with potassium, because heavier members of the group have more violent results.   Potassium is also easily oxidized.  That is why, in the laboratory, potassium (in ore or in sticks) must be protected from contact with air to prevent their instantaneous oxidation and also to minimize the possibility of fires (Silberberg 980).

            Potassium’s uses are quite limited because it does not have any major uses for which sodium cannot be substituted.  The major use of potassium at present is in an alloy with sodium to be used as a heat exchanger in chemical and nuclear reactors.  When potassium is allowed to have direct contact with oxygen (O₂), it forms a superoxide, KO₂ (eq 2):

            K _(s)   +   O_{2 (g)}                              KO_{2 (s)}                                         (eq 2)

This material is utilized as an emergency source of O₂ in breathing masks for miners, divers, submarine crews, or firefighters (Silberberg 980).  Furthermore, fertilizers usually contain large amounts of potassium, as this is needed to meet the needs of growing plants (Brown et al. 256).

         Although potassium in its pure form may have limited uses, it does form compounds with other elements to produce substances which are widely used in industries.  Potassium hydroxide (KOH) for example, is very alkaline and is considered to be a strong base. Because of this characteristic, it is used as saponifying agent for liquid soaps.  Furthermore, an aqueous solution of KOH functions as an electrolyte, which finds extensive uses in alkaline batteries (“Potassium Hydroxide”).

Thus, potassium hydroxide is a primary contributor in powering flashlights and other battery-operated equipment.  Potassium hydroxide is also the main ingredient in the production of potassium carbonate (K₂CO₃).  Potassium carbonate, likewise, has a variety of uses.  Among these is the production of specialty glasses, including television tubes (ERCO Worldwide).  Furthermore, fire extinguishing powders, dyes and pigments, dehydrating agents, boiler compounds, and titanium enamels are produced using potassium carbonate.

  Potassium phosphates are another important potassium compounds produced from potassium hydroxide.  These compounds have excellent solubility and find their uses in industrial detergents and cleaners, as well as in water treatment chemicals (ERCO Worldwide).  Antifreezes are also manufactured with the use of potassium phosphates.

    Potassium salts also have wide range of uses in several sectors of the community.  Potassium nitrate (KNO₃) is an important component of agricultural fertilizer as well as the oxidizing component of gunpowder.  The major sources of potassium nitrate are animal manure and the drainings of decomposing organic materials.

Potassium bromide (KBr), on the other hand, is used as an anti-convulsive and anti-seizure medication, usually administered only to dogs (Dodds and Carson).  Potassium iodide (KI), potassium chloride (KCl) and potassium bromide, compounds formed by potassium with halides, are used in photography and radiation treatments.  Potassium chromate (K₂CrO₄), with its bright yellow-red color is used in dyeing and staining, as well as in explosives and fireworks.

Works Cited

Brown, Theodore, LeMay, Eugene Jr., Bursten, Bruce and Burdge, Julia. Chemistry the Central Science. New Jersey: Pearson Education Inc., 2006.

Dodds, Jean and Carson, Joanne. “Potassium Bromide (KBr)” 7 Feb 2008.

Holtzclaw, Henry and Robinson, William. College Chemistry with Qualitative Analysis. Canada: D.C. Heath and Company, 1988.

Silberberg, Martin. Chemistry: The Molecular Nature of Matter and Change, USA: McGraw-Hill Companies, Inc., 2006.