Facts about calcium, and the reaction of calcium with oxygen

Physical and chemical properties of calcium

[Deposit Photos]

In an­cient times, peo­ple used cal­ci­um com­pounds for build­ing, main­ly cal­ci­um car­bon­ate, which was found in rocks, or lime, which was made from bak­ing it. They also used mar­ble and gyp­sum. Right up un­til the end of the 18th cen­tu­ry, sci­en­tists mis­tak­en­ly be­lieved that lime, which is cal­ci­um ox­ide, was an el­e­ment, un­til An­toine Lavoisi­er pro­posed his the­o­ries about the sub­stance.

Limestone quarry [Flickr]

In the ear­ly 19th cen­tu­ry, the Eng­lish sci­en­tist Humphrey Davy dis­cov­ered cal­ci­um in pure form, us­ing elec­trol­y­sis. He ob­tained a cal­ci­um amal­gam from slaked lime and mer­cury ox­ide. By re­mov­ing the mer­cury, he got metal­lic cal­ci­um.

Cal­ci­um plays an im­por­tant role in bi­ol­o­gy, and is a wide­spread mi­croele­ment found in small quan­ti­ties in the hu­man body, in an­i­mals and also in plants. A de­fi­cien­cy of this el­e­ment leads to var­i­ous dis­eases.

As the met­al dis­plays high ac­tiv­i­ty, in na­ture cal­ci­um is not en­coun­tered in free form. The ma­jor­i­ty of com­pounds are in rocks, for ex­am­ple sil­i­cates and alu­mi­nosil­i­cates. It is also found in sed­i­ment rocks, for ex­am­ple in lime­stone and chalk, which are made of cal­ci­um car­bon­ate, and also in sea­wa­ter and un­der­ground wa­ter in the earth’s crust.

Ap­pli­ca­tion of cal­ci­um

This el­e­ment is of­ten used in the met­al­lur­gi­cal in­dus­try, where cal­ci­um is used as a re­duc­er to ob­tain sev­er­al met­als, for ex­am­ple stain­less steel.

Oth­er cal­ci­um com­pounds used in in­dus­try:

  • cal­ci­um ox­ide CaO or quick­lime, which is used in build­ing and re­pair works;

  • cal­ci­um hy­dro­sul­fate Ca(HSO₃)₂ con­sists of col­or­less crys­tals, and is a preser­va­tive, which is also used in the pa­per in­dus­try;

  • cal­ci­um sul­fate di­hy­drate or gyp­sum Ca­SO₄·2H₂O is used as a bond­ing ma­te­ri­al in the pa­per and cel­lu­lose in­dus­try, and in medicine for hold­ing frac­tures in place.

Phys­i­cal prop­er­ties of cal­ci­um

  • Ca is a soft met­al, which can eas­i­ly be cut with a knife;

  • Cal­ci­um has a shiny sil­very-white col­or, which grows dull from the for­ma­tion of an ox­ide film when stored in­cor­rect­ly;

  • it has a high melt­ing point – 842 de­grees Cel­sius.

  • it has high elec­tric­i­ty and heat con­duc­tiv­i­ty;

  • it boils at a tem­per­a­ture of 1484 de­grees Cel­sius.

Chem­i­cal prop­er­ties of cal­ci­um

Ca is lo­cat­ed in the sec­ond group of the fourth pe­ri­od in the pe­ri­od­ic ta­ble. Cal­ci­um is an ac­tive al­ka­line earth met­al.

Cal­ci­um should be stored in kerosene, be­cause if the met­al is left in the open air, it swift­ly los­es its metal­lic shine and be­comes dull and grey from the im­pact of wa­ter va­por, oxy­gen and car­bon diox­ide.

Cal­ci­um re­acts vi­o­lent­ly with wa­ter, but with­out ig­ni­tion. The abun­dant re­lease of hy­dro­gen caus­es the piece of cal­ci­um to move around in the wa­ter. Cal­ci­um hy­drox­ide also forms. If phe­nolph­thalein is added to the liq­uid, it turns a bright crim­son col­or, which proves that Ca(OH)₂ is a base.

Ca + 2H₂O → Ca(OH)₂↓ + H₂↑

The re­ac­tion of cal­ci­um with oxy­gen

The re­ac­tion of Ca and O₂ is very in­ter­est­ing, but this ex­per­i­ment must not be con­duct­ed at home, as it is very dan­ger­ous.

Let us ex­am­ine the re­ac­tion of cal­ci­um and oxy­gen, and name­ly the com­bus­tion of this sub­stance in air.

Warn­ing! Don’t try to re­peat this ex­per­i­ment with­out a pro­fes­sion­al su­per­vi­sion! Here you’ll find safe chem­istry ex­per­i­ments to do at home.

As a source of oxy­gen, we will take potas­si­um ni­trate KNO₃. If cal­ci­um is stored in kerosene, be­fore the ex­per­i­ment we must clean it with a spir­it burn­er, hold­ing it above the flame. We then place the cal­ci­um is KNO₃ pow­der, mak­ing sure it is well-cov­ered. Then we place the cal­ci­um with potas­si­um ni­trate in the flame of the burn­er. The potas­si­um ni­trate breaks down into potas­si­um ni­trite and oxy­gen. The oxy­gen re­leased burns the cal­ci­um, and the flame turns red.

KNO₃ → KNO₂ + O₂

2Ca + O₂ → 2CaO

We should note that cal­ci­um only re­acts with cer­tain el­e­ments when heat­ed – they in­clude phos­pho­rus, sul­fur, boron, ni­tro­gen and oth­ers.