Chemical and physical characteristics of calcium, its interaction with water

Why is it stored in a sealed container?

[Deposit Photos]

Cal­ci­um is found in the fourth large pe­ri­od, sec­ond group, main sub-group, with the atom­ic num­ber 20. The atom­ic mass of cal­ci­um, ac­cord­ing to the pe­ri­od­ic ta­ble, is 40.08. The for­mu­la of the high­est ox­ide is CaO. The sym­bol of the el­e­ment is Ca, af­ter the first two let­ters of the word cal­ci­um.

Char­ac­ter­is­tics of ba­sic cal­ci­um

In or­di­nary con­di­tions, cal­ci­um is a met­al with a sil­very-white col­or. With high chem­i­cal ac­tiv­i­ty, the el­e­ment can form many com­pounds that be­long to dif­fer­ent class­es. The el­e­ment has im­por­tance for tech­ni­cal and in­dus­tri­al chem­i­cal syn­the­ses. It is wide­spread in the earth’s crust, with a per­cent­age of around 1.5%. It is an al­ka­line earth el­e­ment, as when it dis­solves in wa­ter it forms al­ka­lis, but in na­ture it is en­coun­tered in the form of nu­mer­ous min­er­als and salts. Sea wa­ter has a high con­cen­tra­tion of cal­ci­um (400 mg/l).

Pure calcium in a protective argon atmosphere [Wikimedia]

The char­ac­ter­is­tics of cal­ci­um de­pend on the ar­range­ment of its crys­tal struc­ture, which ex­ist in two types – cu­bic face-cen­tered and body-cen­tered. The type of bond in the mol­e­cule of cal­ci­um is metal­lic.

Nat­u­ral sources of cal­ci­um:

  • ap­atites;
  • al­abaster;
  • plas­ter;
  • cal­cite;
  • flu­o­rite;
  • dolomite.

Phys­i­cal prop­er­ties of cal­ci­um and meth­ods for ob­tain­ing the met­al

The ag­gre­gate state of cal­ci­um in or­di­nary con­di­tions is sol­id, and the met­al melts at a tem­per­a­ture of 842 de­grees of Cel­sius. The el­e­ment is a good con­duc­tor and heat con­duc­tor, and has a shiny sil­ver-white col­or. When heat­ed, it first moves to a liq­uid, then to a va­porous state, and los­es its metal­lic prop­er­ties. The den­si­ty of the el­e­ment is light – it is a soft met­al that can be cut with a knife. It boils at a tem­per­a­ture of 1484 de­grees Cel­sius.

When pres­sure is ap­plied to cal­ci­um, it be­gins to lose its metal­lic prop­er­ties and ca­pac­i­ty for con­duct­ing elec­tric­i­ty. But when pres­sure is in­creased fur­ther, the metal­lic prop­er­ties are re­stored, and it dis­plays the prop­er­ties of a su­per-con­duc­tor, which ex­ceeds oth­er el­e­ments in these pa­ram­e­ters by sev­er­al times.

For a long time, it was not pos­si­ble to ob­tain the met­al in free form – ow­ing to its high chem­i­cal ac­tiv­i­ty, this el­e­ment is not en­coun­tered in na­ture in pure form. The el­e­ment was not dis­cov­ered un­til the ear­ly 19th cen­tu­ry. Cal­ci­um was first syn­the­sized as a met­al by the British sci­en­tist Humphrey Davy. He was the first to dis­cov­er the na­ture of the in­ter­ac­tion of com­pounds of sol­id met­als and salts with an elec­tric cur­rent. Nowa­days, the elec­trol­y­sis of cal­ci­um salts (mix­tures of cal­ci­um and potas­si­um chlo­rides, mix­tures of cal­ci­um flu­o­ride and chlo­ride) is still the most ef­fec­tive method of ob­tain­ing the met­al. Cal­ci­um is also ex­tract­ed from its ox­ide with the use of alu­minothermy – a method that is wide­spread in met­al­lur­gy.

The chem­i­cal prop­er­ties of cal­ci­um

Cal­ci­um is an ac­tive met­al that en­ters into many in­ter­ac­tions. In nor­mal con­di­tions, it eas­i­ly re­acts with the for­ma­tion of cor­re­spond­ing bi­na­ry com­pounds: with oxy­gen and halo­gens. Click here for learn­ing more about cal­ci­um com­pounds. When heat­ed it re­acts with ni­tro­gen, hy­dro­gen, car­bon, sil­i­con, boron, phos­pho­rous, sul­fur and oth­er sub­stances. In open air, it im­me­di­ate­ly in­ter­acts with oxy­gen and car­bon diox­ide, so it be­comes cov­ered with a grey coat­ing. It re­acts vi­o­lent­ly with acids, some­times burst­ing into flame. Cal­ci­um dis­plays in­ter­est­ing prop­er­ties in the com­po­si­tion of salts. For ex­am­ple, cave sta­lac­tites and sta­lag­mites con­sist of cal­ci­um car­bon­ate which grad­u­al­ly forms from wa­ter, car­bon diox­ide and hy­dro­car­bon­ate un­der the in­flu­ence of pro­cess­es in­side un­der­ground wa­ter.

Ow­ing to its high ac­tiv­i­ty in an or­di­nary state, cal­ci­um is stored in the lab­o­ra­to­ry in a dark glass, with a tight­ly closed lid and un­der a lay­er of paraf­fin or kerosene. The qual­i­ta­tive re­ac­tion to the cal­ci­um ion is that a flame turns a bright, rich brick-red col­or.

Flame test: brick-red color originates from calcium [Wikimedia]

The met­al can be iden­ti­fied in the com­po­si­tion of com­pounds by the undis­solved sed­i­ments of some salts of the el­e­ment (flu­o­ride, cal­ci­um car­bon­ate, sul­fate, sil­i­cate, phos­phate, sul­fite)

The re­ac­tion of wa­ter with cal­ci­um

Cal­ci­um is stored in con­tain­er un­der a lay­er of pro­tec­tive liq­uid. To con­duct an ex­per­i­ment to demon­strate how the re­ac­tion of wa­ter and cal­ci­um takes place, you can­not just take out the met­al and cut off a piece of it. It is eas­i­er to use metal­lic cal­ci­um in the lab­o­ra­to­ry in the form of fil­ings which are pre­pared with a lathe.

If you don’t have any fil­ings, and there are no small pieces of cal­ci­um in the con­tain­er, then you will need pli­ers or a ham­mer. Use the tool to sep­a­rate the re­quired piece of cal­ci­um, and place it in a flask or a glass of wa­ter. A cal­ci­um fil­ing is placed in a dish in a gauze bag.

The cal­ci­um sinks to the bot­tom, and hy­dro­gen is re­leased – first­ly in the place where the met­al was bro­ken off. Grad­u­al­ly, gas is re­leased from the sur­face of the cal­ci­um – the process re­sem­bles vig­or­ous boil­ing. At the same time a sed­i­ment of cal­ci­um hy­drox­ide (slaked lime) is formed.

Slaking Lime [Flickr]

The piece of cal­ci­um ris­es to the sur­face, cov­ered with hy­dro­gen bub­bles.

With­in sev­er­al tenths of a sec­ond the cal­ci­um dis­solves, and from the sus­pen­sion of hy­drox­ide the wa­ter turns cloudy white. If you car­ry out the re­ac­tion in a test tube in­stead of a glass, you can ob­serve the re­lease of heat: the test tube quick­ly be­comes hot. The re­ac­tion of cal­ci­um with wa­ter does not end with an ef­fec­tive ex­plo­sion, but the two sub­stances in­ter­act vig­or­ous­ly, mak­ing for an im­pres­sive sight. The ex­per­i­ment is safe.

If you take the bag with the re­main­ing cal­ci­um out of the wa­ter and keep it in the air for a while, as a re­sult of the con­tin­u­ing re­ac­tion it will heat up in­tense­ly, and the wa­ter re­main­ing in the gauze will boil. If you fil­ter part of the cloudy so­lu­tion through a fun­nel into a glass, then if car­bon diox­ide is passed through the so­lu­tion, a sed­i­ment will form. You do not need car­bon diox­ide gas for this, you can sim­ply blow ex­haled air into the so­lu­tion through a glass straw.