General characteristics of oxygen and the reaction of its combustion

50% of the earth’s crust con­sists of oxy­gen. The el­e­ment is also present in min­er­als in the form of salts and ox­ides. Oxy­gen in bond­ed form makes up around 89% of the mass of wa­ter, and is also present in the cells of all liv­ing or­gan­isms and plants. Oxy­gen is also present in the air in a free state in the form O₂ and its al­lotrop­ic mod­i­fi­ca­tion in the form of ozone O₃, and makes up one fifth of the vol­ume of air.

Phys­i­cal and chem­i­cal prop­er­ties of oxy­gen

Oxy­gen O₂ is a gas with­out col­or, taste or smell. It dis­solves poor­ly in wa­ter, and boils at a tem­per­a­ture of -183 de­grees Cel­sius. Oxy­gen in liq­uid form is light blue in col­or, and in sol­id form the el­e­ment forms dark blue crys­tals. Oxy­gen melts at a tem­per­a­ture of -218.7 de­grees Cel­sius.

On heat­ing, oxy­gen en­ters into a re­ac­tion with var­i­ous sim­ple sub­stances (met­als and non-met­als), form­ing ox­ides as a re­sult of in­ter­ac­tion – com­pounds of el­e­ments with oxy­gen. The in­ter­ac­tion of chem­i­cal el­e­ments with oxy­gen is called the ox­i­da­tion re­ac­tion. Ex­am­ples of equa­tions of re­ac­tions are:

4Na + О₂= 2Na₂O,

S + О₂ = SO₂.

Sev­er­al com­plex sub­stances also en­ter into a re­ac­tion with oxy­gen, form­ing ox­ides – the equa­tion of the re­ac­tion is:

СН₄ + 2О₂= СО₂ + 2Н₂О

2СО + О₂ = 2СО₂.

Oxy­gen as a chem­i­cal el­e­ment is ob­tained at lab­o­ra­to­ries and at in­dus­tri­al plants. Oxy­gen can be ob­tained in the lab­o­ra­to­ry by sev­er­al meth­ods:

• by the re­ac­tion of the break­down of Berthol­let’s salt (potas­si­um chlo­rate);
• in the process of the break­down of hy­dro­gen per­ox­ide, heat­ing the sub­stance in the pres­ence of man­ganese ox­ide as a cat­a­lyst;
• by the break­down of potas­si­um per­man­ganate.

Chem­i­cal re­ac­tion of the com­bus­tion of oxy­gen

Pure oxy­gen has spe­cial prop­er­ties that oxy­gen in the air does not have. Air con­tains five times less oxy­gen than pure oxy­gen in the same vol­ume. In air, oxy­gen is mixed with a large amount of ni­tro­gen – a gas that does not burn it­self and does not sup­port com­bus­tion. For this rea­son, if the oxy­gen in the air around a flame is al­ready ex­pend­ed, the next por­tion of oxy­gen must get through ni­tro­gen and prod­ucts of com­bus­tion. Ac­cord­ing­ly, the more en­er­get­ic com­bus­tion of oxy­gen in the at­mos­phere is ex­plained by the swifter sup­ply of oxy­gen to the place of com­bus­tion. In the course of the re­ac­tion, the process of the com­bi­na­tion of oxy­gen with the burn­ing sub­stance takes place more en­er­get­i­cal­ly, and more heat is re­leased. The more oxy­gen that is sup­plied to the burn­ing sub­stance in a unit of time, the more bright­ly the flame burns, the high­er the tem­per­a­ture is and the more in­tense­ly the com­bus­tion process takes place.

How does the com­bus­tion process of oxy­gen take place? This can be test­ed in an ex­per­i­ment. Take a cylin­der and turn it up­side down. Then put a pipe with hy­dro­gen un­der the cylin­der. The hy­dro­gen, which is lighter than air, will com­plete­ly fill the cylin­der. Ig­nite the hy­dro­gen around the open part of the cylin­der, and in­sert a glass pipe into the cylin­der, through which oxy­gen gas flows. Around the end of the pipe, the flame will blaze up, while a flame will calm­ly burn in­side the cylin­der filled with hy­dro­gen. In the course of the re­ac­tion, it is not the oxy­gen that burns, but the hy­dro­gen in the pres­ence of the small amount of oxy­gen com­ing out of the pipe.

What aris­es as a re­sult of the com­bus­tion of hy­dro­gen, and what ox­ide is formed? Hy­dro­gen ox­i­dizes to wa­ter. On the walls of the cylin­der, drops of con­densed wa­ter va­por grad­u­al­ly set­tle. For the ox­i­da­tion of 2 mol­e­cules of hy­dro­gen, 1 mol­e­cule of oxy­gen is used, and 2 wa­ter mol­e­cules form; the equa­tion of the re­ac­tion is:

2Н₂ + O₂ → 2Н₂O

If the oxy­gen comes out of the pipe slow­ly, it burns ful­ly in the at­mos­phere of hy­dro­gen, and the ex­per­i­ment takes place smooth­ly.

As soon as the sup­ply of oxy­gen in­creas­es so much that it does not man­age to burn com­plete­ly, part of it goes out­side the flame, where mix­tures of hy­dro­gen and oxy­gen form, and small in­di­vid­u­al sparks ap­pear, re­sem­bling ex­plo­sions. A mix­ture of oxy­gen and hy­dro­gen is known as det­o­nat­ing gas.

When det­o­nat­ing gas is ig­nit­ed a large ex­plo­sion takes place: with the com­bi­na­tion of oxy­gen and hy­dro­gen, wa­ter forms, and a high tem­per­a­ture de­vel­ops. Wa­ter va­por ex­pands great­ly with the sur­round­ing gas­es, and pres­sure be­comes high, in which not only a frag­ile cylin­der can ex­plode, but also a more durable ves­sel. For this rea­son, you should be ex­treme­ly care­ful when work­ing with det­o­nat­ing mix­ture.

Ex­pen­di­ture of oxy­gen in the com­bus­tion process.

For this ex­per­i­ment, fill a glass crys­tal­liz­er with a vol­ume of 3 liters two thirds full of wa­ter and add a ta­ble­spoon of sodi­um hy­drox­ide or potas­si­um hy­drox­ide. Col­or the wa­ter with phe­nolph­thalein or an­oth­er suit­able dye. Pour sand into a small flask and ver­ti­cal­ly place a wire in it with cot­ton wool on the end. The flask is placed in the crys­tal­liz­er with wa­ter. The cot­ton wool re­mains 10 cm above the sur­face of the so­lu­tion.

Slight­ly wet the cot­ton wool with al­co­hol, oil, hex­ane or oth­er com­bustible flu­id and light. Care­ful­ly cov­er the burn­ing cot­ton wool with the 3-liter flask and low­er it be­low the sur­face of the al­ka­li so­lu­tion. In the com­bus­tion process, the oxy­gen turns to wa­ter and car­bon diox­ide, and as a re­sult of the re­ac­tion the al­ka­li so­lu­tion in the bot­tle ris­es. The cot­ton wool soon goes out. Care­ful­ly place the bot­tle on the bot­tom of the crys­tal­liz­er. In the­o­ry, the bot­tle should fill by one fifth, as air con­tains 20.9% oxy­gen. In com­bus­tion, the oxy­gen turns to wa­ter and car­bon diox­ide CO₂, which is ab­sorbed by the al­ka­li; the equa­tion of the so­lu­tion is:

2NaOH + CO₂ = Na₂­CO₃ + H₂O

In prac­tice, burn­ing ends be­fore all the oxy­gen is ex­pend­ed, some of the oxy­gen turns to car­bon monox­ide, which is not ab­sorbed by the al­ka­li, and some of the air leaves the bot­tle as a re­sult of ther­mal ex­pan­sion.

Warn­ing! Don’t try to re­peat these ex­per­i­ments with­out a pro­fes­sion­al su­per­vi­sion!

Here you’ll find a safer ex­per­i­ment with burn­ing hy­dro­gen