why does carbon usually form four single bonds & not two?

why does carbon usually form four single bonds & not two?

Grade:Upto college level

4 Answers

52 Points
12 years ago

carbon +4 is more stable than +2 [inert pair effect].

hence it usually form foir single bonds.

Aman Bansal
592 Points
12 years ago

Dear Bipasha,

Since the valency of carbon is 4 so it can form 4 bonds only. In that case it is called saturated hydrocarbon. If on the contrary it forms a double or triple bonds then it is called unsaturated hydrocarbon.

Best Of luck

Plz Approve the answer...!!!!

Cracking IIT just got more exciting,It’s not just all about getting assistance from IITians, alongside Target Achievement and Rewards play an important role. ASKIITIANS has it all for you, wherein you get assistance only from IITians for your preparation and winexciting gifts by answering queries in the discussion forums. Reward points 5 + 15 for all those who upload their pic and download the ASKIITIANS Toolbar, just a simple click here to download the toolbar….
So start the brain storming…. become a leader with Elite Expert League ASKIITIANS


Aman Bansal

Askiitian Expert

bhanuveer danduboyina
95 Points
12 years ago

the tetravalency of carbon is 4 and it can bond 4 single bonds. If it forms 2 single bonds, it cannot be stable 

Godfrey Classic Prince
633 Points
12 years ago

Dear Mehruna Bipasha,

Carbon atom is tetravalent in nature i.e it wants four more electrons to reach the stable noble gas(Neon) configuration of 1s2 2s2 2p6

Carbon atom has 6 electrons in it & has an electronic configuration of 1s2 2s2 2p2

It therefore to become stable must get four more electrons to attain the stable Noble Gas electronic configuration!!!

That is the reason why it forms four single bonds & not just two single bonds !!


Hope you understood now..

All the Best & Good Luck !!

Please approve my answer if you liked it by clicking on "Yes" given below..!!!Smile

Think You Can Provide A Better Answer ?


Get your questions answered by the expert for free