Why Do Atoms Need to Have Free Electrons to Create Covalent Bonds?

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Atoms bond because compounds are more stable or have less energy than individual atoms. Interatomic potential energy depends with distance of atoms and there is a distance at which potential energy has minimum. This distance is a length of the bond.

When forming covalent bonds, why is it important to have free electrons as minimum of potential energy can be obtained regardless of having free electrons?
 

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  • #2
hutchphd
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What is your definition of "free electrons ?"
 
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  • #3
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What is your definition of "free electrons ?"
Non - paired valence electrons
 
  • #4
jim mcnamara
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Covalent bonds share an electron - between two non-metals. The reaction components of covalent bonds are electrically neutral. For ionic bonds the reaction components are both charged.

Do you see why?

I'm unsure what you do not see, so this was the best I could do for you.
 
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  • #5
TeethWhitener
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When forming covalent bonds, why is it important to have free electrons as minimum of potential energy can be obtained regardless of having free electrons?
Can you give an example of what you mean? As written, this makes no sense.
 
  • #7
mjc123
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Since its compounds all contain other elements, the compounds of all other elements combined exceed those of carbon.
 
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  • #8
Delta2
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Since its compounds all contain other elements, the compounds of all other elements combined exceed those of carbon.
E hehe I guess @sysprog means the compounds of all other elements with elements other than carbon.
 
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I think that it's a non-rigorous statement from a good encyclopedia.
 
  • #10
Astronuc
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I think that it's a non-rigorous statement from a good encyclopedia.
I would agree. I suspect that the statement reflects the vast number of organic compounds. However, one usually finds hydrogen along with carbon. One can find metal hydrides as much as one finds carbides, but I'd have to think about elements that would form hydrides and not carbides. There are metal carbonates, as well has metal hydroxides and oxyhydroxides, and hydrated compounds. On the other hand, when it comes to halides, carbon can form CClF3, CCl2F2, CCl3F, as well as CF4 and CCl4. Hydrogen can only do HCl and HF. So, maybe C can outdo H in terms of unique compounds. Perhaps take any compound with H and replace H with C, one can form many different compounds, but the reverse is not true, because of the difference in number of valence electrons 1 (H) and 4 (C).

My work actually involves carbides and hydrides, as well as many exotic alloys and compounds, and how they behave in different systems.
 
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As @Astronuc's illustrative post pointed out, hydrogen and carbon are often found together -- I think that the term 'hydrocarbon' has applicabilty.
Depending on the electronegativity of the elements, Fluorine is the most reactive element on the periodic table.
Yes, flourine is by far the most reactive element; however, even though it can react with an otherwise 'inert' gas, and is stored in liquid-nitrogen-chilled nickel containers, it does not have anywhere near as many reactions as carbon.
 
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  • #13
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Atoms bond because compounds are more stable or have less energy than individual atoms. Interatomic potential energy depends with distance of atoms and there is a distance at which potential energy has minimum. This distance is a length of the bond.

When forming covalent bonds, why is it important to have free electrons as minimum of potential energy can be obtained regardless of having free electrons?
There can be "a minimum" but the next questions will be depth and directionality of the minimum.
H atom has unpaired electron. Two H atoms have a minimum of potential energy. The depth of that minimum is around 435 kJ/mol per the pair.
H2 molecule does not have unpaired electrons. The paired electrons of H2 molecule are attracted to paired electrons of other H2 molecules, and have a minimum of potential energy. The binding energy of a H2 molecule is around 925 J/mol (the heat of evaporation of solid H2), divided between a number of nearest neighbours.
 
  • #14
TeethWhitener
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Thanks to @snorkack for refocusing the thread.
It’s also worth pointing out that unpaired electrons are not, in fact, required for the formation of a covalent bond. Coordinate covalent (or dative) bonds are an example of this. Specific compounds include ammonia borane (NH3BH3) and many metal coordination compounds (e.g., Ni(CO)4). These compounds feature closed shell molecules with lone pairs (Lewis bases) donating electron density into empty orbitals of other compounds (Lewis acids).

The paired electrons don’t even need to be centered on a single atom. ##\eta^2##-ethylene compounds such as Zeise’s salt and higher hapticity compounds such as ferrocene are examples of bonds with a high covalent character where the paired electrons of the Lewis bases don’t reside primarily on a single atom.
 
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  • #15
TeethWhitener
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The paired electrons of H2 molecule are attracted to paired electrons of other H2 molecules
I’ll also point out that this isn’t strictly accurate. There are a few different interpretations of dispersion forces between molecules such as H2, but direct electron-electron interaction is always repulsive without some intermediating effect (cf. BCS theory).
 

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