Chemistry: Electron Configurations & Reactivity

[physicurious]

Trying to remember some chemistry here and think I must me remembering something wrong. I'm remembering that the more unpaired electrons in an atom the more reactive it is, or the more bonds it can form. So hydrogen is reactive ie 1 unpaired electron in the 1s orbital. Helium is stable 2 paired electrons in the 1s orbital.

But now beryllium is plenty reactive and very capable of forming chemical bonds. It's electron configuration is 1s2 2s2 ie no unpaired electrons.

I also know carbon can form four bonds but its electron configuration is 1s2 2s2 2p2 ie only two unpaired electrons. I'm thinking about something wrong. Help!

 

[Mayhem]

Carbon obtains its famous tetravalent property from orbital hybridization, where the 2s and 2p orbitals hybridize into sp, sp² and sp³ orbitals. That is the explanation - the physics may differ.

 

[Dr Bob]

Let's take some examples and go from there. IN GENERAL, atoms like to have their outside shells filled. For elements in group 1 (Li, Na, K, etc) the outside shell is the 2s for Li, 3s for Na, 4s for K etc. By losing that one electron the next lower shell is filled. They do that easily, they are very reactive.

As you move DOWN the table from top to bottom in group 1, the reactivity increases BECAUSE that outside electron is further, distance wise, from the nucleus, the outside electron has less pull on it from the nucleus, that one electron is more easily pulled away, which results in increased reactivity. The same logic holds for elements in group 2 (Mg, Ca, Sr, etc) and group 13 (B, Al, Ga, etc).

You will notice that I didn't talk about H in group 1, Be in group 2 or B in group 13. I omitted those on purpose because, in general, the first element in the series often doesn't follow the general rule and/or has additional rules. Having to include those in a general discussion complicates the discussion to it's easier to talk about them separately.

On the other side of the table we have elements like those in group 17 (F, Cl, Br, I etc). Those elements have SEVEN electrons in their outside shell. They need to ADD one more electron to make that shell full so they do that with vigor. As one moves DOWN the table from F, to I, the reactivity DECREASES. For group 16, (O, S etc) and group 15 (N, P,, As etc) the logic is much the same as for the F block.. The reactivity is decreasing as one moves to the left and down the table.

For reactions between group I or 2 elements with group 16 or 17 elements the compounds formed are ionic in nature. The rules are somewhat different where covalent bonds are the order of the day. I hope I have helped. Follow up as needed.

 

[physicurious]

Let's take some examples and go from there. IN GENERAL, atoms like to have their outside shells filled. For elements in group 1 (Li, Na, K, etc) the outside shell is the 2s for Li, 3s for Na, 4s for K etc. By losing that one electron the next lower shell is filled. They do that easily, they are very reactive.

As you move DOWN the table from top to bottom in group 1, the reactivity increases BECAUSE that outside electron is further, distance wise, from the nucleus, the outside electron has less pull on it from the nucleus, that one electron is more easily pulled away, which results in increased reactivity. The same logic holds for elements in group 2 (Mg, Ca, Sr, etc) and group 13 (B, Al, Ga, etc).

You will notice that I didn't talk about H in group 1, Be in group 2 or B in group 13. I omitted those on purpose because, in general, the first element in the series often doesn't follow the general rule and/or has additional rules. Having to include those in a general discussion complicates the discussion to it's easier to talk about them separately.

On the other side of the table we have elements like those in group 17 (F, Cl, Br, I etc). Those elements have SEVEN electrons in their outside shell. They need to ADD one more electron to make that shell full so they do that with vigor. As one moves DOWN the table from F, to I, the reactivity DECREASES. For group 16, (O, S etc) and group 15 (N, P,, As etc) the logic is much the same as for the F block.. The reactivity is decreasing as one moves to the left and down the table.

For reactions between group I or 2 elements with group 16 or 17 elements the compounds formed are ionic in nature. The rules are somewhat different where covalent bonds are the order of the day. I hope I have helped. Follow up as needed.

Thanks for this reply. When I was thinking of this at first maybe I was thinking about 'reactivity' being synonymous with how many bonds and element can form. But now that I think about it another way of thinking about 'reactivity' could be how easily an element forms a bond which I think, correct me if I'm wrong, sounds more like what you're describing.

What you wrote especially makes sense when I look at an electronegativity table. Those group 1 elements you were talking about have lower electronegativities the more you move down the chart. This makes sense in my head now because I can see that lone electron in that last shell will be not only progressively farther and farther away from the positively charged nucleus AND there will be some shielding from the positive charge by all the other electrons occupying the lower energy levels.

So then I can see how those group 1 elements would easily lose an electron to elements having higher electronegativities. Especially with elements like Nitrogen, Oxygen, Fluorine, and chlorine. Well at least I'm guessing a bit there just by looking at their electronegativities.

It makes similar sense why the electronegativities of group 13 elements tend to go down to as you go down the table because of a combination of distance from the protons and shielding of positive charge from the electrons at lower energies. And also why going down the table would mean less reactivity - tending to have lesser electronegativites they would be less likely to gain an electron.

Although now that I'm thinking in terms of electronegativities, distance of electrons from proton charges, and electron shielding I'm realizing some of my reasoning is starting to break down here and doesn't really explain why group 13 elements tend to have such high electronegativities. By using the same logic I did before I would expect the elements to also have decreasing electronegativities moving from left to right on the table. But of course they do not. Oh well trying to understand that bit is a thought for another day.

Thanks again for the reply and giving me much to think about. Cheers!

 

[Borek]

Thanks for this reply. When I was thinking of this at first maybe I was thinking about 'reactivity' being synonymous with how many bonds and element can form. But now that I think about it another way of thinking about 'reactivity' could be how easily an element forms a bond which I think, correct me if I'm wrong, sounds more like what you're describing.

Yes, that's more like it.

Think about having a piece of elemental silicon and a piece of elemental potassium. Si will produce 4 bonds, K will produce one, so by your first approach Si is more reactive. But drop them both into water - Si won't change (at least not till you get seriously bored), K will react vigorously and produce an explosion in seconds.

So then I can see how those group 1 elements would easily lose an electron to elements having higher electronegativities. Especially with elements like Nitrogen, Oxygen, Fluorine, and chlorine. Well at least I'm guessing a bit there just by looking at their electronegativities.

This is a bit more complicated - elemental N (in the form of N₂) is highly inert and making it react even with alkali metals is not trivial.