I want to know how do metals replace hydrogen in a substance such as water. Isn't water happy with its configuration? It has shared enough electrons so that both hydrogen and oxygen complete their orbitals.

I know that alkali metals are very active and they have very low electronegativity. And elements such as gold have much higher electronegativity while being considered very inactive. Maybe I have a wrong understanding of what electronegativity is but isn't the element with higher electronegativity supposed to be more active and be able to replace those with lower ones?

  • $\begingroup$ What do you mean by Aurum? Shouldn't it be gold? $\endgroup$ Commented May 30, 2017 at 11:19
  • $\begingroup$ Yes that's what i meant. But in my language we refer it to as Aurum instead of gold $\endgroup$
    – Ebrin
    Commented May 30, 2017 at 11:26
  • $\begingroup$ Alright! Nevertheless I'll have to edit your question so that it confirms to IUPAC standards. For your question, do you know what electropositivity is? $\endgroup$ Commented May 30, 2017 at 11:28
  • $\begingroup$ I'm just trying to understand it now, from what i've gathered it's how much element wants to attract electrons and i think it's only relevant in covalent bonds, but i've never got a clear explanation from my teacher. $\endgroup$
    – Ebrin
    Commented May 30, 2017 at 11:32
  • $\begingroup$ 2 Na + 2 H2O -> 2 NaOH + H2 < for instance in this example, Does the sodium want to give away its electrons more than hydrogen does and that causes replacement of hydrogen or are there other reasons? $\endgroup$
    – Ebrin
    Commented May 30, 2017 at 11:39

1 Answer 1


There's electronegativity and then there's electropositivity. Electronegativity is the tendency of a substance to gain electrons and electropositivity on similar grounds is the tendency to donate electrons. What actually is a redox reaction? Exchange of electrons?

A substance with high electronegativity would like to snatch any electron it finds, while a substance with high electropositivity would like to force-fully hand over its electrons, however unwilling the substance on the receiving end may be. Metals are electropositive.

how do metals replace hydrogen in a substance such as water. Isn't water happy with its configuration?

There was water, it was living happily. Now, water had two components. Oxygen and hydrogen. Let's pull up their electronegativities. Oxygen is at $3.44$ and hydrogen is at $2.2$. This means that the bond that oxygen and hydrogen shared in water was a polar bond. Oxygen being powerful (higher electronegativity) kept the bond in it's favor. The electrons of the bond were more towards oxygen. This gave oxygen a partial negative charge ($\delta -$) and hydrogen a partial positive charge ($\delta+$).

In came a metal which was highly electropositive (think alkali metals, as an example you can take up sodium whose electronegativity is at $0.93$). He gave oxygen a tempting offer:

"Oxygen, you seem to be caught up in a tussle, look at hydrogen, he doesn't want a positive charge because he's quite electronegative. You're constantly fighting with him to keep the electrons of the bond pair pulled towards yourself. Why don't you pair up with me? I'm low at electronegativity. This means I'm electropositive. As much as electropositive species like bearing a negative charge, we electropositive species like bearing a positive charge!"

And oxygen was greedy. A greater difference in electronegativities of elements forming the bond meant a bond which would be richer in ionic character, which appealed to oxygen.

  • $\begingroup$ Wow, that's such a good explanation, thank you very much. It's much more clear now, but with that i have few more questions : 1)Do the 'partial charges' only occur during hydrogen bonds? or is it just a bond with smaller charges? 2)My teacher said that metals always have the positive charge when interacting with hydrogen, but what about the ones that have higher electronegativity than hydrogen? for example Gold has 2.4. I now understand from your comment that they woudn't be the best to pair up, but do they not interact at all? and is it possible to somehow make them react? $\endgroup$
    – Ebrin
    Commented May 30, 2017 at 12:06
  • $\begingroup$ @Ebrin The pleasure's all mine. Let me say a bit more about partial charges. Say if you had the same elements forming a bond, i.e. $\ce{A-A}$ , you wouldn't have partial charges at all, when a difference of electronegativities comes into play, there's this thing called bond polarization. It can be due to various other factors, you might want to look up dipoles, induced dipoles etc. It occurs basically in every bond, as long as there's a electronegativity difference between the constituents. $\endgroup$ Commented May 30, 2017 at 12:29
  • $\begingroup$ @Ebrin In a bond, the element which has a higher electronegativity would bear a negative charge, that's true. pubs.rsc.org/en/content/articlelanding/cc/2016/… this article hints at the formation of a $\ce{Au-H}$ bond, win! $\endgroup$ Commented May 30, 2017 at 12:32
  • $\begingroup$ Okay that makes sense, So what about Ionic bonds? Are ionic bonds just the same as covalent bonds, but the difference between electronegativites is much higher? My teacher said the ionic bonds occurs only between metals and nonmetals, she wasn't too clear about why or how it happens. From what I've gathered, there are partial charges between covalent bonds, and actual charges between ionic bonds? Do the covalent bonds give us cations and anions ? $\endgroup$
    – Ebrin
    Commented May 30, 2017 at 12:50
  • $\begingroup$ The thing that confuses me the most, is that in class we would treat Ionic and covalent bonds the same exact way, only difference being that ionic was between metal and nonmetals. And when i was studying on the web, people would say that covalent bonds don't really have a charge and ionics do. So are they pretty much the same? what are the key differences? Sorry for asking so many questions, i just have to understand this once and for all. $\endgroup$
    – Ebrin
    Commented May 30, 2017 at 12:54

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