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Free alkaline & alkali ions in water


gatewood
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6 hours ago, gatewood said:

Would it be possible to have, say, calcium cations in aqueous solution, react with O2 to form calcium oxide, if the concentration of the dissolved gas was high enough?

No. The reason is to do with oxidations states. Ca2+ cations are already "oxidised", in that they have a +2 oxidation state, which is the highest possible for them in normal chemistry. (Ca metal, in an oxidation state of 0, reacts vigorously with oxygen.) 

Another way to think of it is that CaO is comprised of Ca cations with a charge of 2+ and oxide anions with a charge of 2-. In the course of the reaction between a Ca metal atom and and an oxygen molecule, two electrons are transferred from Ca, which binds them weakly, to O which binds them strongly.  This results in a net release of energy, leading to a compound has lower chemical  energy than the starting materials - which is the direction all spontaneous chemical reactions take. Whereas if you try to react Ca2+ cations with oxygen, they have no more electrons to give up, so nothing happens. 

Edited by exchemist
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Yes, you explained it neatly and pretty well, though I'm aware of this, given it's ionic bonding 101.

The core of my question, lies more in the fundamental physics that are going on, during the bonding of 2 ions. The way I think I understand ionic bonds, is akin to how, 2 magnets have high potential energy when they're far apart, which gets turned into kinetic energy when "falling" into each other, then all such energy is released and the system is now at a lower energy state, once the magnets come in contact and are at rest.

So given such model, it is not actually the transferring of electrons, which releases energy in itself, but the actual process of coming in contact. I mean, given how electropositive it is, wouldn't a +2 calcium cation quickly bind with basically anything with an inkling of electronegativity or negative polarity? (which is what I think, makes it so soluble in water, just like how sodium cations and chlorine anions break apart to bond with the negative and positive polarity ends of the H2O molecule, their electric fields have found more positive and negative fields on which to be balanced). The cation remains highly reactive till it finds an electronegative atom/molecule, on which to bond and balance the electric field.

Though, to be fair, the O2 molecule is probably a bad candidate, since it is non-polar (it would require some energy to break the covalent bond).

Edited by gatewood
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5 hours ago, gatewood said:

...which is what I think, makes it so soluble in water,...

The White Cliffs of Dover indicate that the solubility isn't just down to the Ca++ ion.

The Ca++ ions in water are surrounded by a bunch of water molecules which are more or less firmly attached.

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9 hours ago, gatewood said:

Yes, you explained it neatly and pretty well, though I'm aware of this, given it's ionic bonding 101.

The core of my question, lies more in the fundamental physics that are going on, during the bonding of 2 ions. The way I think I understand ionic bonds, is akin to how, 2 magnets have high potential energy when they're far apart, which gets turned into kinetic energy when "falling" into each other, then all such energy is released and the system is now at a lower energy state, once the magnets come in contact and are at rest.

So given such model, it is not actually the transferring of electrons, which releases energy in itself, but the actual process of coming in contact. I mean, given how electropositive it is, wouldn't a +2 calcium cation quickly bind with basically anything with an inkling of electronegativity or negative polarity? (which is what I think, makes it so soluble in water, just like how sodium cations and chlorine anions break apart to bond with the negative and positive polarity ends of the H2O molecule, their electric fields have found more positive and negative fields on which to be balanced). The cation remains highly reactive till it finds an electronegative atom/molecule, on which to bond and balance the electric field.

Though, to be fair, the O2 molecule is probably a bad candidate, since it is non-polar (it would require some energy to break the covalent bond).

Yes you are right, the answer I gave before was a simplistic one, for ionic compounds generally. The case of these oxides is a bit more involved, since in fact the second electron affinity of oxygen is +ve, endothermic.

The stability of compounds like CaO relies on a high lattice energy, i.e. the reduction in electrostatic potential that comes from close approach of a large number of oppositely charged ions, in a crystal lattice. In the case of CaO, the O2- anion and the Ca2+ cation are of very similar size, allowing a very efficient packing arrangement which minimises inter-ionic distances, releasing more energy, which compensates for the energy required to get a second extra electron onto the O atom.

So my first answer was a bit misleading. Sorry about that.     

Edited by exchemist
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14 hours ago, John Cuthber said:

The White Cliffs of Dover indicate that the solubility isn't just down to the Ca++ ion.

The Ca++ ions in water are surrounded by a bunch of water molecules which are more or less firmly attached.

Well, that's precisely the point isn't it? A hydration sphere surrounding a cation (firmly attached due to the high electropositivity of the cation).

11 hours ago, exchemist said:

Yes you are right, the answer I gave before was a simplistic one, for ionic compounds generally. The case of these oxides is a bit more involved, since in fact the second electron affinity of oxygen is +ve, endothermic.

The stability of compounds like CaO relies on a high lattice energy, i.e. the reduction in electrostatic potential that comes from close approach of a large number of oppositely charged ions, in a crystal lattice. In the case of CaO, the O2- anion and the Ca2+ cation are of very similar size, allowing a very efficient packing arrangement which minimises inter-ionic distances, releasing more energy, which compensates for the energy required to get a second extra electron onto the O atom.

So my first answer was a bit misleading. Sorry about that.     

Its ok, I wasn't all too clear myself either.

Hmmm... interesting. Forgive my ignorance, but you gave me a lot of questions buddy:

1. Could you elaborate a bit more on your first paragraph, if you would? (that last part)

2. Why does the close proximity of the ions produce a reduced electrostatic potential?

3. Why exactly does the second electron (to complete the valence shell) requires an extra kick to orbit oxygen?

4. Aren't the alkaline and alkali oxides, actually rather reactive? If exposed to the atmosphere, they will form metal carbonates... if I understand it correctly.

Edited by gatewood
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On 6/13/2021 at 6:03 AM, gatewood said:

Yes, you explained it neatly and pretty well, though I'm aware of this, given it's ionic bonding 101.

Look around you.

Water, alcohol, oil, resins, methane, nitrogen, oxygen, carbondioxide, benzene; These are common liquids or gases.
And they are all covalently bonded, albeit some have straightforward covalency some have dative covalency (ie are polar).

Compare this with sodium chloride, iron oxide, copper sulphate; These are all common solids.
And they are all ionically bonded.

This situation represents the very large majority of cases.
This is no accident, there are good reasons for this.

That is to observe that ionic compounds tend to from solids whilst covalent compounds often appear as fluids.

Yes one of these reasons is molecular weight.
But there are plenty of examples of ionic solids with a lower molecular weight than covalent liquids, eg sodium chloride is 58 whilst benzene is 78, both from my list.

So it is instructive to consider what is different.

The difference is that in a fluid the molecules have a degree of autonomy not present in a solid.
They can move about as a molecule.
And most important you can for instance identify one particular carbon atom with two particular oxygen atoms forming the 'molecule'.
However you cannot identify a particular sodium atom (ion) with one particular chlorine atom (ion) in the solid.
In fact electric forces link one  (each) sodium+ to 6 chloride-
The coordination number is said to be 6.
So the intensity of the charge difference is distributed that way.

https://courses.lumenlearning.com/cheminter/chapter/ionic-crystal-structures/

This is the key difference of importance to your question.
The chemical implications of these can be very complex indeed as both John Cuthber and exchemist are trying to tell you.

 

 

 

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10 hours ago, gatewood said:

Well, that's precisely the point isn't it? A hydration sphere surrounding a cation (firmly attached due to the high electropositivity of the cation).

Its ok, I wasn't all too clear myself either.

Hmmm... interesting. Forgive my ignorance, but you gave me a lot of questions buddy:

1. Could you elaborate a bit more on your first paragraph, if you would? (that last part)

2. Why does the close proximity of the ions produce a reduced electrostatic potential?

3. Why exactly does the second electron (to complete the valence shell) requires an extra kick to orbit oxygen?

4. Aren't the alkaline and alkali oxides, actually rather reactive? If exposed to the atmosphere, they will form metal carbonates... if I understand it correctly.

Like all these things it gets more complex when you delve into it. Electron affinity is the energy released by an electronegative atom when it gains an extra electron and becomes an anion. For example, all the halogens release energy on gaining an electron, meaning the anion has lower energy than the neutral atom. The same is true for oxygen when it gains one electron.  However when it gains a second, that is energetically unfavourable, due to the repulsion from the net -ve charge of the anion towards a second electron. (I had forgotten this, and only remembered after looking it up.)

However electron affinity is only a measure of the energy change when a free atom or ion gains an extra electron. In the case of metal oxides, the oxygen atom is not free. It is sitting in a crystal lattice, in the present case (CaO) surrounded by 6 nearest neighbour Ca2+ ions. That makes its environment much more energetically attractive for oxygen to pick up a second electron and form O2-. Hence it is common to find metal oxides with O2- anions even though, if the oxygen atom were free, you would have to "force" it to accept a second electron. (It's significant that if you put these oxides in contact with water you never get hydrated O2- ions. When they come out of the crystal lattice they pinch an H+ ion from water to make OH- (hydroxide) - plus another OH- from what is left of the water molecule:

O2- +H2O -> 2OH- .

As for the question about close proximity lowering potential, that's just replaying what you said, in effect,  about magnets. You bring opposite magnetic poles together, or opposite electric changes together, and you lower the magnetic or electrostatic potential energy. That is reflected in the fact that you have to do work to pull them apart again.  And, as they come together, magnets can gain kinetic energy, just as you said, at the expense of the magnetic potential energy. Similarly, ions with opposite charges approaching one another gain kinetic energy at the expense of electrostatic potential energy - which, in the context of molecular scale processes, means the heat energy given off in an exothermic reaction.

The alkaline earth (Group II) metal oxides react with CO2 in the air, yes. I'm less sure about the alkali (Group I) metals. The carbonate anion has a charge of 2- and this means it needs 2 M+ atoms to go with it, so I'm not sure how the kinetics and thermodynamics of that work out. 

    

Edited by exchemist
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16 hours ago, exchemist said:

Like all these things it gets more complex when you delve into it. Electron affinity is the energy released by an electronegative atom when it gains an extra electron and becomes an anion. For example, all the halogens release energy on gaining an electron, meaning the anion has lower energy than the neutral atom. The same is true for oxygen when it gains one electron.  However when it gains a second, that is energetically unfavourable, due to the repulsion from the net -ve charge of the anion towards a second electron. (I had forgotten this, and only remembered after looking it up.)

However electron affinity is only a measure of the energy change when a free atom or ion gains an extra electron. In the case of metal oxides, the oxygen atom is not free. It is sitting in a crystal lattice, in the present case (CaO) surrounded by 6 nearest neighbour Ca2+ ions. That makes its environment much more energetically attractive for oxygen to pick up a second electron and form O2-. Hence it is common to find metal oxides with O2- anions even though, if the oxygen atom were free, you would have to "force" it to accept a second electron. (It's significant that if you put these oxides in contact with water you never get hydrated O2- ions. When they come out of the crystal lattice they pinch an H+ ion from water to make OH- (hydroxide) - plus another OH- from what is left of the water molecule:

O2- +H2O -> 2OH- .

As for the question about close proximity lowering potential, that's just replaying what you said, in effect,  about magnets. You bring opposite magnetic poles together, or opposite electric changes together, and you lower the magnetic or electrostatic potential energy. That is reflected in the fact that you have to do work to pull them apart again.  And, as they come together, magnets can gain kinetic energy, just as you said, at the expense of the magnetic potential energy. Similarly, ions with opposite charges approaching one another gain kinetic energy at the expense of electrostatic potential energy - which, in the context of molecular scale processes, means the heat energy given off in an exothermic reaction.

The alkaline earth (Group II) metal oxides react with CO2 in the air, yes. I'm less sure about the alkali (Group I) metals. The carbonate anion has a charge of 2- and this means it needs 2 M+ atoms to go with it, so I'm not sure how the kinetics and thermodynamics of that work out. 

    

- First paragraph

Ok, then its just like stacking a bunch of N magnets onto a single S magnet, the atom loses the electron affinity, simply because its electric field becomes more balanced with the new electron. E.g. the Ca++ has higher energy than a Ca+ because its electric field is even more unbalanced, and an O- has more trouble accepting an electron than a neutral O, because... well, it is even more electronegative. Simple.

- Second paragraph (first half)

But wouldn't the Ca++ actually compete with the O-- for the electrons to reduce themselves?

- Second paragraph (second half)

You taught me something rather interesting, thank you very much :)

- Third paragraph

Sure, I realized too late, how easy that question was 🤪 . Thanks for taking the time to clarify anyhow :)

- Fourth paragraph

Yes, it does pick up 2 alkaline metals. E.g. 2Na+:

https://www.google.com/search?q=sodium+carbonate+molecule&tbm=isch&ved=2ahUKEwixjbOsopjxAhUJTqwKHRw7CSIQ2-cCegQIABAA&oq=sodium+carbonate+molecule&gs_lcp=CgNpbWcQAzIECCMQJzICCAAyAggAMgIIADICCAAyAggAMgIIADIGCAAQCBAeMgYIABAIEB4yBggAEAgQHjoECAAQQzoGCAAQBRAeUJtiWLFsYM1taABwAHgAgAGRBIgB3RKSAQswLjEuNS4wLjEuMZgBAKABAaoBC2d3cy13aXotaW1nwAEB&sclient=img&ei=-OTHYLGxDomcsQWc9qSQAg&bih=798&biw=1600&client=-b-d

17 hours ago, studiot said:

Look around you.

Water, alcohol, oil, resins, methane, nitrogen, oxygen, carbondioxide, benzene; These are common liquids or gases.
And they are all covalently bonded, albeit some have straightforward covalency some have dative covalency (ie are polar).

Compare this with sodium chloride, iron oxide, copper sulphate; These are all common solids.
And they are all ionically bonded.

This situation represents the very large majority of cases.
This is no accident, there are good reasons for this.

That is to observe that ionic compounds tend to from solids whilst covalent compounds often appear as fluids.

Yes one of these reasons is molecular weight.
But there are plenty of examples of ionic solids with a lower molecular weight than covalent liquids, eg sodium chloride is 58 whilst benzene is 78, both from my list.

So it is instructive to consider what is different.

The difference is that in a fluid the molecules have a degree of autonomy not present in a solid.
They can move about as a molecule.
And most important you can for instance identify one particular carbon atom with two particular oxygen atoms forming the 'molecule'.
However you cannot identify a particular sodium atom (ion) with one particular chlorine atom (ion) in the solid.
In fact electric forces link one  (each) sodium+ to 6 chloride-
The coordination number is said to be 6.
So the intensity of the charge difference is distributed that way.

https://courses.lumenlearning.com/cheminter/chapter/ionic-crystal-structures/

This is the key difference of importance to your question.
The chemical implications of these can be very complex indeed as both John Cuthber and exchemist are trying to tell you.

 

 

 

Well, I'm aware of the melting and boiling point difference between covalent bonds, which have low melting and boiling points, because the energy that binds the molecules is low, and ionic ones, which bind their molecules far more strongly, that's why, huge temperatures are needed to even begin to have them behave as fluids.

And about the molecular weight, I was kinda aware of that, but it'll be a whole different topic to talk about more complex molecules such as aromatic compounds (I mean, cellulose is a huge polymer and it'll decompose way before even melting).

Finally, I'm definitely gonna study that last part, thanks for sharing it :)

18 hours ago, John Cuthber said:

So, your point was that the cliffs are mainly  made from a highly soluble thing that doesn't dissolve.

The dover white cliffs are an entirely different creature, they're made mostly of chalk (calcite minerals), product of ancient coccolithophores and other microorganisms, their shells are made of calcium carbonate (most sea shells are made of it), a metal carbonate that is NOT soluble (the very part of hardwood ashes that won't dissolve with water).

Edited by gatewood
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10 hours ago, gatewood said:

 

- Second paragraph (first half)

But wouldn't the Ca++ actually compete with the O-- for the electrons to reduce themselves?

 

That strikes me as a rather penetrating question. +1. 

The answer, I think, must be that for Ca++ to pinch an electron from O-- would involve it getting a lot bigger, because the electron would have to go into the next shell (4s), which is at a greater distance from the nucleus than the 3s and 3p subshells, which are already full in Ca++. As I mentioned in an earlier post, Ca++ and O-- are of similar size and can pack efficiently. A larger Ca+ ion would pack less efficiently. The larger size would push the neighbouring ions apart, reducing the strength of the ionic bonding and leading to a higher energy state overall for the crystal. In other words, it would reduce the so-called "lattice energy". 

 

 

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16 hours ago, gatewood said:

The dover white cliffs are an entirely different creature, they're made mostly of chalk (calcite minerals), product of ancient coccolithophores and other microorganisms, their shells are made of calcium carbonate (most sea shells are made of it), a metal carbonate that is NOT soluble (the very part of hardwood ashes that won't dissolve with water).

Was there any part of that which you thought I didn't already know?

 

16 hours ago, gatewood said:

But wouldn't the Ca++ actually compete with the O-- for the electrons to reduce themselves?

Yes and no.

While we usually consider CaO to be "ionic", if you look at the electron distribution (by Xray diffraction) you see that, on average, not all of the 2 electrons are on the oxygen.

If you choose an easier ion to oxidise- for example iodide and a metal that's harder to oxidise (or easier to reduce) like lead or silver you get an "ionic" solid, but (unlike the iodide, or silver, ions) it's yellow.

That's because the energy from a blue photon is high enough to kick an electron from the anion back onto the cation.

https://en.wikipedia.org/wiki/Charge-transfer_band

The TiO2 which is used as a pigment in most white paint is similar- the electron can be moved from the oxide ion back to the titanium ion, but only by light in the extreme blue end of the spectrum, and that's not normally noticeable.
 

 

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19 hours ago, gatewood said:

Well, I'm aware of the melting and boiling point difference between covalent bonds, which have low melting and boiling points, because the energy that binds the molecules is low, and ionic ones, which bind their molecules far more strongly, that's why, huge temperatures are needed to even begin to have them behave as fluids.

 

I'm gald you found something useful in my offering.

However you should avoid this myth that ionic bonds are 'stronger than covalent ones'.

Some are , some are not.

The hardest natural substance known is a covalent crystal and is used to make special scientific equipment for containing very very high pressures (greater than at the base of the solid part of the Earth) called the diamond anvil.

As John and exchemist are noting there are also situations of intermediacy where a bond is mostly ionic (covalent) but a little bit covalent (ionic).

 

As regards your original question,

Perhaps you should study the reaction of quicklime with water to produce slaked lime.

Although bonding other aspects of physchem are very modern and greeat fun I still urge you to get a good grasp of 'the stuff of matter'.
Chemistry is very much about the stuff of matter and how it behaves and the theory follows knowledge of this, rather than the other way round.

Yes I know there are advanced chemists who will tell you about nucleophilic reaction sites and so on , but I guarantee they have a good knowledge of 'stuff' and its behaviour to work from.

 

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