alevels2k17
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it is mentioned in my book that the stability of an ionic compound increases as the charge on the ions increases......... how are group 1 metals more stable than group 2 then ( since their carbonates dont decompose due to high stability) i dont get it isnt there a contradiction here
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samb1234
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(Original post by NoorL)
it is mentioned in my book that the stability of an ionic compound increases as the charge on the ions increases......... how are group 1 metals more stable than group 2 then ( since their carbonates dont decompose due to high stability) i dont get it isnt there a contradiction here
Group 1 ions have a +1 charge so polarise the co3 2- ion less
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alevels2k17
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(Original post by samb1234)
Group 1 ions have a +1 charge so polarise the co3 2- ion less
but the book says stability of ionic compounds increases as the charge on the ions increases ughhh im so confused.. isnt what u just said the opposite
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RMNDK
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(Original post by NoorL)
it is mentioned in my book that the stability of an ionic compound increases as the charge on the ions increases......... how are group 1 metals more stable than group 2 then ( since their carbonates dont decompose due to high stability) i dont get it isnt there a contradiction here
Stability gets a bit fiddly when you start considering polarisation as well.

The reason why Group 1 metal carbonates have a higher temperature for thermal decomposition is, as sam1234 pointed out, they have a lower charge.

The carbonate ion, CO32- has three oxygen atoms; two of them have the negative charge. In this case, it's best to think of all three oxygen atoms being negative and the cloud of negative charge being spread out over the whole of these three oxygen atoms.
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This is called resonance. In actual fact, if you were to freeze time, you'd find that at any time, only two oxygen atoms posses the negative charge, but because the oxygen atoms keep passing around this negative charge so quickly, we just think of the whole molecule being negative.

http://www.chemguide.co.uk/inorganic...carbonate3.gif
If you get a Mg2+ ion, it's going to strongly polarise the carbonate ion and strongly attract the negative charge at one oxygen atom. This is going to leave the carbon and the other two oxygen atoms less negative and more ready to leave as CO2

But if you get a Na+ ion, it's not going to polarise the carbonate ion as much. This means the carbon atom and two oxygen atoms are still quite negative, and it's not going to be released as CO2 so easily. Thus, in order to liberate it, we need to add a higher temperature. That makes Na2CO3 more stable.

Normally, between two ions there is an electrostatic force of attraction that binds them together. And you're right, if you increase the charge of the ions, the electrostatic force will increase which increases strength of bonding, which makes the ionic lattice much more stable. That's ionic bonding.

But there is a point where the electrostatic force of attraction becomes so great that it actually starts bringing back the electrons on the negative atom. It distorts the electron cloud of the anion and forms, what seems to be, a covalent bond.

What we actually say is that the bond shows covalent character. (you might come across this term if you're doing A2, or AS even, if not, don't use it)

So as you can see there is a point in which the charge of the ion actually decreases the stability of a compound (there's a lot of factors in play which is why it muddles people up).
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Here's another example if you're interested.

The melting point of MgO is much higher than Na2O. That's expected because Mg2+ ions have a higher charge, so the charge density is much greater than that of Na+ thus the ionic bonding is stronger.

But the melting point of Al2O3 is lower than that of MgO. You would think that Al3+ ions would form even stronger ionic bonding.

In fact, because the Al3+ has such a huge charge density, it actually polarises the oxide ion by distorting the electron cloud. The electrons it just gave away, it now attracts them again. This forms a "somewhat" covalent bond. Because there is less ionic bonding, the melting point is lower.
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alevels2k17
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(Original post by RMNDK)
Stability gets a bit fiddly when you start considering polarisation as well.

The reason why Group 1 metal carbonates have a higher temperature for thermal decomposition is, as sam1234 pointed out, they have a lower charge.

The carbonate ion, CO32- has three oxygen atoms; two of them have the negative charge. In this case, it's best to think of all three oxygen atoms being negative and the cloud of negative charge being spread out over the whole of these three oxygen atoms.
Spoiler:
Show
This is called resonance. In actual fact, if you were to freeze time, you'd find that at any time, only two oxygen atoms posses the negative charge, but because the oxygen atoms keep passing around this negative charge so quickly, we just think of the whole molecule being negative.

http://www.chemguide.co.uk/inorganic...carbonate3.gif
If you get a Mg2+ ion, it's going to strongly polarise the carbonate ion and strongly attract the negative charge at one oxygen atom. This is going to leave the carbon and the other two oxygen atoms less negative and more ready to leave as CO2

But if you get a Na+ ion, it's not going to polarise the carbonate ion as much. This means the carbon atom and two oxygen atoms are still quite negative, and it's not going to be released as CO2 so easily. Thus, in order to liberate it, we need to add a higher temperature. That makes Na2CO3 more stable.

Normally, between two ions there is an electrostatic force of attraction that binds them together. And you're right, if you increase the charge of the ions, the electrostatic force will increase which increases strength of bonding, which makes the ionic lattice much more stable. That's ionic bonding.

But there is a point where the electrostatic force of attraction becomes so great that it actually starts bringing back the electrons on the negative atom. It distorts the electron cloud of the anion and forms, what seems to be, a covalent bond.

What we actually say is that the bond shows covalent character. (you might come across this term if you're doing A2, or AS even, if not, don't use it)

So as you can see there is a point in which the charge of the ion actually decreases the stability of a compound (there's a lot of factors in play which is why it muddles people up).
Spoiler:
Show
Here's another example if you're interested.

The melting point of MgO is much higher than Na2O. That's expected because Mg2+ ions have a higher charge, so the charge density is much greater than that of Na+ thus the ionic bonding is stronger.

But the melting point of Al2O3 is lower than that of MgO. You would think that Al3+ ions would form even stronger ionic bonding.

In fact, because the Al3+ has such a huge charge density, it actually polarises the oxide ion by distorting the electron cloud. The electrons it just gave away, it now attracts them again. This forms a "somewhat" covalent bond. Because there is less ionic bonding, the melting point is lower.
oh okay so ur saying that the covalent character that resulted from high polarisation in mg+2 reduced the stability of the compound and that is why na+ appeared more stable (bcz the polarisation was low in it na+???
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RMNDK
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(Original post by NoorL)
oh okay so ur saying that the covalent character that resulted from high polarisation in mg+2 reduced the stability of the compound and that is why na+ appeared more stable (bcz the polarisation was low in it na+???
Precisely.
The fact that the electrons are more shared between the ions in Magnesium Carbonate means it will more readily decompose. Sodium will not polarise the carbonate ion as much.
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alevels2k17
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(Original post by RMNDK)
Precisely.
The fact that the electrons are more shared between the ions in Magnesium Carbonate means it will more readily decompose. Sodium will not polarise the carbonate ion as much.
alright.. thanks a lot for clarifying it xx
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alevels2k17
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(Original post by RMNDK)
Precisely.
The fact that the electrons are more shared between the ions in Magnesium Carbonate means it will more readily decompose. Sodium will not polarise the carbonate ion as much.
one last question... why are group 1 metals able to decompose no3^-1 and not carbonates... carbonates have a charge of -2 so arent they easier to decompose??
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RMNDK
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(Original post by NoorL)
one last question... why are group 1 metals able to decompose no3^-1 and not carbonates... carbonates have a charge of -2 so arent they easier to decompose??
Does your book/specification say that Group 1 carbonates do not decompose period? They do decompose, just at higher temperatures.

But I do accept that nitrates are more easier to decompose than carbonates.

Let's compare NaNO3 and Na2CO3

NaNO3 decomposes to NaNO2 and O2
Spoiler:
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Ignore lithium, it decomposes to lithium oxide, nitrogen dioxide and oxygen because it's polarising enough.
This image might help:
https://puffthemutantdragon.files.wo...rate.png?w=584

When Na+ approaches the NO3- it's going to attract the electron pair on the oxygen atom. This will cause the other oxygen atom to be removed as a free radical. This happens twice, so the two free radicals combine to form oxygen, i.e. 2NaNO3 ---> 2NaNO2 + O2

When Na+ approaches the CO32-, it's going attract the oxygen atom and form the oxide. This will liberate CO2, i.e. Na2CO3 ---> Na2O + CO2

Although the CO32- has a higher negative charge, you're having two Na+ ions polarising it, instead of one like in sodium nitrate. So the two factors effectively cancel out.

What makes the nitrate more easier to decompose is that you're liberating an oxygen radical. In sodium carbonate, you need to remove a whole CO2 molecule which is considerably more difficult. You need to polarise the oxygen atom so much in order to expel carbon and the other oxygen atoms.

This is why Li+ will form the oxide. It's attractive enough to obtain an oxide ion and expel Nitrogen and the other oxygen atoms as NO2
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It's mechanism is a bit more complicated because it also forms Oxygen, but I'm not in a position to comment on that.
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alevels2k17
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(Original post by RMNDK)
Does your book/specification say that Group 1 carbonates do not decompose period? They do decompose, just at higher temperatures.

But I do accept that nitrates are more easier to decompose than carbonates.

Let's compare NaNO3 and Na2CO3

NaNO3 decomposes to NaNO2 and O2
Spoiler:
Show
Ignore lithium, it decomposes to lithium oxide, nitrogen dioxide and oxygen because it's polarising enough.
This image might help:
https://puffthemutantdragon.files.wo...rate.png?w=584

When Na+ approaches the NO3- it's going to attract the electron pair on the oxygen atom. This will cause the other oxygen atom to be removed as a free radical. This happens twice, so the two free radicals combine to form oxygen, i.e. 2NaNO3 ---> 2NaNO2 + O2

When Na+ approaches the CO32-, it's going attract the oxygen atom and form the oxide. This will liberate CO2, i.e. Na2CO3 ---> Na2O + CO2

Although the CO32- has a higher negative charge, you're having two Na+ ions polarising it, instead of one like in sodium nitrate. So the two factors effectively cancel out.

What makes the nitrate more easier to decompose is that you're liberating an oxygen radical. In sodium carbonate, you need to remove a whole CO2 molecule which is considerably more difficult. You need to polarise the oxygen atom so much in order to expel carbon and the other oxygen atoms.

This is why Li+ will form the oxide. It's attractive enough to obtain an oxide ion and expel Nitrogen and the other oxygen atoms as NO2
Spoiler:
Show
It's mechanism is a bit more complicated because it also forms Oxygen, but I'm not in a position to comment on that.
thank you very much xx
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