Nagging problems


by Bubonic Plague
Tags: nagging
Bubonic Plague
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#1
Sep20-03, 07:52 AM
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After 2 years of chemistry, i have some nagging problems which glare at me everytime.

When atoms combine to form compounds, they combine to form compounds with zero valence right? But in that case why are there compounds such as NO3^-? or SO4^2- or OH^-? If such compounds exist, then why can't there be compounds like such as PS^2-, or BBr^3-, etc with any combinations like these?

And since the bonds are all essentially metallic, covalent and ionic, then ultimately couldn't i have any combination of elements, example KMgFeBOHSe, after all, it would just be a matter of placing the right number of atoms at the right places, so it all adds up.

Also, my teacher says that sulphuric acid is a stronger acid then hydrochloric acid, because it has more hydrogen atoms in one molecule(roughly like that). But H2SO4 has a lot of bonds, so for it to dissociate in water, it needs a lot of energy right? But for HCl, there is only 1 bond, so it doesn't need a lot of energy to dissociate. So shouldn't my teacher's statement be true only at high temperatures?
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FZ+
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#2
Sep20-03, 07:06 PM
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When atoms combine to form compounds, they combine to form compounds with zero valence right?
They don't.

It is worthwhile to note that atoms don't try to do anything. They don't hold chemistry textbooks. What atoms do, in the process of random motion, is to end up in states of high stability. Usually, this is with neutral valencies, but in certain situations, they can instead form complex ions. It is true that conservation of charge applies, so the charge of the whole lot is still neutral, but it is possible to seperate them.

Note the difference between more covalent bonding, in which case we can generally distinguish individual molecules, and more ionic forms of bonding in which case the whole idea of "compound" becomes a bit nebulous as we just get a rough mix of charged particles.

And since the bonds are all essentially metallic, covalent and ionic, then ultimately couldn't i have any combination of elements, example KMgFeBOHSe, after all, it would just be a matter of placing the right number of atoms at the right places, so it all adds up.
Yes, but can you expect it all to stay together for longer than a nanosecond? ie. be stable.
On Radioactive Waves
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#3
Sep21-03, 12:20 PM
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NaSO4 dissociates 4 times, so you get 4 times the hydronium ions when it completly dissociates as compared to HCl

edit: ummmmm- this was late at night, and I dont know what I was thinking.

radagast
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#4
Sep22-03, 12:07 PM
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Nagging problems


Originally posted by On Radioactive Waves
NaSO4 dissociates 4 times, so you get 4 times the hydronium ions when it completly dissociates as compared to HCl
I have no idea what you're trying to get across, but:

NaSO4 doesn't exist.

Na2SO4 does exist and

NaHSO4 exists and

NaSO4- exists.


NaHSO4 can dissociate into Na+, [ HSO4- OR SO42- and H3O+ ], but that's a max of only one hydronium ion.

Na2SO4 doesn't produce any hydronium ions,

and pure sulfuric acid, in water, H2SO4 can dissociate into a maximum of two hydronium ions, and that's the most you'll get from one of these molecules.
HazZy
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#5
Sep23-03, 12:12 PM
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however the second dissociation is sort of minor.
radagast
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#6
Sep23-03, 12:44 PM
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And don't even talk about the third dissociation of phosphoric...
Another God
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#7
Sep28-03, 04:44 AM
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Originally posted by Bubonic Plague

Also, my teacher says that sulphuric acid is a stronger acid then hydrochloric acid, because it has more hydrogen atoms in one molecule(roughly like that).
Chemistry is not my strong point, but I think the answer to this would have to do with the fact that H2SO4 has two hydrogen ions which can be dissociated, while HCl only has one. They are both Strong acids, so the first H dissociates completely: That is, if you mix these acids with water, than all of the H ions dissociate. If you had equal amounts of each type of acid, then at this first dissociation, they would be equal strength. But H2SO4 has a second H doesn't it? After losing the first H, it becomes HSO4- which is a weak acid. As such, it doesn't dissociate completely, only partially. This partial dissociation is enough though, to make the final solution more acidic then tha HCl solution.


(That makes sense to me....in my experience though, what makes sense to me, in chemistry, is very rarely the correct answer [:(]. So I would await Bystander of CSF or Boulderheads affirmation on that one)
FZ+
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#8
Sep28-03, 03:55 PM
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So shouldn't my teacher's statement be true only at high temperatures?
I am not really sure about this, but when we gauge the strength of acids, we refer to it's acidity at standard conditions of 25C, or so. Thus, it may well be that at that temperature, H2SO4 is better as a proton donor than HCl.

H2SO4(aq) <-> 2H(aq) (+) + SO4(aq) (-)

What is significant is the position of this equilibrium at various temps. At higher temps, the equilibrium should decrease the pH value of the acidic solution.

EDIT: Damn, I keep making this mistake... Low pH = more H+...
Chemicalsuperfreak
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#9
Sep28-03, 05:23 PM
P: 324
Originally posted by Another God
Chemistry is not my strong point, but I think the answer to this would have to do with the fact that H2SO4 has two hydrogen ions which can be dissociated, while HCl only has one. They are both Strong acids, so the first H dissociates completely: That is, if you mix these acids with water, than all of the H ions dissociate. If you had equal amounts of each type of acid, then at this first dissociation, they would be equal strength. But H2SO4 has a second H doesn't it? After losing the first H, it becomes HSO4- which is a weak acid. As such, it doesn't dissociate completely, only partially. This partial dissociation is enough though, to make the final solution more acidic then tha HCl solution.


(That makes sense to me....in my experience though, what makes sense to me, in chemistry, is very rarely the correct answer [:(]. So I would await Bystander of CSF or Boulderheads affirmation on that one)
That's correct, sir. Often when dealing with acids, people will deal with the concept of "normality" instead of molarity. Normality is just molarity times the number of donatable hydrogen atoms. So a 1M solution of HCl is 1N, while a 1M solution of H2SO4 is 2N. Which sort of implies that H2SO4 is twice as acidic as HCl, and for some purposes this holds true; say for example a chemist wanted to neutralize a strongly basic solution.
Chemicalsuperfreak
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#10
Sep28-03, 05:24 PM
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Originally posted by FZ+
I am not really sure about this, but when we gauge the strength of acids, we refer to it's acidity at standard conditions of 25C, or so. Thus, it may well be that at that temperature, H2SO4 is better as a proton donor than HCl.

H2SO4(aq) <-> 2H(aq) (+) + SO4(aq) (-)

What is significant is the position of this equilibrium at various temps. At higher temps, the equilibrium should increase the pH value of the acidic solution.
Careful, are you saying the equilibrium would go to the left? Because if it goes to the right the pH would go down. pH = -log[H+]


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