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Why do strong acids dissociate completely in water?

by aleksbooker
Tags: acids, completely, dissociate, strong, water
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aleksbooker
#1
Mar4-14, 06:00 PM
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I'm confused about the dissolution of strong acids (like HCl) in water.

We're told that they dissolve completely because
  1. They are strong electrolytes.
  2. The hydrogen atom that is lost during dissociation is *not* strongly bound to the rest of the acid molecule.
  3. Therefore, the solvent (usually water) pulls at the H+ atom more strongly than the rest of the acid molecule.
As far as I can tell, this is completely contradictory. If it's a strong electrolyte, wouldn't the rest of the acid molecule (the Cl- part) have a *stronger* hold on the H+ ion than a weaker acid/weaker electrolyte? Wouldn't it be *harder* to separate the H+ ion from the molecule?
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Chestermiller
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Mar4-14, 06:13 PM
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Quote Quote by aleksbooker View Post
I'm confused about the dissolution of strong acids (like HCl) in water.

We're told that they dissolve completely because
  1. They are strong electrolytes.
  2. The hydrogen atom that is lost during dissociation is *not* strongly bound to the rest of the acid molecule.
  3. Therefore, the solvent (usually water) pulls at the H+ atom more strongly than the rest of the acid molecule.
As far as I can tell, this is completely contradictory. If it's a strong electrolyte, wouldn't the rest of the acid molecule (the Cl- part) have a *stronger* hold on the H+ ion than a weaker acid/weaker electrolyte? Wouldn't it be *harder* to separate the H+ ion from the molecule?
This is just a matter of semantics. A strong electrolyte in one that completely dissociates so that the dissociated ions can support an electric current by migrating toward the electrodes. That's why it's called an electrolyte. A weak electrolyte is one that cannot support an electric current as well. That's because the solution is not fully ionized, so that they cannot migrate toward the electrodes.

Chet
aleksbooker
#3
Mar4-14, 06:15 PM
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So, by nature, the molecules of strong electrolytes have *weaker* holds on their protons? That way more of them can disassociate into solution and actually act as electrolytes.

And more, strong acids have weaker holds on their protons (H+ ions), while weak acids have stronger holds on their protons.

But I thought the whole point of a halide (e.g. Cl-) and a metal ion (H+) getting together was that they had *strong* electronegativity, and *didn't* separate easily.

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Mar4-14, 06:45 PM
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Why do strong acids dissociate completely in water?

Quote Quote by aleksbooker View Post
So, by nature, the molecules of strong electrolytes have *weaker* holds on their protons? That way more of them can disassociate into solution and actually act as electrolytes.

And more, strong acids have weaker holds on their protons (H+ ions), while weak acids have stronger holds on their protons.

But I thought the whole point of a halide (e.g. Cl-) and a metal ion (H+) getting together was that they had *strong* electronegativity, and *didn't* separate easily.
I don't know how to answer your last question regarding ionic bonding, because my chemistry background is not strong enough. Guess we'll have to wait for someone more knowledgeable to help. Sorry.

chet
Yanick
#5
Mar4-14, 07:28 PM
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The trick is to consider the medium you are thinking of the dissociation happening in. You can look up the Coulomb Potential or Coulomb's Law and a table of dielectric constants to see how the dissociation of ions is favored in a medium such as water as opposed to in a vacuum or gas phase.
aleksbooker
#6
Mar4-14, 09:09 PM
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Hmm... I'll save that for a later date. I want to know, but it's not essential to this class and I unfortunately can't spare the time to research. Hopefully I'll run into it again later, though. Appreciate your help guys!
Borek
#7
Mar5-14, 02:55 AM
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Actually HCl is quite covalent, not ionic. However, when water dipoles come into account, they change the picture, as they solvate both H+ and Cl- and shift the dissociation far to the right.

Edit: BTW, the way you asked the question is putting things on the head. While we say "strong acids dissociate fully in the water" what we really do is we classify those acids that dissociate fully as strong acids - so it is dissociation first, classification as strong later. Answer to the question "why do strong acid dissociate fully?" is "because we decided to call substances that behave this way strong acids".
Yanick
#8
Mar5-14, 08:22 AM
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Quote Quote by Borek View Post
Actually HCl is quite covalent, not ionic. However, when water dipoles come into account, they change the picture, as they solvate both H+ and Cl- and shift the dissociation far to the right.
Isn't this essentially the same as saying that water has a higher dielectric constant compared to vacuum or air thus favoring dissociation molecules to form ions? Or am I missing something?

Quote Quote by Borek View Post
Edit: BTW, the way you asked the question is putting things on the head. While we say "strong acids dissociate fully in the water" what we really do is we classify those acids that dissociate fully as strong acids - so it is dissociation first, classification as strong later. Answer to the question "why do strong acid dissociate fully?" is "because we decided to call substances that behave this way strong acids".
That's a great way of putting it.
Borek
#9
Mar5-14, 08:29 AM
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Quote Quote by Yanick View Post
Isn't this essentially the same as saying that water has a higher dielectric constant compared to vacuum or air thus favoring dissociation molecules to form ions? Or am I missing something?
Related, but different IMHO. First of all, dielectric constant of water is a property of a bulk liquid, with randomly oriented molecules. When it comes to solvation water molecules become ordered, so in the microscale the local dielectric constant has a different value. Plus, ordering these molecules means strong entropic and enthalpic effects.
abitslow
#10
Mar6-14, 09:22 AM
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Borek nailed it. A strong acid is DEFINED as one which will completely dissociate in the medium (solvent) under discussion. (So, while water is the "normal" solvent considered, it could be something else; like ethanol(anhydrous), gasoline, or even a solid or a gasseous "solvent") (Confusingly, some will call the acid "strong" if it is strong in water, even if it is in some other solvent).
Now to your 3 points:
1. Exactly, keep in mind that there are 3 major definitions of what an acid is. They aren't the same. Bronsted, Lewis, and Arrhenius. The Arrhenius definition is the one you first learn about and it is specific for water solutions and ONLY water solutions. So your first point is correct just like a square is a rectangle but not necessarily the other way around. An electrolyte is not necessarily an arrhenius acid.
2. This is pretty much wrong. The fact is that for the reaction HA → H(+) + A(-) to occur spontaneously the energy ("free energy" which includes an entropy term, sorry) of the products must be MORE than the energy of the reactants (which must include the water, although it is not shown in the "net reaction" above, but has an extremely important effect on the reaction (obviously)). So the energy keeping the reactants together just has to be less than keeping the products "together" (actually, keeping them as ions in solution - the ions are surrounded by water molecules and it is these complexes that stabilize the ions and contribute the energy of formation). So the amount of energy can be small as long as the products' is just a bit more, or the amount can be HUGE, as long as the products' are just a bit more. It is relative. Its their difference you should think about, not their absolute value. (Their absolute value comes in when you're doing the actual mixing in the lab, if the energy is large then you need to be really careful not to go boom :)
3. Is nonsense. Well, almost complete nonsense. As I explained in #2 it is the total free energy that is important. If the energy of the solvated anion is small, that contributes very little to the total for the products and so makes it less likely that the products' energy will be more than the reactants ...meaning that it is less likely that the products will form...meaning less likely to be strong. If you think about it, you will realize that an H(+) ion is an H(+) ion, so if #3 were right, it wouldn't matter whether you were talking about HBr or HCl or H2O or CH4....they would all be strong acids since they all have H's to ionize (to be "pulled" into solution). The fact is the more the anion (Br(-), Cl(-), HO(-) or CH3(-) ) is pulled, the more acidic it will be...there's nothing that says the H(+) needs to be more strongly "pulled" than the A(-).
aleksbooker
#11
Mar6-14, 02:46 PM
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This is awesome. I'm learning so much (or at least, I think I am). Here's what I've gathered from you guys so far:

HCl and other compounds that dissociate completely in solution are called strong acids because of this behavior, not the other way around.

This behavior (complete dissociation) is caused by the relative difference between the energy holding the reactants together (enthalpy of dissolution?) and the energy holding the products together (enthalpy of formation?). If the energy holding the products together is higher, than the reaction favors the products and the compound will dissociate.

The "pull" of the solvent/energy of solvation (I'm assuming these all mean the same thing in this context) contribute most to these enthalpies. Since the H+ ion and its energy of solvation is the same whether it comes from HBr, HCl, HNO3, that particular reaction is not the determining factor in whether or not a compound will dissociate completely into its component ions. Instead, its the energy of solvation of the A(-) anion. If the anion has an enthalpy of solvation with a high enough magnitude that it, combined with the solvation of the H+ anion, is greater than the magnitude of the energy holding the compound together, *that's* when the compound dissociates.

If that's all true, I'm still somewhat back at the original question. Why does HCl dissociate completely while H2O or HC2H3O2 (acetic acid) do not? Is it that Cl- has a greater energy of solvation, so that:

1. It dissociates more readily and is held more tightly by the solvent?
2. Because it's held so tightly, it is less likely to "react backwards" to reform aqueous HCl?

Whereas the anions of H2O and acetic acid, OH- and C2H3O2 have lower energies of solvation, which means:

1. It's harder for the solvent to separate them from the compound and into solvated ions?
2. The ions that are solvated are more susceptible to recombining to form H2O and acetic acid?
Yanick
#12
Mar6-14, 03:28 PM
P: 380
abitslow made a mistake when he was explaining the free energy concept. All systems tend toward the lowest possible energy configuration. The statement that the products have to have a higher energy than the reactants is exactly backwards, unless something else was meant.

ΔG = ΔH + TΔS describes the free energy of a system, the equation has two terms. The first involves the enthalpy which, on a microscopic scale, can refer to electrostatic interactions between molecules, solvents etc. The second term involves temperature and entropy. Without getting too crazy we'll just consider entropy as a measure of "disorder" of a system. I sometimes explain entropy by telling people to imagine how many unique "pictures" they can take of a system, the more unique pictures the higher the entropy. Don't worry about this too much for now, you'll just confuse yourself even more. Also lets not get into the little degree sign in the equation, its a not exactly relevant to the question at hand.

In the simplest terms, a process/reaction with ΔG < 0, as written, will happen spontaneously (fancy way of saying, you just mix the stuff and it will go) while the opposite is true when ΔG > 0. When ΔG = 0, the system is already at equilibrium. As you can see a process is favorable when the ΔH term is negative and the ΔS term is positive (remember the negative sign in front of the TΔS). You can also consider how each term can "overpower" the other. For example we can have a small positive change in enthalpy but a large (also positive) change in entropy which will still lead an overall negative ΔG.

The link between an abstract quantity such as free energy (ΔG) and something tangible, like the ratio of products/reactants in a system is straightforward: ΔG = -RTlnKeq. The Keq is called an equilibrium constant and only varies with temperature. It is simply the ratio of concentrations of products divided by the concentrations of reactants, each raised to a power corresponding to their stoichiometric coefficient in a balanced chemical equation. With acids and bases chemists needed to be difficult and defined something called Ka and Kb where they each describe a very specific reaction (HA + H2O → H3O+ + A- for Ka and the reverse for Kb). Just remember that Ka is simply a special case of Keq which is linked with the abstract concept of free energy (ΔG).

Now we can start evaluating things based on energy. Chemists have come up with empirical rules for qualitatively evaluating systems and making educated guesses about what should happen. So in the case of something like HCl, we say that the free energy of the system decreases when HCl (in water) dissociates into a proton and the chloride anion, it is tough to describe how exactly the entropy behaves but we do know that the enthalpy is very negative because throwing something like HCl into water will heat up the water considerably. We can go a step further and talk about the energetics of the products, but that is always in relation to the energetics of the reactants. We can make some trends up to show that the anion which gave up the proton plays an important role in the energetics of the system, all else being equal. For haloacids (HI, HBr, HCl etc) we can talk about the charge density of the anion which decreases down the periodic (same amount of charge with a larger area). You can look up the Ka's, most likely the pKa's which are -log(Ka), for the various haloacids and the ionic radii of the anions which are produced.

If we take the example of acetic acid, we come to a charge density argument once again. A good comparison for the acetic acid case is something like ethanol. Each have a proton which can leave and an oxygen based anion is produced. Acetic acid is ~1010 more acidic than ethanol. Well, the reasoning is that the negative charge produced in acetic acid is delocalized over two oxygen atoms whereas in ethanol its "stuck" on a single oxygen. Charge density at work once again.

This can get much hairier but I will end the discussion now by reiterating that free energy is minimized in a system. If a reaction happens, such as HCl very nearly fully dissociating, then the system has a lower free energy. Dissecting the specific energetics is not really a simple process but we can look at trends and such and come up with rules of thumb but at this level its not going to be the "full picture."
aleksbooker
#13
Mar11-14, 11:32 PM
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Mmk. It took me a minute to get back to this, but here's my understanding. Strong acids disassociate completely not necessarily because their components are held "more tightly" in compounds with the solvent particles, but rather because the compounds they form with the solvent particles (or the formation of solvated ions) is much more stable than their acidic compound form.

The amount of energy required to force the components back into acid form is too high.

For weak acids, on the other hand, the difference in enthalpies is not so great, and the product is not so stable, so the compounds tend to go back and forth. Is this right, or did I misread your explanations?
DrDu
#14
Mar12-14, 03:27 AM
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P: 3,563
You should also take in mind that in hydrolysis, a strong new covalent bond is formed, namely between H and O in the hydronium ion, H3O+.


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