SplendidAcylation
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Silver nitrate + hydrochloric acid - metathesis, or acid-base reaction?
Hi,
I have recently been playing with silver salts, in the process of quantitative determination of cyanamide ion in my homemade calcium cyanamide (a
write-up is on its way)
Anyway, something was bothering me slightly, and I was kept awake at night simulating various precipitation reactions in my head...
"Does silver chloride react with nitric acid?"
My first assumption was that it would definitely react, "hydrochloric acid being a weaker acid than nitric acid"
But this is not so!
Nitric acid is actually much weaker an acid than hydrochloric acid, as can be seen here
So if this reaction does not take place, the opposite reaction must take place, and, upon investigation, sure enough, it does:
Silver nitrate solution reacts with hydrochloric acid, forming a precipitate of silver chloride, and leaving in solution the nitric acid.
But what is bothering me now, is how no-one refers to this as an acid-base reaction.
The overall reaction, with state symbols, is:
AgNO3 (aq) + HCl (aq) -> HNO3 (aq) + AgCl (s)
Upon closer examination, this appears to me to be an acid-base reaction:
NO3- + HCl <--> HNO3 + Cl-
The forward reaction is favoured, as HCl is the stronger acid.
As chloride ions form, they are removed from the equilibrium by the formation of solid AgCl, which drives the equilibrium forward.
However, the position of equilibrium would still lie heavily to the right, even if no precipitate formed.
However, this reaction is often referred to instead as a double-displacement, or metathesis, reaction, with no mention of acid-base reactions, see
this thread for example:
https://chemistry.stackexchange.com/questions/15394/reaction...
It is stated "Formation of the solid is the driving force for this reaction"
This makes sense, and indeed the formation of the solid pushes the reaction to completion, however it does not apply in a seemingly identical
scenario, imagine we replace the hydrochloric acid with water:
AgNO3 (aq) + HOH (aq) -> HNO3 (aq) + AgOH(s)
It seems, based on the assumption that the formation of a solid drives the reaction forward, that this would take place, however, of course, it
doesn't, as silver hydroxide would react immediately with nitric acid.
In this case, the acid-base equilibrium lies heavily to the left:
NO3- + HOH <--> HNO3 + OH-
Therefore, it seems to me to be the case that two factors cause the silver nitrate + hydrochloric acid reaction to take place:
Hydrochloric acid is the stronger acid
The product is insoluble
If we replaced the hydrochloric acid with sodium chloride, then we would have a pure metathesis reaction, where the reaction was governed solely by
the insolubility of the product.
So, to be nit-picky, it is incorrect to say this is a metathesis reaction...
So what do you think, am I correct, or am I just being a fuss-pot? 
[Edited on 28-1-2023 by SplendidAcylation]
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Tsjerk
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When you say "strong" acid, you say the Ka is > 1. This means that at concentrations we are talking about here, close to 100% of the acid is
dissociated. So it doesn't really matter which acid is stronger; both are strong.
The equilibrium is effected though by AgCl precipitation, this takes it out of the equation, driving the reaction towards AgCl and nitric acid.
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blogfast25
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Basically the latter
This is a simple ionic precipitation reaction: don't overthink it!
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CharlieA
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I generally think of an acid-base reaction as producing a salt + water.
And I think (as I taught my students many years ago) of a double metathesis reactions is swap of ions (or as I used to teach "a Bob and Carol and Ted
and Alice" reaction.
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DraconicAcid
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A metathesis reaction is a swapping of ions. This includes both acid-base reactions and precipitation reactions.
"Acid and base give salt and water" only works if you stick to Arrhenius' definition of a base, which is a hydroxide and only a hydroxide. A useless
definition even for high school chemistry.
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
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Rainwater
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Just throwing this out there to complicate things a little more
| Quote: | | AgNO3 (aq) + HCl (aq) -> HNO3 (aq) + AgCl (s) |
This is true in diluted solutions.
If you used an exact stoichiometric amount of pure Ag metal, 98% RFNA and 32% HCl,
in a test tube at stp protected from water.
It will take a lot of time but you will see the silver dissolve into solution,
bubbles H2 gas and be left with a
Clear solution.
Purified KNO3 can be used as a nitric source aswell.
As soon as you add water, you will get a percipitate.
If you impact the test tube hard enough you will get a percipitate.
Quantitative Chemical Analysis by: Harris, Daniel C.
ISBN: 9781429218153 will explain why
"You can't do that" - challenge accepted
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SplendidAcylation
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Thanks for all the responses.
I was going to reply by saying that weaker acids than nitric acid, such as acetic acid, will not precipitate silver salts from a solution of silver
nitrate, however, to my immense surprise, they will!
Even weak acids like oxalic acid and acetic acid, when added to a solution of silver nitrate, will cause silver oxalate/acetate to precipitate,
leaving nitric acid in solution!
This is mind-blowing to me, I would have definitely assumed that the reverse would be true, i.e., that silver acetate would react with nitric acid,
since nitric acid is a much stronger acid.
However, evidently, it seems that the formation of a precipitate will not always drive a reaction to completion, after all, if it did, then
no insoluble base would ever dissolve in acid!
If we add a sufficiently weak acid to a solution of silver nitrate, such as water or carbon dioxide, we will not obtain a precipitate of silver
hydroxide, or carbonate.
Indeed, the reverse reaction does take place, of course, with silver hydroxide or silver carbonate readily dissolving in nitric acid, as
would be expected.
There is no qualitative difference between water, as an acid, and acetic acid, except that acetic acid is a much stronger acid.
Would it therefore be somewhat correct, as a sweeping statement, to say that the formation of a precipitate drives the reaction forward, only if the
difference in strengths of the acid added, and the acid left behind in the solution, is not too great.
At some point, the acid-base equilibrium, rather than the formation of a precipitate, must become the primary factor in driving the reaction in one
direction or the other, once the difference in Ka becomes too great.
I'm fairly good at understanding general concepts, such as Le Chatelier's principle, but not so good at mathematical aspects, so I really have no
intention of calculating the positions of equilibrium when a weak acid is added to a solution of silver nitrate...
I overthink everything!
Quote: Originally posted by DraconicAcid  | A metathesis reaction is a swapping of ions. This includes both acid-base reactions and precipitation reactions.
"Acid and base give salt and water" only works if you stick to Arrhenius' definition of a base, which is a hydroxide and only a hydroxide. A useless
definition even for high school chemistry. |
An oxide would also class as a base by this definition! 
@Rainwater, that's interesting.
How can you have 32% HCl in the absence of water, isn't this 68% water by mass?
Another thing that interests me greatly is the acidic nature of transition-metal salts.
A solution of silver nitrate, for instance, is acidic.
The most basic explanation of this usually goes something like:
"A salt of a weak base and a strong acid is acidic"
This explanation works with AgNO3, NH4Cl, etc
The antithesis being:
"A salt of a strong base and a weak acid is basic"
This explanation works with CH3COONa, NaF, etc.
However, in the case of AgNO3 and NH4Cl, the cause of the acidity of these solutions is quite different.
In the case of NH4Cl, the acidity is due to the ammonium ion being a weak acid, protonating water:
NH4+ + H2O <--> NH3 + H3O+
In the case of AgNO3, however, there is no source of protons to donate!
This is often explained by the equilibrium:
Ag+ + H2O <--> AgOH + H+
However this still doesn't explain the phenomenon...
I believe that the actual explanation has to do with the formation of metal aquo complexes:
M(H2O)6 + <--> M(H2O)5-OH + H+
Also, another problem with "A salt of a weak base and a strong acid is acidic", is the definition of a "weak base" in this context.
Both AgOH and NaOH contain hydroxide ions, yet AgNO3 is acidic and NaNO3 is neutral.
How can AgOH be a weak base, if they both contain the same anion?
In reality, the idea of AgOH being a weak base fits nicely with its observed properties, it doesn't form alkaline solutions (because it is insoluble),
it isn't caustic, etc, but it isn't really a weak base, it's just insoluble.
However, insoluble bases like calcium carbonate, which would fit the above criteria for a "weak base" do not form acidic salts!
So, of course, such generalizations don't work.
Anyway, on the topic of metal aquo complexes, you can see that removing the H+ from the above equilibrium will push the equilibrium to the right,
towards the hydroxy derivative.
Would it be reasonable, then, to guess that the following reaction will take place:
AgCl(s) + OH- (aq) --> AgOH (s) + Cl- (aq)
Since a solid is reacting, forming a solid precipitate, it isn't so easy to see how this would take place, however, when silver chloride is added to
water, a solubility equilibrium is quickly set up:
AgCl (s) <--> Ag+ (aq) + Cl- (aq)
The Ag+, in turn, hydrates, forming an aquo complex:
Ag(H2O)6 + (probably!)
Which then dissociates, forming H+
If we add hydroxide ions to this system, the H+ will be removed, pushing the equilibrium towards Ag(H2O)5OH
Since AgOH readily converts to Ag2O, it would be the oxide that would precipitate...
My mind gets a bit fuzzy when it comes to figuring out how Ag(H2O)5OH turns into Ag2O... Presumable it loses the water of hydration, and then two
moles of AgOH lose one mole of water, yielding Ag2O?
I believe this type of reaction also takes place when insoluble Hg2Cl2 is added to a solution containing hydroxide ions, however I can't find any
reference to this...
[Edited on 29-1-2023 by SplendidAcylation]
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DraconicAcid
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| Quote: Originally posted by SplendidAcylation | Thanks for all the responses.
I was going to reply by saying that weaker acids than nitric acid, such as acetic acid, will not precipitate silver salts from a solution of silver
nitrate, however, to my immense surprise, they will!
Even weak acids like oxalic acid and acetic acid, when added to a solution of silver nitrate, will cause silver oxalate/acetate to precipitate,
leaving nitric acid in solution! |
It depends on the relative concentrations.
The Ksp for silver acetate (AgC2H3O2(s) == Ag+ + C2H3O2-) is 1.9 x 10-3.
Ka for acetic acid (HC2H3O2 == H+ + C2H3O2- )is 1.8 x 10-5. (I'm too lazy to type out hydronium ions.)
So the reaction of silver ions with acetic acid:
Ag+ + HC2H3O2 == AgC2H3O2(s) + H+
has an eq'm constant of 0.0095
It's a small number, but it's not zero.
Silver oxalate has a Ksp of 5 x 10-12. Ka2 values for oxalic acid 5.4 x 10-5.
2 Ag+ + HC2O4-(aq) == Ag2C2O4(s) + H+ K = 1.1 x 10+7.
If you add silver nitrate to a solution of oxalic acid or sodium hydrogen oxalate, you'll get practically all of the silver precipitating. But if you
add concentrated acid, you'll force the eq'm back to dissolve the silver.
| Quote: Originally posted by DraconicAcid | A metathesis reaction is a swapping of ions. This includes both acid-base reactions and precipitation reactions.
"Acid and base give salt and water" only works if you stick to Arrhenius' definition of a base, which is a hydroxide and only a hydroxide. A useless
definition even for high school chemistry. |
An oxide would also class as a base by this definition!
That's true, but I've met several otherwise-competent chemists who insisted that oxides were "basic anhydrides", rather than actual bases!
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
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Rainwater
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You cant. Yes it is.
The precipitation requires an abundance of water.
The way I understand it, as the metal began to dissolve,
it forms a super saturated solution
which is sustained by the excess of Cl in solution thanks to the common ion effect
Even tho the exact amounts of reagents are in solution, all the metal will not dissolve because some of the ions are to busy keeping the solution
saturated.
If extra water is added, the reaction will go to completion.
"You can't do that" - challenge accepted
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SplendidAcylation
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| Quote: Originally posted by DraconicAcid |
It depends on the relative concentrations.
The Ksp for silver acetate (AgC2H3O2(s) == Ag+ + C2H3O2-) is 1.9 x 10-3.
Ka for acetic acid (HC2H3O2 == H+ + C2H3O2- )is 1.8 x 10-5. (I'm too lazy to type out hydronium ions.)
So the reaction of silver ions with acetic acid:
Ag+ + HC2H3O2 == AgC2H3O2(s) + H+
has an eq'm constant of 0.0095
It's a small number, but it's not zero.
Silver oxalate has a Ksp of 5 x 10-12. Ka2 values for oxalic acid 5.4 x 10-5.
2 Ag+ + HC2O4-(aq) == Ag2C2O4(s) + H+ K = 1.1 x 10+7.
If you add silver nitrate to a solution of oxalic acid or sodium hydrogen oxalate, you'll get practically all of the silver precipitating. But if you
add concentrated acid, you'll force the eq'm back to dissolve the silver.
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Thanks for the explanation!
I'll re-read this and figure out exactly how you're calculating these things.
It's probably simple, I wish I were better at mathematical operations!
| Quote: Originally posted by DraconicAcid |
That's true, but I've met several otherwise-competent chemists who insisted that oxides were "basic anhydrides", rather than actual bases!
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Hmmm, I've heard of the "basic anhydride" description, indeed, there won't be much oxide ion in solution, mostly converting to hydroxide ions, still,
oxides do behave as a base, for instance if you were to add sodium oxide to glacial acetic acid, I'd imagine you'd end up with sodium acetate and
water!
One of my favourite acid-base reactions that does not fit into the Brønsted–Lowry theory is:
CaO + CO2 <--> CaCO3
It's obviously an acid-base reaction, yet there are no protons! 
I was also confused for a while about how CO2 can act as a Lewis acid when it has a full octet on all its atoms, eventually I figured out it has to do
with a resonance contributor that has an electron deficient carbon.
Resonance is another thing I would like to understand better, that is, predicting resonance structures.
[Edited on 2-2-2023 by SplendidAcylation]
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Texium
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| Quote: Originally posted by SplendidAcylation | | Quote: Originally posted by DraconicAcid | That's true, but I've met several otherwise-competent chemists who insisted that oxides were "basic anhydrides", rather than actual bases!
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Hmmm, I've heard of the "basic anhydride" description, indeed, there won't be much oxide ion in solution,
mostly converting to hydroxide ions, still, oxides do behave as a base, for instance if you were to add sodium oxide to glacial acetic acid, I'd
imagine you'd end up with sodium acetate and water! |
As you would if you were to add sodium hydroxide to acetic anhydride.
It’s all semantics. Almost everything you’re talking about in this thread is. If you stretch the definition enough, you can call practically
anything an acid-base reaction, because by the Lewis definition the only requirement is that a pair of electrons are moved from one atom to another…
which approaches the very definition of a chemical reaction. The only true exceptions are free-radical reactions and pericyclic reactions
(Diels-Alder, sigmatropic reactions, etc).
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