Triprotic Acid: Why Not Ammonia & Hyronium?

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In summary, the reason why ammonia and hydronium are not considered triprotic acids is because their conjugate bases are not stable enough due to the delocalization of charges. This makes them weaker acids compared to acids with delocalized electrons in their conjugate bases. Additionally, the stability of the conjugate base is determined by factors such as enthalpy and entropy changes, and can be calculated using the Gibbs free energy equation.
  • #1
pivoxa15
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Why isn't ammonia and hyronium considered triprotic acids? Both nitrogen and oxygen are highly electronegative so other negative charges would like to snactch their hydrogens. Is it because it would take too much energy due to its tetrahedral geometry which allows their charges to be distributed so they don't become very polar molecules?

I understand that In reality, the former is a weak base and the latter is a weak acid.
 
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  • #2
pivoxa15 said:
Why isn't ammonia and hyronium considered triprotic acids?
What is hyronium? Are you referring to the hydronium ion, H3O+?
Both nitrogen and oxygen are highly electronegative so other negative charges would like to snactch their hydrogens.
That is not a complete argument. Youu need to specifically determine what the partial positive charge on the hydrogen is, due to this electronegativity, and that's not all. In the case of NH3, there is a partial positive charge on the H atoms, though it isn't very big, being shared by three of them. But more important is the lone pair on the N atom, which would be happy to attract a proton. It is this desire to accept protons that makes NH3 a Bronsted base.

Is it because it would take too much energy due to its tetrahedral geometry which allows their charges to be distributed so they don't become very polar molecules?
It doesn't have very much to do with the actual geometry, though it is important that the partial positive charge is being shared by the three H atoms.
 
  • #3
With H3PO4 for instance, the resulting negative charges of the conjugate base is delocalized within the phosphate...this is not the case with Hydronium.
 
  • #4
GCT said:
With H3PO4 for instance, the resulting negative charges of the conjugate base is delocalized within the phosphate...this is not the case with Hydronium.

This a molecule with a delocalised electron tend to give up its Hydrogen more easily? The base of hydronium is H2O which doesn't have a delocalised electron.
 
  • #5
Delocalization increases the stability (lowers the potential energy) of the conjugate base,
thus decreasing the pKA of your acid (i.e., your acid is stronger).

Examples of acids with delocalized electrons in their conjugate bases include:
-Nitric acid (HNO3)
-Sulphuric acid (H2SO4)
-Carbonic acid (H2CO3)
-Perchloric acid (HClO3)
-etc...etc...

Delocalized electrons in conjugate bases, however, do not make acids "strong";
Take carboxylic acids (e.g., acetic acid), for example...
pivoxa15 said:
The base of hydronium is H2O which doesn't have a delocalised electron.
Well, GCT was referring to acids whose conjugate bases are charged molecules.
Given two such acids, the acid on whose conjugate base the charge is also delocalized will generally be stronger than the other acid.
 
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  • #6
bomba923 said:
Delocalization increases the stability (lowers the potential energy) of the conjugate base,
thus decreasing the pKA of your acid (i.e., your acid is stronger).

Examples of acids with delocalized electrons in their conjugate bases include:
-Nitric acid (HNO3)
-Sulphuric acid (H2SO4)
-Carbonic acid (H2CO3)
-Perchloric acid (HClO3)
-etc...etc...

Delocalized electrons in conjugate bases, however, do not make acids "strong";
Take carboxylic acids (e.g., acetic acid), for example...

GCT was referring to acids whose conjugate bases are charged molecules.
Given two such acids, the acid on whose conjugate base the charge is also delocalized will generally be stronger than the other acid.


How does the acid know (before it dontes H+) that after it donates the H+, it (the conjugate base) will be more stable (i.e have lower potential energy).
 
  • #7
pivoxa15 said:
How does the acid know (before it dontes H+) that ...
(Certainly seems like chemistry nowadays is taught anthropomorphically, doesn't it? :biggrin:)

*pivoxa15, two slightly more general questions:
How do any chemical reagents "know" how to react?
How do they "know" what will happen afterwards?

A more complete answer may involve some quantum physics, and a workable definition of what it means for a simple compound to "know" something (perhaps have some analogous "molecular cognition")...

Anyhow,
pivoxa15 said:
... after it donates the H+, it (the conjugate base) will be more stable (i.e have lower potential energy).

The explanation you may be looking for (concerning reactions proceeding to increase thermodynamic stability) may simply be an application of spontaneity-->a property of chemical reactions that release Gibbs free energy (usually as heat) and move toward a lower, more thermodynamically stable, energy state. Such reactions will proceed without any extra energy input from the surroundings into the system.

At constant temperature and pressure, the net change in the free energy ([itex]\Delta G[/itex]) of your system can be calculated as [tex]\Delta G = \Delta H _{net} - T \Delta S _{net}[/tex],
where [itex]\Delta H _{net}[/itex] and [itex]\Delta S _{net}[/itex] are the net changes in enthalpy and entropy, respectively, of your system (and T is the absolute temperature of your system (units Kelvin, obviously)).

For example, the dissociation of a general acid HA (i.e., HA→H+(aq) + A-(aq))
in aqueous solution can be considered as the sum of reactions
[tex]\begin{gathered}
{\text{HA}} \to {\text{H}}^ + + {\text{A}}^ - \hfill \\
{\text{H}}^ + + {\text{H}}_{\text{2}} {\text{O}} \to {\text{H}}^ + \left( {aq} \right) \hfill \\
{\text{A}}^ - + {\text{H}}_2 {\text{O}} \to {\text{A}}^ - \left( {aq} \right) \hfill \\ \end{gathered} [/tex]

and so
[tex]\begin{gathered}
\Delta H_{net} = \Delta H_{{\text{HA}} \to {\text{H}}^ + + {\text{A}}^ - } + \Delta H_{{\text{H}}^ + + {\text{H}}_{\text{2}} {\text{O}} \to {\text{H}}^ + \left( {aq} \right)} + \Delta H_{{\text{A}}^ - + {\text{H}}_2 {\text{O}} \to {\text{A}}^ - \left( {aq} \right)} \hfill \\
\Delta S_{net} = \Delta S_{{\text{HA}} \to {\text{H}}^ + + {\text{A}}^ - } + \Delta S_{{\text{H}}^ + + {\text{H}}_{\text{2}} {\text{O}} \to {\text{H}}^ + \left( {aq} \right)} + \Delta S_{{\text{A}}^ - + {\text{H}}_2 {\text{O}} \to {\text{A}}^ - \left( {aq} \right)} \hfill \\
\end{gathered} [/tex]

After measuring your solution's temperature, you can determine via
[tex]\Delta G = \Delta H_{net} - T\Delta S_{net} [/tex]
the change in free energy of your system caused by dissociation.

A spontaneous reaction corresponds to a release of free energy, which implies [itex]\Delta G < 0[/itex].
If this is the case in your solution, your acid will dissociate.

Regarding my previous post,
bomba923 said:
Well, GCT was referring to acids whose conjugate bases are charged molecules.
Given two such acids, the acid on whose conjugate base the charge is also delocalized will generally be stronger than the other acid.
_as delocalization would decrease [tex]\Delta H_{{\text{HA}} \to {\text{H}}^ + + {\text{A}}^ - } [/tex]
 
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  • #8
I think its more me not having properly learned chemistry.

So its to do with there being more probable the fact that dissociation occurs and it will occur. Its more probable because after occurring, it will be at a lower energy state hence its likely that it will spontaeously occur. There are reaons (at the atomic level) why it will be at a lower energy state. These reasons are why reformation occurs and the spontaeous nature of it.

Can you explain
For example, the dissociation of a general acid HA (i.e., HA→H+(aq) + A-(aq))
in aqueous solution can be considered as the sum of reactions
[tex]\begin{gathered}
{\text{HA}} \to {\text{H}}^ + + {\text{A}}^ - \hfill \\
{\text{H}}^ + + {\text{H}}_{\text{2}} {\text{O}} \to {\text{H}}^ + \left( {aq} \right) \hfill \\
{\text{A}}^ - + {\text{H}}_2 {\text{O}} \to {\text{A}}^ - \left( {aq} \right) \hfill \\ \end{gathered} [/tex]

Wouldn't you get hyrodium ions and hydroxide ions as products with the last two reactions?
 
  • #9
Continuing this line of thinking (NH3 being triprotic acid) looks like CH4 is tetraprotic ;)

Borek
 
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  • #10
But NH3 is not a triprotic acid and CH4 is not acidic not basic because the C and H are too close in electronegativities.
 
  • #11
pivoxa15 said:
Wouldn't you get hyrodium ions and hydroxide ions as products with the last two reactions?
The second reaction listed
[tex]{\text{H}}^ + + {\text{H}}_{\text{2}} {\text{O}} \to {\text{H}}^ + \left( {aq} \right)[/tex]
does produce hydronium ions.

The label (aq) denotes that the species is in aqueous solution (it is hydrated);
H+(aq) is essentially H3O+. You can usually omit the (aq) label if you write out the formula of the hydrate (H+(aq) is equivalently H3O+)
pivoxa15 said:
...and hydroxide ions as products with the last two reactions?
Why? From what would you "produce hydroxide" ?

Remember that distilled water already contains hydronium and hydroxide
(in equal concentrations of 10-7M :wink:), as
[tex]{\text{H}}_2 {\text{O}} \rightleftharpoons {\text{H}}^ + \left( {aq} \right) + {\text{OH}}^ - \left( {aq} \right)[/tex]
This is why distilled water has a pH = -log(10-7) = 7,
which we define as neutral.

Adding an acid does not "produce hydroxide";
The second reaction
[tex]{\text{H}}^ + + {\text{H}}_{\text{2}} {\text{O}} \to {\text{H}}^ + \left( {aq} \right)[/tex]
produces hydronium (your "hydrogen cation hydrate")
and the third reaction
HA→H+(aq) + A-(aq)<<(sorry no LaTeX)
produces a conjugate base hydrate.
 
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  • #12
bomba923 said:
The second reaction listed
[tex]{\text{H}}^ + + {\text{H}}_{\text{2}} {\text{O}} \to {\text{H}}^ + \left( {aq} \right)[/tex]
does produce hydronium ions.

The label (aq) denotes that the species is in aqueous solution (it is hydrated);
H+(aq) is essentially H3O+. You can usually omit the (aq) label if you write out the formula of the hydrate (H+(aq) is equivalently H3O+)

Why? From what would you "produce hydroxide" ?

Remember that distilled water already contains hydronium and hydroxide
(in equal concentrations of 10-7M :wink:), as
[tex]{\text{H}}_2 {\text{O}} \rightleftharpoons {\text{H}}^ + \left( {aq} \right) + {\text{OH}}^ - \left( {aq} \right)[/tex]
This is why distilled water has a pH = -log(10-7) = 7,
which we define as neutral.

Adding an acid does not "produce hydroxide";
The second reaction
[tex]{\text{H}}^ + + {\text{H}}_{\text{2}} {\text{O}} \to {\text{H}}^ + \left( {aq} \right)[/tex]
produces hydronium (your "hydrogen cation hydrate")
and the third reaction
[tex]{\text{A}}^ {-} + {\text{H}}_{\text{2}} {\text{O}} \to {\text{H}}^ + \left( {aq} \right)[/tex]
produces a conjugate base hydrate.
In your previous post you had for your third reaction
[tex]{\text{A}}^ {-} + {\text{H}}_{\text{2}} {\text{O}} \to {\text{A}}^ - \left( {aq} \right)[/tex]
Is your the one in your recent post wrong?

Would A- have hydronium ions surrounding it? If so should you show it in the formula?
 
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  • #13
pivoxa15 said:
In your previous post you had for your third reaction
[tex]{\text{A}}^ {-} + {\text{H}}_{\text{2}} {\text{O}} \to {\text{A}}^ - \left( {aq} \right)[/tex]
Is your the one in your recent post wrong?

Would A- have hydronium ions surrounding it? If so should you show it in the formula?

No, that's just a typo from copy->pasting the [tex]\LaTeX[/tex] too much :devil:...
(but thanks for noticing :wink:)

The third reaction (correct as originally posted) is
[tex]{\text{A}}^ {-} + {\text{H}}_{\text{2}} {\text{O}} \to {\text{A}}^ - \left( {aq} \right)[/tex]
and it produces a conjugate base hydrate.

Regarding your other question,
pivoxa15 said:
Would A- have hydronium ions surrounding it?
some hydrogen ions dissociated from acids may still be "close" to the conjugate bases simply due to charge difference.

However, remember that
strong acids have weak conjugate bases; while [tex]\Delta H_{{\text{HA}} \to {\text{H}}^ + + {\text{A}}^ - } [/tex] may be relatively lower, it may not be favorable (in terms of enthalpy) for many cations (i.e., hydronium) to be so close to one another, especially around such a weak anion (i.e., your conjugate bases).

Weak acids, on the other hand, have strong conjugate bases. As [tex]\Delta H_{{\text{HA}} \to {\text{H}}^ + + {\text{A}}^ - } [/tex] would be somewhat (relatively speaking) large, only a few hydrogen will dissociate from the acid. Compared to the previous case (stronger acid), you have a stronger anion and fewer cations; hence, your hydroniums may well be (somewhat) closer to your conjugate bases.

Remember that, without energy input from the environment, spontaneity generally decides how much dissociation will take place. While there may exist some 'optimal' distance between given hydronium and conjugate base ions (weighing the relative enthalpies and entropies associated with hydroniums existing near one another, hydroniums existing near the conjugate anion, hydroniums' distance from other dissociated conjugate base anions, etc...), dissociated hydroniums are generally free to travel anywhere within your solution.
 
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1. What is a triprotic acid?

A triprotic acid is a type of acid that has three ionizable hydrogen atoms in its chemical structure. This means that when the acid is dissolved in water, it can release three hydrogen ions (H+) into the solution.

2. How is a triprotic acid different from a monoprotic acid?

A monoprotic acid, such as hydrochloric acid, only has one ionizable hydrogen atom and can only release one H+ ion into the solution. On the other hand, a triprotic acid, like phosphoric acid, can release three H+ ions. This makes triprotic acids more acidic and able to donate more protons.

3. Why is ammonia not considered a triprotic acid?

Ammonia (NH3) is not considered a triprotic acid because it does not have any ionizable hydrogen atoms. In order for a compound to be classified as an acid, it must have a hydrogen atom that can be easily donated to a base. In ammonia, the hydrogen atom is bound too tightly to the nitrogen atom to be released as an H+ ion.

4. What is the role of hydronium in triprotic acids?

Hydronium (H3O+) is formed when a hydrogen ion (H+) reacts with a water molecule (H2O). In triprotic acids, hydronium ions are produced when the first two hydrogen ions are released. The third hydrogen ion is then released as a hydronium ion, which helps to stabilize the negative charge on the remaining ions.

5. How do triprotic acids behave in solution?

Triprotic acids behave differently in solution depending on their strength. Strong triprotic acids, like sulfuric acid, will completely dissociate in water, releasing all three H+ ions. Weak triprotic acids, like phosphoric acid, will only partially dissociate, leaving some of the H+ ions bound to the acid molecules. This behavior also affects the pH of the solution, as strong acids will have a lower pH than weak acids.

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