pivoxa15 said:
How does the acid know (before it dontes H+) that ...
(Certainly seems like chemistry nowadays is taught
anthropomorphically, doesn't it?
)
*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 (\Delta G) of your system can be calculated as \Delta G = \Delta H _{net} - T \Delta S _{net},
where \Delta H _{net} and \Delta S _{net} 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
\begin{gathered}<br />
{\text{HA}} \to {\text{H}}^ + + {\text{A}}^ - \hfill \\<br />
{\text{H}}^ + + {\text{H}}_{\text{2}} {\text{O}} \to {\text{H}}^ + \left( {aq} \right) \hfill \\<br />
{\text{A}}^ - + {\text{H}}_2 {\text{O}} \to {\text{A}}^ - \left( {aq} \right) \hfill \\ \end{gathered}
and so
\begin{gathered}<br />
\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 \\<br />
\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 \\ <br />
\end{gathered}
After measuring your solution's temperature, you can determine via
\Delta G = \Delta H_{net} - T\Delta S_{net}
the change in free energy of your system caused by dissociation.
A spontaneous reaction corresponds to a release of free energy, which implies \Delta G < 0.
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 \Delta H_{{\text{HA}} \to {\text{H}}^ + + {\text{A}}^ - }
), as
... 
