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Analytical Chemistry 2

الكلية كلية العلوم للبنات     القسم قسم الكيمياء     المرحلة 1
أستاذ المادة محمد حامد سعيد الدهيمي       07/01/2017 17:54:20

LECTURE 2
Acid- Base Equilibrium in Water
The chemistry of aqueous solutions is dominated by the equilibrium between neutral water molecules and the ions they form.
2 H2O(l) <---> H3O+(aq) + OH-(aq)
Strict application of the rules for writing equilibrium constant expressions to this reaction produces the following result.
This is a legitimate equilibrium constant expression, but it fails to take into account the enormous difference between the concentrations of neutral H2O molecules and H3O+ and OH- ions at equilibrium.
Measurements of the ability of water to conduct an electric current suggest that pure water at 25oC contains 1.0 x 10-7 moles per liter of each of these ions.
[H3O+] = [OH-] = 1.0 x 10-7 M
At the same temperature, the concentration of neutral H2O molecules is 55.35 molar.
The ratio of the concentration of the H+ (or OH-) ion to the concentration of the neutral H2O molecules is therefore 1.8 x 10-9.


In other words, only about 2 parts per billion (ppb) of the water molecules dissociate into ions at room temperature.

The equilibrium concentration of H2O molecules is so much larger than the concentrations of the H3O+ and OH- ions that it is effectively constant. We therefore build the [H2O] term into the equilibrium constant for the reaction and thereby greatly simplify equilibrium calculations. We start by rearranging the equilibrium constant expression for the dissociation of water to give the following equation.

[H3O+][OH-] = Kc x [H2O]2

We then replace the term on the right side of this equation with a constant known as the water dissociation equilibrium constant, Kw.

For more information See the attached file


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