PH, POH, AND K

Adding an acid to water increases the H3O+ion concentration and decreases the OH- ionconcentration. Adding a base does the opposite. Regardless ofwhat is added to water, however, the product of theconcentrations of these ions at equilibrium is always 1.0 x 10-14at 25oC.

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=1.0 x 10-14

The table below lists pairs of H3O+ andOH- ion concentrations that can coexist at equilibriumin water at 25oC.

Pairs of EquilibriumConcentrations of H3O+and OH- Ions That Can Coexist in Water

Data from this table are plotted in the figure below over anarrow range of concentrations between 1 x 10-7 Mand 1 x 10-6 M. The point at which theconcentrations of the H3O+ & OH-ions are equal is called the neutral point. Solutions inwhich the concentration of the H3O+ ion islarger than 1 x 10-7 M are described as acidic.Those in which the concentration of the H3O+ion is smaller than 1 x 10-7 M are basic.

It is impossible lớn construct a graph that includes all thedata from the table given above. In 1909, the Danish biomagdalenarybarikova.comistS. P. L. Sorenson proposed using logarithmic mathematics tocondense the range of H3O+ và OH-concentrations to lớn a more convenient scale. By definition, thelogarithm of a number is the power lớn which a base must be raisedto obtain that number. The logarithm to the base 10 of 10-7for example, is -7.

log (10-7) = -7

Since the concentrations of the H3O+ andOH- ions in aqueous solutions are usually smaller than1 M, the logarithms of these concentrations are negativenumbers. Because he considered positive numbers more convenient,Sorenson suggested that the sign of the logarithm should bechanged after it had been calculated. He therefore introduced thesymbol "p" to lớn indicate the negative of thelogarithm of a number. Thus, pH is the negative of thelogarithm of the H3O+ ion concentration.

pH = - log

Similarly, pOH is the negative of the logarithm of theOH- ion concentration.

pOH = - log

pH + pOH = 14

The equation above can be used khổng lồ convert from pH lớn pOH, orvice versa, for any aqueous solution at 25C, regardless of howmuch acid or base has been added to lớn the solution. By convertingthe H3O+ & OH- ionconcentrations in the table above into pHand pOH data, we can fit the entire range of concentrations ontoa single graph, as shown in the figure below.

Acid-DissociationEquilibrium Constants

There is a big difference between strong acids such ashydrochloric acid & weak acids such as the acetic acid invinegar. Both compounds satisfy the Brnsted definition of anacid. (They are both H+ ion, or proton, donors.) Butthey differ in the extent to which they donate H+ ionsto water.

By definition, a strong acid is any substance that is good atdonating an H+ ion to water.

Example: 99.996% of the HCl molecules in a 6 M solutiondissociate when the following reaction comes khổng lồ equilibrium. Thisequilibrium lies so far to the right that we write the equationfor the reaction with a single arrow, suggesting thathydrochloric acid dissociates more or less completely in aqueoussolution.

HCl(aq)	+	H2O(l)	" width="17" height="9">	H3O+(aq)	+	Cl-(aq)
0.004%				99.996%
at equilibrium				at equilibrium

Weak acids are relatively poor H+ ion donors.

Example: Acetic acid is a Brnsted acid because it can donatean H+ ion to water. But it isn"t a very good H+ion donor. Only about 1.3% of the acetic acid molecules in an0.10 M solution lose a proton to lớn water.

A quantitative feeling for the difference between strong acidsand weak acids can be obtained from the equilibrium constants forthe reactions between acids and water. Because it istime-consuming to write the formula CH3CO2Hfor acetic acid, magdalenarybarikova.comists commonly abbreviate this formula asHOAc and describe the dissociation of the acid as follows.

HOAc(aq) + H2O(l)

Using this convention, the equilibrium constant expression forthe reaction between acetic acid và water would be written asfollows.

Like the equilibrium constant expression for the dissociationof water, this is a legitimate equation. But most acids are weak,so the equilibrium concentration of H2O is effectivelythe same after dissociation as before the acid was added. Becausethe term has no effect on the equilibrium it isbuilt into the equilibrium constant for the reaction as follows.

The result is an equilibrium constant for this equation knownas the acid-dissociation equilibrium constant, Ka.For this reaction:

In general, for any acid HA:

Values of Ka can be used khổng lồ estimatethe relative strengths of acids. The larger the value of Ka,the stronger the acid. By definition, a compound is classified asa strong acid when Ka is larger than 1.Weak acids have values of Ka that aresmaller than 1. A danh sách of the acid-dissociation equilibriumconstants for some common acids is given in the table below.

Values of Ka forCommon Acids

The table above provides us with the basis for understandingthe difference between strong acids & weak acids. Think aboutthe reaction between a very strong acid và water.

HCl is a much stronger acid than the H3O+ion. This means that H2O is a stronger base than theCl- ion. It isn"t surprising to find that the strongerof a pair of acids reacts with the stronger of a pair of bases togive a weaker acid and a weaker base.

Let"s consider the reaction between acetic acid and water.

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In this case, the reaction tries to convert the weaker of apair of acids & the weaker of a pair of bases into a strongeracid & a stronger base. It isn"t surprising to lớn find that thisreaction occurs lớn only a minor extent.

As the value of Ka decreases further the extent towhich the acid will react with water must decrease as well.Inevitably, we should encounter acids that are so weak they can"tcompete with water as a source of the H3O+ion.