General characteristics of nonelectrolytes solutions. Water solutions of nonelectrolytes. Acid-base characteristics of solutions, страница 6

HCI + NH3  NH4+ + CI-

acid 2         base 1                 acid 1         base 2

Thus, it is possible to take to the group of acids and bases anions (CI -, CH3COO-), cations (H+, NH4+) and also substances not containing hydroxyl groups (for example, NH3).

Examples of conjugate acid-base pairs:

NH4+    NH3 + H+

conj.acid                 conjug. base

CH3COOH    CH3COO-  + H+

conj.acid                                   conjug. base

H3O+    H2O  + H+

conj.acid                 conjug. base

From a position of protolytic theory, the meaning acid and base has a relative character:  everything depends on the level of affinity of observed compound with proton.  In  NH3 - H2O – HF proton affinity decreases.  The water split out a proton in the mixture with ammonia, therefore, is being acid and in mixture with HF add proton and function like the base, that is to reveal amphoteric characteristics:

               H+    

NH3    +     H2O      NH4+ + OH -

                              acid

                    H+

     H2O     +     HF       H3O+   + F -

          base

          Thus,  protonic process is going in the direction of transition of proton from stronger acid (weakly holding proton) to a stronger base (stronger joining proton than conjugate base of the previous pair).

3.2. IONIC PRODUCT OF WATER

 A dissociation of water as a weak amphoteric electrolyte can be presented by the following way:

          H+

H2O + H2O  H3O+   + OH –

Usually for easy use instead of H5O+ we put H+:

H2O  H+ + OH –

According to the law of active mass, the value of constant of water dissociation express by the following way:

K H2O  = [H+] * [OH –]     =  1,8 * 10–16 (at 25°C)

                    [H2O]

Due to insignificant dissociation of water molecules the concentration of nondissociated molecules can be equate with the concentration of water molecules, that is,

[H2O] * 1000 / 18 = 55,5 mol/L

then:

[H+] * [OH –]    =K H2O [H2O]=1,8 *  10–16* 55,5 = 10–14

This value is designated Kw  and has the name of ionic product of water:

Kw = [H+] * [OH –]    = 10–14

From the equation follows that the concentration of ions of hydroxide and hydrogen never can be equal to 0.  That’s why in the solutions of even strong acids ions  OH – are present and in solutions of strong base ions  H+ are present.  The reaction of the solution can be characterized by the concentration of hydrogen ions.  The solution will have the neutral reaction if it has:

[H+] = [OH –] = √ Kw = √10–14 = 10–7 mol/L

At [H+] > 10–7 mol/L solution has got an acid reaction

At [H+] < 10–7 mol/L solution has got an alkaline reaction

Concentration of ions [H+] at known concentration of ions  [OH –]    always can be defined by ionic product of water.

3.3. HYDROGEN AND HYDROXIDE INDEXES

The degree of acidity or alkalinity of solutions is expressed not by the concentration of ions  [H+], but its common logarithm, taken with the opposite character.  This value is called hydrogen  index and designate pH.

pH = -lg[H+]

For example, in 0,025m solution of hydrochloride acid:

[H+] = 2,5*10–2 mol/L, where pH = -lg2,5*10–2 = 1,6

In neutral sphere pH = 7, because [H+] = 10–7 mol/L.  In sour sphere pH <7, in alkaline sphere pH>7.

With increase of acidity of the solution pH decreases and with increase of alkalinity  - increase.  Last values pH meet the concentration of hydrogen ions 1 M of the solution of a strong acid  (pH ≈0) and 1M of the solution of the strong base (pH ≈14).  In some cases it is better to use the value of hydroxyl index. pOH = lg[OH –] . When we find the logarithm of an equation [H+] * [OH –] = Kw, we will get pH* pOH =pKw, because at the temperature of 25°C:

pKw = -lg Kw = -lg10–14   =14, then pH+pOH=14.

The dependence between values pH an pOH and a reaction of sphere for 25°C can be expressed by the following scheme:

                                                                                                            pH=-lg[H+]

 

                    Sour sphere                     neutral sphere                  alkaline sphere

 

pH= -lg[OH –] 

3.4 TYPES OF PROTOLYTIC REACTIONS

             Consideration of various protolytic reaction (reactions of neutralization, hydrolysis, ionization) will begin with a consideration of the balance in solutions of acids and bases.  The equation of dissociation of a weak acid in water can be written by the following way:

          H2O

HA        H+ + A-

Then, according to the law of active mass the constant of an acid dissociation (Ka) can be written:

Ka = [H+] * [A –]    

          [HA]

Ka values for weak acids are quite small ( from 10–2 to 10–12) and that’s why it’s easier like the concentration of hydrogen ions express: -lg Ka =p Ka

If the constant of dissociation characterize the process of  split out of protons (alkaline reaction) then they use the index pKB.  In water solutions of weak bases (ammonia, organic amines) the equation of dissociation can be expressed like this:

B+ H2O  BH++ OH –

The constant of balance:

K = [BH+] * [OH –]    

          [B]*   [H2O]

Concentration of water is constant then,

KB = [BH+] * [OH –]    

          [B]

Ka for conjugate acid according to the equation

BH+  B + H+ will be:

Ka = [H+] * [B]

          [BH+]

Multiplying Ka and KB will get Kw

Ka * KB =[H+] * [OH –]    = Kw or pKa + pKB = pKw

For precise definition  Ka and KB values of activity should be used.