PH





pH is a measure of the concentration of protons (H+) in a solution and,
therefore, its acidity or alkalinity. The concept was introduced by S.P.L.
S¿rensen in 1909. The p stands for the German potenz, meaning power or
concentration, and the H for the hydrogen ion (H+).

The formula for calculating pH is:

     [\mbox{pH} = -\log_{10} \left[ \mbox{H}^+ \right]

where [H+] indicates the concentration of H+ ions (or also written [H3O+],
concentration of the equivalent hydronium ions), measured in moles per litre
(also known as molarity).

In aqueous solution at standard temperature and pressure, a pH of 7
indicates neutrality (e.g. pure water) because Water naturally disassociates
into H+ and OH- ions with equal concentrations of 1×10-7M. A lower pH
number (for example pH 3) indicates increasing strength of acidity, and a
higher pH number (for example pH 11) indicates increasing strength of
alkalinity. Most substances have a pH in the range 0 to 14, although
extremely acidic or basic substances may have pH < 0, or pH > 14.

In nonaqueous solutions or non-STP conditions, the pH of neutrality may not
be 7. Instead it is related to the disassociation constant for the specific
solvent used.

There is also pOH, in a sense the opposite of pH, which measures the
concentration of OH- ions. Since water self ionizes, and notating [OH-] as
the concentration of hydroxide ions, we have

     Kw=[H+][OH-]=10-14

where Kw is constant, the ionization constant of water.

Now, since

     log Kw=log [H+] + log [OH-]

by logarithmic identities, we then have the relationship

     14 = log [H+] + log [OH-]

and thus

     pOH = log [OH-] = 14 - log [H+]

Some common aqueous pH's

   * 3.5: orange juice (slightly acidic)
   * 7.0: pure water
   * 7.34 - 7.45: human blood (slightly alkaline)
   * 11.0: household ammonia (very alkaline)

Measuring

pH can be measured by addition of a pH indicator or using a pH meter.
Universal Indicator changes colour depending on the pH of the solution it is
added to. Electronic pH meters consist of an electrolytic cell in which an
electric current is created due to the hydrogen cations completing the circuit.

Calculation of pH for weak and strong acids

Values of pH for weak and strong acids can be approximated using certain
assumptions. It is assumed that for strong acids, the dissociation reaction
goes to completion (i.e., no unreacted acid remains in solution). Dissolving
the strong acid HCl in water can therefore be expressed:

     HCl(aq) → H+ + Cl-

This means that in a 0.01 M solution of HCl it is approximated that there is
a concentration of 0.01 M dissolved hydrogen ions. From above, the pH is: pH
= -log10 [H+(aq)]:

     pH = -log(0.01)

which equals 2.

For weak acids the dissociation reaction does not go to completion, an
equlibrium is set up between the ions and the acid. The following shows the
equilibrium reaction between methanoic acid and its ions:

     HCOOH(aq) ↔ H+(aq) + HCOO-(aq)

It is necessary to know the value of the equilibrium constant of the
reaction for each acid in order to calculate its pH. In the context of pH,
this is termed the acidity constant of the acid but is worked out in the
same way (see chemical equilibrium):

     Ka = [hydrogen ions (aq)][acid ions (aq)] / [acid (aq)]

For HCOOH, Ka = 1.6 × 10-4

Two assumptions are made in the calculation of pH for a weak acid. It is
assumed that the water the acid is dissolved in does not provide any
hydrogen ions. Water is a very weak acid and in general it supplies far
fewer than the acid dissolved in it. Consequently in the above reaction the
concentration of hydrogen ions equals the concentration of methanoate ions:

     [H-(aq)] = [HCOO-(aq)]

It is also taken that the the amount of undissociated acid at equilibrium is
equal to the amount originally added to the solution. Although this is
obviously untrue (otherwise the pH would remain 7!) this amount can be
neglected because the fraction of hydrogen ions given is again very small.

With a 0.1 M solution of methanoic acid (HCOOH), the acidity constant is
equal to:

     Ka = [H+(aq)][HCOO-(aq)] / [HCOOH(aq)]

So:

     1.6 × 10-4 = [H+][HCOO-] / 0.1

     1.6 × 10-4 × 0.1 =[H+][HCOO-]

As [H-(aq)] = [HCOO-(aq)]:

     1.6 × 10-4 × 0.1 =[H+]2

The concentration of hydrogen ions is: 4 × 10-3. The pH, therefore,
is: 2.3.