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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. Sørensen in 1909. The p stands for the German potenz, meaning power or concentration, and the H for the hydrogen ion (H+). 

In aqueous solution at standard temperature and pressure, a pH of 7 indicates neutrality (e.g. pure water) because water naturally dissociates 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 dissociation 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 values 

3.5: orange juice (slightly acidic) 
5.6: unpolluted rain water (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 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. 

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