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PH
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PH

The title of this article given above is incorrect, due to technical limitations. The correct title is pH.

Alternate uses: see Ph

pH

Acids and Bases:
Acid-base reaction theories
pH
Self-ionization of water
Buffers
Systematic_naming
Redox reactions
Electrochemistry
Strong acids
Weak acids
Weak bases
Strong bases
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 layman's terms , the "pH" value is an approximate number between 0 and 14 that indicates whether a solution is acidic (pH < 7), basic (pH > 7) or neither (pH = 7).

Table of contents
1 Definition
2 Some common pH values
3 Measuring
4 pOH
5 Calculation of pH for weak and strong acids
6 Neutralisation
7 See also

Definition

The formula for calculating pH is:

[H+] indicates the concentration of H+ ions (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 dissociates into H+ and OH- ions with equal concentrations of 1×10-7 M. 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.

Some common pH values

pH
Battery acid 1.0
Stomach 2.0
Lemon juice 2.4
Cola 2.5
Vinegar 2.9
Orange or apple juice 3.5
Beer 4.5
Coffee 5.0
Tea 5.5
Acid rain < 5.6
Milk 6.5
Pure water 7.0
Blood 7.34 - 7.45
Sea water 8.0
Hand soap 9.0 - 10.0
Household ammonia 11.5
Bleach 12.5
Household lye 13.5

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.

pOH

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 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+] = 14 - pH

Calculation of pH for weak and strong acids

Values of pH for weak and strong acids can be approximated using certain assumptions.

Under the Brønsted-Lowry theory, stronger or weaker acids are a relative concept. But here we define a strong acid as a species which is a much stronger acid than the hydronium (H3O+) ion. In that case the dissociation reaction (strictly HX+H2O↔H3O++X- but simplified as HX↔H++X-) 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+]:
pH = -log(0.01)
which equals 2.

For weak acids, the dissociation reaction does not go to completion. An equilibrium is reached between the hydrogen ions and the conjugate base. The following shows the equilibrium reaction between methanoic acid and its ions:

HCOOH(aq) ↔ H+ + COOH-

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][acid ions] / [acid]

For HCOOH, Ka = 1.6 × 10-4.

When calculating the pH of a weak acid, it is usually assumed that the water does not provide any hydrogen ions. This simplifies the calculation, and the concentration provided by water, 1×10-7 M, is usually insignificant.

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

Ka = [H+][COOH-] / [HCOOH]

Given that an unknown amount of the acid has dissociated, [HCOOH] will be reduced by this amount, while [H+] and [COOH-] will each be increased by this amount. Therefore, [HCOOH] may be replaced by 0.1-x, and [H+] and [COOH-] may each be replaced by x, giving us the following equation:

Solving this yields 3.9×10-3, which is the concentration of hydrogen ions after dissociation, so the pH is 2.4.

Neutralisation

Neutralisation can be summed up by the formula:

H+ + OH- = H2O
(acid + alkali = water)

See also