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  • An Introduction to pH


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    The concept of pH was first introduced by the Danish chemist Søren Peder Lauritz Sørensen at the Carlsberg Laboratory in 1909. It’s impact has been nothing short of spectacular. Aqueous solutions can be quantitatively defined by their hydrogen ion concentration using a simple mathematical formula. A large amount of analysis begins with accurate pH testing. Laboratory professionals globally work with chemicals and reagents that can be measured to provide a pH value using the Sorensen pH scale. The pH scale is measured from 0 to 14. The term pH is derived from "p," the mathematical symbol of the negative logarithm, and "H", the chemical symbol of Hydrogen. The formal definition of pH is the negative logarithm of the Hydrogen ion activity.

    pH = -log[H+]

    The pH of an aqueous solution is directly related to the amount of hydrogen ion [H+] and the hydroxyl ion [OH-] concentrations within that solution. An atom or molecule that has either gained or lost an electron(s) is called a charged particle - or an ion. Electrical conduction through aqueous solutions is possible due to the presence of these ions. The number of electrons lost or gained determines the charge on the compound in the solution. Different compounds have the ability to lose or gain varying numbers of electrons meaning solutions of differing charge levels can be formed. Generally these compounds will dissociate (or ionize) in solution to form hydrogen [H+] or hydroxyl [OH-] ions. The ease at which these molecules dissociate determines the strength of acidity or basicity when in aqueous solution. Hydrochloric acid (HCl) and sodium hydroxide (NaOH) are examples of compounds that dissociate with ease forming the charged ions shown in the equations below:

    HCl + H2O H3O+ + Cl-

    NaOH Na+ + OH-

    In aqueous solutions, [H+]ions combine with water molecules to form the hydronium ion [H3O+]. (The terms "hydronium ion" and "hydrogen ion" can be used interchangeably in pH measurement applications.) The measurement of pH involves the measurement of the [H+]/[H3O+] ion concentrations within the aqueous solution.

    The compounds that do not dissociate very easily form weak acids or bases. This means that only a very small percentage of [H+] or [OH-] ions are present in solution. Acetic acid is a good example of this, forming less than one H+ ion for every one hundred acetic acid molecules:

    H2O + CH3COOH H3O+ + CH3COO-

    Even pure water will dissociate very weakly. At 25°C, as little as 1 x 10-7 [H+] and 1 x 10-7 [OH-] are formed for every one water molecule:

    2H2O H3O+ + OH-

    When an acid is added to water the concentration of [H+] increases whereas the concentration of [OH-] decreases, which leads to an overall decrease in pH. Conversely when a base is added to water, it has the opposite effect. The concentration of [OH-] increases, reducing the concentration of [H+], leading to an overall increase in pH.

    Pure water at room temperature (25 ℃) is considered to be ‘neutral’ with a pH of 7. This implies that the [H+] and [OH-] ion concentrations are equal. An aqueous solution with a pH value above 7 is considered alkaline meaning the [OH-] ion concentration is higher than [H+]. Conversely, an aqueous solution with a pH value below 7 is considered acidic, stemming from the [H+] ion concentration being higher than [OH-]. Solutions measuring pH 0 and 14 are considered the strongest acidic and alkaline solutions respectively.

    For a given set of conditions the relationship between [H+] and [OH-] concentrations is constant. Therefore if the concentration of either is known, then the other can be determined. pH can thus be viewed as a measurement of both acidity and alkalinity, even though by definition it is a selective measurement of hydrogen ion activity.

    pH measurement is a logarithmic function which means that a change of one pH unit represents a ten-fold change in hydrogen ion concentration. This implies that as the pH decreases, the solution becomes 10 times more acidic. For example, pH 4 is 10 times more acidic than pH 5 and 100 times (10 x 10) more acidic than pH 6. Conversely for pH values higher than 7, each is 10 times more alkaline than the next lower whole value. For example, pH 10 is 10 times more alkaline than pH 9 and 100 times (10 x 10) more alkaline than pH 8.

    It is crucial to take the determination of pH seriously as there is are a wide variety of applications that rely on accurate pH measurement. pH control is important when purifying drinking water and processing sewage to name but a few. Plants require the soil to be roughly neutral in order to grow properly, and animals can be adversely affected if their blood pH level is not within the correct limits. From keeping the ecosystem safe to reducing the effects of pollution to transferring wastewater safely, these processes would be impossible without accurate pH testing.