pH is an expression (in a dimensionless number) of the acidity of a liquid solution. The pH of a neutral solution lies (at room temperature) around 7.0. Acid solutions have a pH lower than 7, basic solutions have a pH higher than 7.
The concept of pH was introduced in 1909 by Søren Sørensen. The p stands for the German Potenz, meaning power, and the H stands for the hydrogen (H+) (Latin: Pondus hydrogenii and/or Potentia hydrogenii).
The pH is equal to the negative logarithm (with base number 10) of the concentration of hydrogen ions (H+). The unit of concentration is mol/litre.
Formally, it is better to describe the pH as the negative logarithm of the hydrogen ion activity. By doing so the factor (f) gives the activity of the hydrogen ions as a value between 0 and 1. In normal conditions the factor is approximately equal to 1.
In water or in a watery solution, a certain amount of the water molecules are present in the form of ions. Two H2O molecules are split into a positive H3O+ ion and a negative OH- ion. The solubility product of both ions in water is 10-14 (at 22°C), which is to say that in every watery solution the product of the concentration of OH- and the concentration of H+ is always the same: 10-14. So if we presume pure water, then sufficient water is split into ions that the concentration of H+ as well as the concentration of OH- is the same at 1:10,000,000 = 10-7. The pH of this according to the formula is: ? log10 ? 7 = ? ( ? 7)log10 = 7.
All solutions with a pH of 7 are called neutral solutions. Such a solution is neither acidic nor alkaline (basic).
The pH scale is a logarithmic scale, which for watery solutions in practice runs from 0 to 14. Lower than 7 means that the solution is acidic (the lower the more acidic). Above 7 means that the solution is basic. Values below 0 and above 14 are possible, but only in extreme situations and such solutions are generally very dangerous: concentrated acids and concentrated alkalis.
The logarithmic character of the scale means that even within a scale of 0 - 14 very extreme values can be displayed. In a solution with a pH of 8 there are more than 100 times as many OH- ions as H+ ions, and in gastric acid with a pH of 1 there are 1,000,000,000,000 times as many H+ ions as OH- ions.
In addition to the pH scale there is also the pOH scale, which is precisely the opposite of the pH scale. Where the pH scale measures the activity of H3O+ ions, the pOH scale shows the activity of OH- ions. This comparison only applies at a temperature of 293.15 K. As the temperature increases, so the levels of both pH and pOH decline.
The pH of a solution can be measured in a variety of ways.
There are dyes (pH indicators) that change colour if H+ ions are being taken up or released. This colouration takes place at a certain pH. If you use a variety of colouring agents (in solution or impregnated in paper for ‘pH strips’) the colour change can be used to give a rough estimate of the pH.
The Litmus colouration is well known, but in the kitchen there are also natural pH indicators, such as when red cabbage (the pan turns blue in the basic washing up water; recipes for red cabbage often include some acidic apple, lemon and/or vinegar to prevent it discolouring while cooking). Also well known is the way black tea turns lighter when a slice of lemon is added.
There are electrochemical reactions in which H+ ions are involved and in such reactions the voltage varies as a function of the pH. One can measure the pH with a pH meter by measuring the electrochemical reaction under controlled conditions. This is done by a process of titration, using a strong base (often caustic soda). Titration involves dripping the base into the solution under investigation until this solution contains no more ampholytes. In order to make this visible a pH indicator is used with a transition point close to 7. This method is more accurate than is actually needed in practice if you only want a measurement of the pH.
pH 14: caustic soda solution (1 mol/l)
pH 13: caustic soda solution (0.1 mol/l)
pH 11,5: domestic ammonia
pH 10,5: soapsuds
pH 8,5: seawater, intestinal fluid
pH 7,4: human blood
pH 7: pure water (neutral)
pH 6,7: milk
pH 6: rain water
pH 5: lightly acid rain
pH 4,5: acid rain, tomato juice
pH 3: domestic vinegar
pH 2,5: cola
pH 2: stomach acid, lemon juice
pH 1: sulphuric acid (battery acid)
pH 0: hydrochloric acid (1 mol/l)
A solution that can hold its H3O+ concentration constant, even if H+ ions are produced or used by a chemical reaction, is called a buffer. A buffer can be made by adding a weak acid and a correspondingly weak base to a watery solution. If H+ ions are made in such a buffered solution, they can absorbed by the weak base that is present, with the formation of a weak acid, without changing the pH. The reverse is also true; if H+ ions are taken out of the solution, new H+ ions are freed up by the weak acid in the formation of a weak base.
Each combination of a weak acid and a correspondingly weak base has its own ideal pH at which the buffer functions best (it can absorb or release the most with the smallest deviation in pH). In addition, the so-called buffer capacity depends on the concentration of the buffering substances in the solution.
Examples of substances we can use as buffering agents include acetic acid (vinegar) and hydrogen sulphate (pH about 4.77), or H2PO4 and HPO42.
What is the pH of a solution with a concentration of 6.5 · 10-4 mol/l H+-ions?
The concentration is a figure between 10-4 and 10-3. This will let you estimate the pH already: it will be between 3 and 4.
The term ‘pH neutral’ stated on cosmetic products means something other than a pH of 7. It means that the product has a pH that matches the natural pH of the skin. The natural pH of the skin is about 5.5.