Site author Richard Steane
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pH and hydrogen ion concentration [H+]

Water - the basis for aqueous solutions, and life

Water consists of molecules built from 2 atoms of hydrogen attached to an oxygen atom: H2O.
The hydrogen is covalently bonded to the oxygen.

However a very small number of these molecules interact to exchange an electron, which causes them to break down and dissociate into hydrogen ions: H+ and hydroxide ions: OH-:
H2O ⇌ H+ + OH-
This is written with a reversible arrow as it is possible for H+ and OH- to re-form H2O.

The simplest ion - with several names

Hydrogen ions are also known as protons. This term is used in biochemistry, e.g. proton pumps.

In many (chemical) texts you will see reference to a hydronium ion, also known as hydroxonium or oxonium: H3O+
This is effectively the same as a water molecule H2O combined with a hydrogen ion H+.
It probably portrays the origin of these ions more accurately from 2 water molecules:
2 H2O ⇌ H3O++ OH-
However the concept of pH revolves around the concentration of H+ which is identical to H3O+.

Concentration time

Pure water is almost all H2O; the concentration of H+ (and OH-) is very small indeed.
In fact the ratio of H2O:H+ is approximately 5.5 x 10 8(: 1).
In other words, there are 550 million molecules of water for each hydrogen ion.

The H+ concentration of pure water is actually 10-7 M.

This is actually another way of expressing the fraction:
           1           M

10-7 is 'ten to the power of minus seven', and this (the term 'power') gives the 'p' in pH.

Concentrations are sometimes written in square brackets, so the hydrogen ion concentration is shown as [H+].


Concentrations are usually expressed in terms of molarity.

A molar solution (1.0 M) can be made by dissolving the relative molecular mass of a substance in water and diluting it to exactly 1 litre - 1 dm3.
It is also sometimes written as 1.0 mol dm-3, 1.0 mol/L or 1.0 mol L-1

Of course we cannot obtain hydrogen ions from a container in the prep room, unlike other substances served up as solutions in Biology lessons!

Water - strong stuff

Since the relative molecular mass of water is 18, 1 litre of water (1000g) contains 55.55 moles, so water is effectively a 55.55 molar solution of H2O!

pH - Let's start in neutral

pH is a numerical way of defining the hydrogen ion concentration of a liquid.
It is actually the negative logarithm (to the base 10) of the hydrogen ion concentration.

Pure water is neutral. So in order to work out the pH of pure/neutral water, we proceed as follows:
Taking the value above ( 10-7 M) we extract the power (-7) and ignore the -sign.
This leaves 7, which is the pH value for pure water.

Up and down - one unit at a time

Acids increase the concentration of hydrogen ions in a solution.
This has the effect of lowering the pH.
An acidic solution with the pH 6.0 has a hydrogen ion concentration of 10-6 M;
Click to see/ hide more:
So each reduction in pH reading of 1.0 results in a 10x increase in hydrogen ion concentration, 2.0 gives a 100x increase and 3.0 gives 1000x increase and so on - a logarithmic scale.
Alkalies increase the concentration of hydroxide ions in a solution and decrease the concentration of hydrogen ions.
This has the effect of increasing the pH.
An alkaline solution with the pH 8.0 has a hydrogen ion concentration of 10-8 M;
Click to see/ hide more:

Working out pH - starting from hydrogen ion concentration

If you are given the hydrogen ion concentration - [H+] - as power of 10, then the pH is just the number of the power (ignoring the - sign)
e.g. [H+] = 10-8.5
gives you the pH as 8.5

If you are given the hydrogen ion concentration in standard form you can calculate the pH using the log function on a scientific calculator:
[H+] = 3.2 x 10-9 M
Key in 3.2 ÷ 1000000000 [or however you input 10-9] then =, then 'log' also gives you the pH as 8.5 (or maybe 8.494850022! )

pH indicators

There are a variety of chemical substances that can be used to give information about the pH of a solution, either in liquid form or as paper.

Natural compounds

litmus_paper (23K) Litmus consists of (a mixture of) compounds obtained from lichens. It is red in acidic conditions, blue in alkaline and more or less purple in neutral. It has been used for hundreds of years as a simple test for acidity/alkaliniity.

Other plant dyes react to pH. Anthocyanins are red, purple, or blue water-soluble pigments found in (vacuoles within cells of) many plants, notably in flowers but also in leaves and fruits, and their colour varies with pH. Anthoxanthins are similar yellow or white compounds.

Synthetic compounds

There are many of these.
phth (4K) Phenolphthalein is a compound with quite a distinct colour change, often used in Biology, especially with digestion of lipids. It is also used as an end-point indicator in the titration of acids. For all practical purposes it is colourless below pH 8.5-9, above this it is a bright pink fuchsia colour, although it apparently changes to other colours at extreme ends of the pH scale.

Universal indicator - of which there are several versions - consists of a mixture of several compounds, which react to a wide range of pH values giving a rainbow-like spectrum of colours.
UIpHscale (52K)

pH meters

These consist of a glass probe connected to an electronic meter.
- advantages: - disadvantages:

Acids and alkalies - weak and strong

Acids dissociate to give H+ and another anion.
Strong acids dissociate completely:
H2SO4 → 2 H+ + SO4-
sulphuric acid --> hydrogen ions + sulphate ion

HCl → H+ + Cl-
hydrochloric acid --> hydrogen ion + chloride ion

and strong alkalies also dissociate completely:
NaOH → Na+ + OH-
sodium hydroxide --> sodium ion + hydroxide ion

whereas weak acids only dissociate partly:
ethanoic acid <--> hydrogen ion + ethanoate ion
This is described as being in chemical equilibrium.


This partial dissociation has consequences in buffering action. A pH buffer is a solution that maintains the pH at a constant value, despite additions of (small amounts of) acid or alkali which would otherwise cause a change in pH. Buffers are generally mixtures of a weak acid and its salt, or a weak alkali and a corresponding salt. The proportion of these ingredients can be varied to give a set pH. Because there is an equilibrium between the dissociated and undissociated components, hydrogen ions will enter or leave solution in response to added ions, so the pH is kept constant.

Environmental variations in pH

Water in the laboratory (distilled/deionised) does not usually have a pH of 7.0.
If it is in contact with carbon dioxide in the air (0.035%) some of it forms carbonic acid, which then dissociates:
CO2 + H2O ⇌ H2CO3 H+ + HCO3-
As a result, its pH becomes 5.7.

In theory, the same should apply to rainwater, which is the result of a natural distillation process.

But if there are any soluble gases in the air which give rise to an acidic solution, the result is acid rain. This is mostly due to sulphur dioxide (SO2) and nitrogen oxides (NOX), originating from burning of fossil fuels to generate electricity, as well as exhaust gas from motor vehicles and other industrial processes. A small amount comes from volcanoes. The result is sulphuric and nitric acid, which brings the pH of rain (and other precipitation e.g. snow, fog, or hail) down to 4.2-4.4. As well as damage to buildings, the acid may solubilise aluminium in soils leading to toxicity to plants and aquatic life. Base cations, such as calcium and magnesium, are leached by acid rain and this can have wide-ranging impacts on some ecosystems.

Once in contact with the earth's surface, water may dissolve other ions from rocks and soil. This is true of most domestic water supplies, whether taken from underground boreholes, rivers or lakes which are often used as reservoirs.
Calcium carbonate and magnesium carbonate dissolve to make hard water with a pH 7-8.5.

The pH of seawater is generally taken to be 8.1 ( range between 7.5 and 8.4) and this has been reducing gradually as carbon dioxide dissolves in sea water.

Effects of pH on living organisms

Many marine organisms are adversely affected by acidification. For example, corals and many other animal groups have a "skeleton" of calcium carbonate and become threatened as it becomes soluble.

Enzymes perform a variety of roles in living organisms. Digestive enzymes (mostly extracellular) are known to work at different pHs. Enzymes may be affected by changes in the pH of their environment. Groups on sidechains on amino acid residues may change in charge with result that substrate may not bind or be converted efficiently, or in extreme conditions the 3-D nature of the active site may be permanently changed so that the enzymes is effectively denatured.

Variation of pH within the digestive system

The pH of the various parts of the digestive system varies according to the type of chemical conversion that is in progress there and enzymes effectively turn on and off accordingly. It is worth remembering that the gut wall is covered with a mucus layer which protects it from the pH and digestive enzymes.

The mouth is an environment which is fairly close to neutral - pH 5.6 to 7.9. Salivary amylase starts working on starch here.
Sugars, especially those from sweets, may be converted into acids by the action of bacteria, and this is a significant factor in tooth decay. Chewing gum may reduce this, along with increasing the flow of saliva.

In the stomach the pH drops to 2-3. This is due to secretions from oxyntic cells (parietal cells) which produce H+ and Cl-. This acidity does not of itself dissolve food, but it assists enzymes to carry out digestion, as well as killing bacteria on food. Pepsin is secreted in an inactive form, pepsinogen, and once mixed with the acidic stomach contents it is activated by the removal of a section of the polypeptide chain covering the active site. Digestion of proteins is optimised by the acid-generated unfolding of polypeptide chains, allowing access by proteolytic enzymes. The proteolytic enzyme pepsin has an optimum pH of 2.0.
Salivary amylase becomes much less active in the stomach, but it has been shown that it remains more active in the presence of its substrate starch.
And if there is an excess of acid, then the next stage is compromised, leading not to digestion but indigestion ...

In the duodenum, the pH rises due to secretion of bile (produced in the liver, stored in the gall bladder). Bile salts emulsify fats and oils in alkaline conditions. Pancreatic juice is also quite alkaline, and it contains 3 types of enzymes with an optimum pH to match: The protease trypsin has an optimum pH of 7.8-8.7 which is normally found in the ileum. Pancreatic lipase, active in the same area, has an optimum pH of 8.0. Curiously, the optimum pH for pancreatic amylase is quoted to be 6.7 - 7.0

pH of blood and body fluids

Blood's pH is normally 7.4, and it is maintained very close to this level: +/- 0.05 pH units. Body fluids derived from it: tissue fluid, lymph have a similar pH.

However blood pH may rise due to alkalosis or fall due to acidosis.
Alkalosis causes the serum pH to be 7.45 or higher (alkalemia). Respiratory alkalosis is caused by hyperventilation, resulting in a loss of carbon dioxide, and metabolic alkalosis can be caused by loss of hydrochloric acid within the stomach content due to vomiting, dehydration, diuretic drugs or endocrine disorders.
Conversely, acidosis causes the serum pH to be 7.35 or lower (acidemia). Respiratory acidosis reflects lung problems and metabolic acidosis can be caused by problems with the kidneys.

Other related topics on this site

(also accessible from the drop-down menu above)

Water - in a class of its own



Web references

Dissociation, self-ionization of water, Kw from Online Introductory Chemistry - Dr. Walt Volland

Definition of pH, pOH, "p", sample calculations - same again

Carbonic acid From Wikipedia, the free encyclopedia

Litmus From Wikipedia, the free encyclopedia

Rising Acidity in the Ocean: The Other CO2 Problem from Scientific American

Seawater From Wikipedia, the free encyclopedia

What is Acid Rain? from US Environmental Protection Agency

Effect of gum chewing on the pH of dental plaque.

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