I often get comments about the chemistry of community water fluoridation which make clear the need for a simple explanation of some of the chemical concepts involved. Here is an article I wrote some time ago on this but, for the life of me, I cannot remember where I put it on-line. So, I might be repeating myself – but, at least, it makes the article available. (t can also be downloaded as a pdf – just click on the title
Some claims made by critics of community water fluoridation (CWF) are chemically wrong. However, they may seem convincing to people without a chemical background.
Here I discuss some of the chemistry involved in the fluoridation debate and show how these claims are wrong.
What happens when a solid dissolves in water?
In a solid like ordinary salt (NaCl) the atoms exist as positively (Na+ cations) and negatively (Cl- anions) ions in a rigid lattice structure. This structure is generally stable as the ions are held together by electrostatic forces. But the structure can often be disrupted by water. The water molecular (H2O) is polar – it has a negatively charged end (the O atom) and a positively charged end (the H atoms). These ends are attracted to the oppositely charged ions, surrounding them and bringing them into solution.
Once in solution, the ions are free from the rigid lattice and move about by themselves. The cations and anions are randomly mixed up through the volume of solution.
The ions are also hydrated. Depending on the chemical nature of the ion and its charge there may be different numbers of water molecules in the primary hydration shell right next to the ion. But other water molecules are also weakly associated outside the primary shell so we can think of anions and cations in solution as being sheathed by jackets of water molecules.
Can calcium fluoride exist in solution?
Some people suggest that natural forms of calcium fluoride are not toxic because the calcium modifies the fluoride. But technically there is no such thing as calcium fluoride in solution.
In nature fluoride is usually present as solid fluorite (calcium fluoride) or fluorapatite (a calcium phosphate containing fluoride and other ions). But when calcium fluoride dissolves the ions separate and the resulting solution is a random mixture of hydrated fluoride anions and hydrated calcium cations.
Fluorite (“natural” calcium fluoride) dissolves to form hydrated calcium cations and fluoride anions.
So our “natural” water containing “natural” fluoride actually does not contain calcium fluoride. Calcium fluoride does not exist as a separate species in solution. It contains a random mixture of hydrated fluoride anions and hydrated calcium cations.
We can describe this with the chemical formula:
CaF2 → Ca2+(aq) + 2F–(aq)
Where the (aq) notation identifies the ion as being hydrated in the solution.
What ions are in your drinking water?
In the real world, our “natural” water source contains more than this, though. It contains other ions which have dissolved from minerals or from other sources like rain and runoff.
In reality, our “natural” water should be considered as a solution of a range of randomly distributed anions and cations. Because of the nature of dissolved ions and the multiple ions present we cannot describe our “natural” water as containing “calcium fluoride,” “sodium chloride” or any other common chemical. These names are really only applicable to the ionic solids. Rather the water is a solution of hydrated Ca2+(aq), Na+(aq), F–(aq), Cl–(aq), etc. We have to characterise the water by the amounts of each ion present in solution.
The drinking water you get after treatment may contain less of some of the natural ions, or more if extra is added during (eg. F– is naturally in the water source but sometimes supplementary fluoride is added to provide concentration optimum for dental health).
Your drinking water contains a random mixture of hydrated anions and cations
You may think I have missed some obvious ions. For example – H+(aq) and OH–(aq). These are usually understood as present (at extremely low concentrations) and easily derived from the H2O molecule anyway.
H2O (aq) ↔ H+(aq) + OH–(aq)
In practice, water treatment plants adjust the pH (degree of acidity or alkalinity) of your water to very near neutral where the concentration of H+(aq) and OH–(aq) are approximately the same and extremely low. They may do this by adding lime (containing Ca2+), ammonia (containing (NH4+) or other chemicals.
What about Al3+(aq)?- after all, chemicals like aluminium sulphate are added to remove colloidal material? However, this procedure works because in dilute solution Al3+(aq) hydrolyses (reacts with water) to form solid Al(OH)3 – so removing Al3+(aq) from solution.
These are sometimes added during water treatment. Their purpose is to increase the fluoride (F–) concentration to levels which are optimum of dental health. The chemicals used are generally fluorosilicic acid, sodium fluorosilicate or sodium fluoride.
Some critics of fluoridation argue these chemicals are toxic and calcium fluoride, a “natural” form of fluoride, is safe. They have even argued that community water fluoridation would be OK if CaF2 was used. But this argument is faulty for a number of reasons.
- The lower toxicity of CaF2 is a result of its lower solubility. This is why some studies show the toxicity of high concentrations of fluoride can be reduced by addition of calcium salts.
- Despite its low solubility CaF2 is sufficiently soluble to maintain a fluoride concentration of about 8 ppm (mg/L) – still far higher than the optimum concentrations aimed for in CWF (0.7 ppm).
- The low solubility of CaF2 makes it impractical as a fluoridating chemical as if added as a solid uniform equilibrium concentrations would be difficult to achieve. If added as a liquid we would need a container almost as large as the water reservoir itself to store the near saturated CaF2
- “Natural” CaF2 would be too impure for use in water treatment. Expensive processing (involving conversion to hydrofluoric acid and precipitation of CaF2) would be required to reduce the impurities.
Sometimes critics argue that “natural” fluoride in water is in the form of CaF2 which makes it safe because of the presence of Ca. But remember that CaF2 does not exist in solution which contains a random mixture of cations and anions. The hydrated Ca2+ ion is present in water naturally because it is derived from a range of sources besides fluoride minerals. It is also often added to water during treatment. So your drinking water already contains calcium, and usually at higher concentrations than if all the fluoride had been derived from “natural” CaF2
What about fluorosilicates?
Some critics of CWF claim that fluoride is not the problem. That because the most commonly used fluoridating chemicals are fluorosilicic acid (H2SiF6) and sodium fluorosilicate (Na2SiF6) the problem is the fluorosilicate species. They will even claim that we are drinking fluorosilicic acid and claim that there has been no testing of the safety of this chemical in drinking water.
But this claim is wrong. In fact, fluorosilicates react with water when diluted. They decompose to form silica and the hydrated fluoride anion. Consequently, safety studies made with sodium fluoride are completely relevant to these fluoridating chemicals when diluted.
A small amount of silica is normally present in drinking water. There is a tendency for this to polymerise and end up as solid SiO2.
Because of the extreme dilution of the fluorosilicate the liberated H+(aq) does not have a measurable effect on the pH mainly because of the equilibrium:
H2O (aq) ↔ H+(aq) + OH–(aq)
Anyway, the pH of the water is adjusted during treatment to neutral values (by the addition of acids, soda ash or lime) to prevent acid attack on pipes.
Chemicals in drinking water are extremely dilute
Critics will often wave pictures of bags of chemical being added to drinking water. Often they will illustrate their claimed danger of fluoridating chemicals by referring to safety data sheets. But these data sheets provide information on the storage, handling and disposal of the concentrated chemicals and have no relevance to the extremely dilute nature of the final drinking water.
The recommended optimum concentration of fluoride in drinking water is 0.7 ppm. Humans have difficulty imagining such extreme dilutions but the following figure provides some idea in day-to-day concepts.