Written by Andy Connelly, published 26th March 2017. Last updated 15th July 2017.
Aqueous hydrofluoric acid is not a strong acid ; however, it is very dangerous. The first question you should always ask is. Do you need to use HF? If there is an alternative, take it [1, 2].
Although hydrofluoric acid (HF) is a weak acid it does attack silicates very effectively, and most oxides-although not, for example, spinels and zircons . For this reason, its primary laboratory use is to attack and dissolve silicate minerals.
It is thought of as a weak acid because, although HF dissociates strongly in water, it does not form free H3O+ (H+) ions but forms a strongly bound ion pair, H3O+.F- . The lack of free H3O+ ions makes it a weak acid as shown in the acid dissociation constant (pKa) values (see Table 1):
We tend to use HF directly in its aqueous form; however, there are alternatives. Many fluoride-containing solid chemicals (e.g. ammonium bifluoride, sodium fluoride, etc.) may react with acid or water to produce HF. In this way HF, can be produced in situ to help reduce handling of the aqueous acid. However, once the reaction has started the same HF precautions are required so the advantages are limited.
DISCLAIMER: I am not an expert in health and safety. The content of this blog is what I have discovered through my efforts to understand the subject. I have done my best to make the information here in as accurate as possible. If you spot any errors or admissions, or have any comments, please let me know. HANDLING OF HYDROFLUORIC ACID OR ANY OTHER CHEMICAL MENTIONED IN THIS POST SHOULD BE CARRIED OUT UNDER LOCAL H&S POLICIES AND WITH APPROPRIATE RISK AND COSHH ASSESSMENTS. ALL STATEMENTS AND INFORMATION GIVEN IN THIS BLOG POST ARE PRESENTED WITHOUT GUARANTEE, WARRANTY, OR RESPONSIBILITY OF ANY KIND. THE READER SHOULD NOT ASSUME THAT ALL SAFETY MEASURES ARE INDICATED HEREIN.
Aqueous hydrofluoric acid is a contact-poison with the potential for deep, initially painless, burns and ensuing tissue death. Once burns become apparent they produce a characteristic excruciating pain.
At concentrations of 30 parts per million (around 0.003%) HF is immediately dangerous to life . Not only is the fluoride ion itself toxic but it interferes with the body’s calcium metabolism. These effects can cause systemic toxicity, and eventual cardiac arrest. It is often said that there is a serious potential for death when skin burns exceed 160cm2  (see Figure 1). However, this is fairly unhelpful. what volume of acid does this refer to? What concentration? Smaller areas than this can also result in serious medical problems (e.g. serious cardiovascular problems) and potentially death if not treated. For HF, all skin contact should be treated has potentially hazardous to life.
It is particularly dangerous in contact with eye and in all instances of inhalation or ingestion. Even short period contact with low concentrations can cause serious medical problems in these areas . What is particularly shocking is how quickly the HF can cause death.
“…a petroleum worker whose face and neck were in contact with 10 per cent HF for only 30-35s. Despite intravenous and subcutaneous calcium gluconate he died 2 h later. Post-mortem findings showed thick mucus with blood in the bronchi, edematous lungs with hemorrhagic patches, and sub-serous hemorrhages in the pleurae.”
“The first recorded fatality from fluoride poisoning was in 1873 when King reported the case of a 35-year-old man who died within 35min after drinking about 15ml of hydrofluoric acid.”
For information on treatment and control of HF please see your local HF policy and consult you local H&S officer.
Storage of HF
HF should not be stored in glass bottles as it attacks silicates. While other suitable materials are available (see ) plastics are the most appropriate for most laboratories. Polyethylene, polypropylene, and poly(vinyl chloride) show good resistance to dilute solutions containing up to 50% of HF at ambient temperature. Polytetrafluoroethylene (PTFE) shows good resistance. However, PTFE is not completely impermeable to HF. 
Hydrofluoric acid etches glass, due to the strong bond formed between fluoride anions and the silicon molecules in glass (see below). HF will also react with glazes, enamels, pottery, concrete, rubber, leather, many metals (especially cast iron) and many organic compounds. Even when it has been “neutralised” by boric acid the result can be very corrosive to glass (see below).
Uses of HF
There are two main uses of HF in physical science laboratories:
HF is an ideal acid for etching of glassware due to reactive nature with silicates but relatively low acidity (see Table 1). The fact it is a weak acid allows a masks to be used on some areas of glass so they can remain protected from the HF and so are not etched.
The reactivity of HF with silicates follows two main reactions :
The release of SiF4 gas in the first pathway drives the reaction forward. However, SiF4(g) is a toxic and corrosive gas which can form HF on contact with water and so raises its own H&S issues.
The most common use of HF in physical chemistry and analytical laboratories is probably for the complete, or partial, dissolution of silicate materials (e.g. rocks, sediments, etc.). HF is one of several chemicals commonly used in such extractions. It is used for silicate rock digest due, in part, to the solubility of hexafluorosilicate ion (SiF62-) in acidic solutions . It also produces solutions that are relatively easy to analyse by ICP-MS or AAS [14,15]
The silicate attacking role of HF is particularly important for palaeontology where organic materials need to be separated from silicate rocks [7,21]. However, for most digests, hydrofluoric acid is not used alone. This is because HF is a non-oxidising acid and because some salts do not have high solubility in HF; including alakaline (earth), actinide and lanthanide elements e.g. calcium and potassium. This is partly due to the formation of low solubility fluorides (e.g. CaF).
Another issue with using HF on its own is that the resultant solution is likely to contain some elements (e.g. Fe) in mixed oxidation states making analysis of the resulting solutions more difficult . For these reasons: nitric, perchloric, and other acids are normally used in conjunction with HF .
Other acid used in digestions
Some other common acids used in digests are shown below and in Table 1 [7,9]. These are often used in conjunction with HF to create the ideal mixture depending on the material you wish to digest. These chemicals all have hazards associated with them and full assessment of the hazards involved (COSHH or similar) must be carried out before their use (see Table 2).
Perchloric acid is a powerful oxidising agent when hot. It has the potential for explosive reaction with organic matter. To avoid this nitric acid is often added to the digest if organic matter is present. The nitric acid reacts with more reactive organic components at lower temperatures than the perchloric so ensuring the reaction is less violent .
Perchloric acid salts are generally very soluble in water (except K, Cs and Rb). Perchloric acid can also become explosive when anhydrous which is a serious danger.
Nitric acid has poor oxidising strength below 2 M but is powerful in its concentrated form. Its oxidising potential is enhanced with increased temperature and pressure. Metal salts of nitric acid are generally very soluble in water except gold, platinum group metals, Al, Ti, Cr, and Zr. Silicates are not as soluble in nitric acid-hence need for HF. Fe, Cu and Ca can also be difficult and hydrochloric acid is sometimes added to help their dissolution.
Nitric acid is commonly used for the digestion of organic matrices :
Hydrochloric acids is not normally used alone for digesting silicates. It is a strong non-oxidising acid that exhibits reducing properties during dissolution. These properties allow hydrochloric acid to be used to remove carbonates from soils/sediments without major damage to organic materials.
However, HCl’s strong complexing nature allows for complete dissolution of numerous metals: platinum group metals, Au, Fe, Cd & Sn. It is also useful to help dissolution of Al, In, and Sb. In general, the addition of conc. HCl is appropriate for the stabilization of Ag, Ba, and Sb, and high concentrations of Fe and Al in solution. The amount of HCl required will vary depending on the matrix and the concentration of the analytes.
There are also insoluble metal chlorides which can hamper dissolution. These include: AgCl, HgCl, and TiCl. Also, PdCl2 which is partial solubility. There are also some chlorides that are volatile including Ge, Hg, Sb,and Sn.
When hot and concentrated sulphuric acid is oxidising. Its high boiling point (see Table 1) can be used to help accelerate dissolution of metallic materials and minerals. However, it is not generally used for rock/soil digests due to low solubility of many inorganic sulphates. It is used for decomposition by dehydration and oxidation of organic compounds. At room temperature sulphuric acid is a relatively non-oxidising acid although it forms a strong oxidising agent with nitric acid.
This is used to help remove organics from rocks/soil. Its oxidizing properties increase as acidity of the solution increases.
Used for complexing geological elements in more specific circumstances.
Addition of water helps to reduce pressure build up in sealed digestion systems. It can also provide greater solubility of many metals.
As established above, digestions are rarely carried out with one acid alone. They are usually done with mixtures. Below are a few common mixtures used in digests [8,9]:
- Aqua regia (3 parts hydrochloric acid to 1 part nitric acid) – oxidising agent, used to dissolve metals, especially platinum and gold
- Reverse aqua regia (3 parts nitric acid to 1 part hydrochloric acid) – used for most organic materials
- HNO3:HCl:HF Used, in various ratios, for the dissolution of alloys, ores, silicates, ash
- Caro’s acid (addition of hydrogen peroxide to sulfuric acid) – forms persulfuric acid used to oxidize organic samples.
- Sulphuric/phosphoric: commonly used for dissolution of alumina and materials containing alumina such as ceramics, catalysts, slags, and ores. Provides high temperature at low pressure.
- Nitric/sulphuric: Combination used to enhance decomposition of organic samples
The hazards inherent in handling HF mean that a method of neutralisation is vital in case of spillage. Also, the solutions that result from a HF extraction will often need to be neutralised, or in some way made safe, to allow an analyst to handle the samples. See Table 3 for some of the chemicals that are used for neutralisation.
The neutralisation process you choose will depend on what you are planning to do with the solution after. The following are two methods which are commonly used. NONE OF THEM GUARANTEE COMPLETE NEUTRALISATION. It is up to you to make an assessment as to whether you are happy that the samples are “safe”.
Boric acid is used to “neutralize” excess HF by complexing the fluorine, forming tetrafluoroboric acid in solution . Boric acid is not normally used to neutralise spills as the complexing reaction is slow.
An example method for neutralisation by boric acid is shown in reference . It is often said that, 1 gram of boric acid should be added for every 1 mL of 49% HF [unknown source]. As boric acid is normally used as a 4% or 4.5% solution that would mean around 25ml of 4% boric acid solution per 1ml HF would be required. However, the stoichiometric amount required is much less than this (see Table 3). Whatever volume you use the above does not guarantee neutralisation and appropriate assessments must be carried out.
Neutralisation using boric acid has two added advantages:
- It can be used to prevent volatilization of silicon (as SiF4) so improving Si recovery
- It prevents formation of, or helps resolubilise, insoluble metal fluoride complexes (e.g. Ca, Al, and Mg).
The rate-determining step in the neutralisation process is the creation of fluoroboric acid (HBF4) which can be a slow process and generally requires a heating step (see ).
Despite these positives the neutralisation of HF by boric acid does have a down side. There is a general misunderstanding that the addition of boric acid will eliminate HF attack of glassware. Unfortunately, fluoroboric acid will attack glass and the attack of silicates, in general, is not greatly altered.
Alkaline materials (bases)
Alkaline materials are the ideal for HF neutralisation as the reaction is fast and the products are much less harmful than those for boric acid neutralisation. However, if you are trying to analyse your sample for sodium then adding lots of sodium carbonate may not be ideal! Also, for many of these chemicals the purchasing of pure versions can be expensive and so there is a large possibility of contamination with an element of interest. For these reasons, alkaline materials tend to be used for neutralisation of spillages only.
If they are used then it should be in excess as there is potential for a solution of alkaline fluoride to be returned to a solution of HF if pH is reduced through addition of acid (see below). For example, there is potential for sodium fluoride containing cleaning materials to react with stomach acid to produce HF in the body where they have been accidentally ingested .
There are disadvantages with using some alkaline materials (see Table 3):
- Sodium or Potassium Carbonate: The reaction of these with HF generate sodium or potassium hydrogen bifluoride (NaHF2 or KHF2) as intermediates, which release gaseous HF when exposed to heat .
- Potassium or Sodium Hydroxide: The neutralization of HF with these is more exothermic than with sodium or potassium carbonate and also generates potassium or sodium hydrogen bifluoride (NaHF2 or KHF2) as intermediates, which release gaseous HF when exposed to heat .
- Low solubility of calcium carbonate or calcium hydroxide mean they are not ideal to be used in a solution for neutralisation (see Table 3).
The high vapour pressure of hydrogen fluoride (see Figure 2) and the rapid reduction in boiling point with increased concentration (see Figure 3) would suggest that heating HF for a long period would significantly reduce concentration of HF in solution. As such, if the solution containing HF is heated to dryness (all liquid evaporated) then the concentration of HF has, in theory, been minimised. However, fluoride salts (e.g. NaF) dissolve in water and/or low pH acid solutions to form HF (depending on the salt). This means that soluble fluoride salts in solution with acids (that are stronger than HF) should be treated as solutions of HF .
The result of this is that a boric acid or alkaline step is recommended after heating to dryness to neutralise/complex any remaining F ions in the solution (e.g. from NaF, CaF2, etc.)
Figure 2: Vapour pressure of hydrogen fluoride as a function of temperature .
It is assumed that no level of dilution in the laboratory can ever make HF safe. However, the Short Term Exposure Limit (STEL) for HF  is 3ppm which is equivalent to approximately 0.0001M. So, in theory, dilution to this level will drop the concentration below the STEL. However, this means that for 1ml of 40% (22.6M) HF you would need to dilute with 226L of water to make it safe!
While STEL is an inhalation measure based on hydrogen fluoride (as F)  this does give an indication of the difficulties in trying to dilute HF to make it safe. Skin contact with HF, even dilute solutions (0.1%) can cause painful second and third degree burns that heal very slowly . In general, DILUTION IS NOT RECOMMENDED AS A METHOD OF NEUTRALISATION – EVEN DILUTE HF CAN BE DANGEROUS. Appropriate H&S assessment must be carried out for any such operations.
You must consult the Faculties HF policy on spills. However, silicon-based absorbent materials should not be used. These materials (e.g. vermiculite) are common in most solvent spill kits but they react with HF to generate silicon tetrafluoride, which is a toxic and corrosive gas. 
Hydrofluoric acid is a dangerous chemical and local policies must be followed when handling it. However, if you need to use it then it can be handled effectively and safely with appropriate precautions.
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