Is HF a Strong Acid?


Whenever we study chemical compounds, one of the first classifications we come across is acids and bases. There are many types of criteria available to classify a given substance as acidic, basic, or neutral. Today we will discuss the nature of HF.

So is HF a strong acid? HF, or hydrofluoric acid, is not a strong acid. It is weakly acidic because it does not completely dissociate in water due to the extremely strong H-F bond. Fluorine has a strong affinity for hydrogen atoms that take a large amount of energy to break the H-F bond.

Let us look at this in details.

There are many definitions for an acid, however, let us consider the simplest one. According to Arrhenius, an acid is a substance that increases the concentration of H+ ions in a solution.

This means that an acid furnishes H+ ions in an aqueous solution by dissociating into H+ions and the corresponding conjugate base.

For example, in the case of HF: HF + H2O ⇆ H3O+ + F-, here HF gets ionized into H+ and F ions, and the H+ ions form a coordinate bond with another water molecule to yield H3O+ions in solution. The higher the concentration of H3O+ ions, the stronger the acid.

So, naturally, those substances in which the bond between the conjugate base and H is easier to break will be the stronger acid.


Types of Acids

Typically, there are 3 main types of definitions for acids and bases.

Arrhenius Theory

The first definition, as we saw above, is the Arrhenius theory of acids and bases.

It states that a substance that increases the concentration of H+ions in solution is an acid, and a substance that increases the concentration of OH ions in water is a base.

The greater the increase in concentration, the stronger the acid or the base.

Going by this example, we can easily see that HCl, H2SO4, HNO3 can be classified as acids and substances like NaOH, KOH, Ca(OH)2 can be classified as bases.


Brønsted-Lowry concept of acids and bases

This theory identifies acids and bases based on the transfer of protons between these species.

According to this theory, a Brønsted acid is a substance that donates protons, or H+ions, and a base is a substance that accepts these donated protons.

This means that for a substance to be a Brønsted acid, it must contain hydrogen atoms that can dissociate as H+ions. Further, after the H+ion comes into the reaction mixture, the base must contain a lone pair of electrons to donate to the H+ion to form a bond with it.

Also, this indicates that a Brønsted acid or a Brønsted base always exists in a conjugate acid-base pair.
For example, if HCl is a Brønsted acid, then Cl- will be its conjugate base. Let us look at a typical Brønsted acid-base reaction:

NH3 (g)   +   HCl(g)    ——>       NH4Cl(s)

In the above reaction, HCl donates its proton to the ammonia molecule, to form ammonium chloride.

Therefore, here HCl is the Brønsted acid and ammonia is the Brønsted base since it is accepting the donated proton by donating its lone pair present on nitrogen to form an N-H bond.

Interestingly, according to Arrhenius’s theory, the above reaction would not be an acid-base reaction because the reaction mixture does not contain H+ or OH- ions.


Lewis Acid-Base theory

Since the Brønsted concept focuses only on the proton-donating or accepting capacity of a substance, Lewis proposed an alternative definition for acids and bases.

This theory helps us to classify compounds as acids or bases even when they don’t have protons to donate or accept, or are not available as aqueous solutions, like solids and gases.

According to this theory, a Lewis acid is a substance that accepts an electron pair and a Lewis base is a substance that donates an electron pair. Therefore, Lewis acids are also called “electrophilic”, or “electron-loving” and Lewis bases are known as “nucleophilic” or “nucleus-loving”.

Since a lone pair of electrons is being donated and accepted, it is evident that a Lewis acid must have a vacant molecular orbital to accept electrons.

This empty orbital is called the LUMO, or the lowest unoccupied molecular orbital. Similarly, a Lewis base must contain a lone pair of electrons available to donate. This lone pair of electrons are present in the HOMO, or the highest occupied molecular orbital of the Lewis Base.

Examples of Lewis acids include metal cations such as Cu2+, Fe2+, Fe3+, compounds with incomplete octets like BF3 and AlF3, etc.

Examples of Lewis bases include OH−, CN−, CH3COO−, :NH3, H2O: , CO: and other species which contain lone pairs of electrons.


Why is HF a Weak Acid?

HF Acidity

Let us now take a deeper look at HF. HF or hydrofluoric acid is a weak acid. It is an acid because it ionizes in solution to furnish H+ions as follows:

HF   +   H2O   ⇆   H3O+   +    F

However, since it is a weak acid, it does not ionize fully in an aqueous solution. But this is not the case with other hydrohalic acids. HCl, HBr, and HI ionize very comfortably in an aqueous solution and are strong acids.

Now, we know that fluorine is the most electronegative element in the periodic table. This indicates that the H-F bond is extremely polar.

So you might argue that because of high polarity, the HF bond must also be easy to break, and so HF must actually be a strong acid. But this is not the case.

Recall the trend of the atomic radius in the second row of the periodic table. As we know, across a period, the number of shells of the elements remains the same, but the number of electrons keeps increasing as we go from left to right.

This means that, for an element located on the right, its electrons face a greater force of attraction from the nucleus and hence get pulled inwards.

This results in the atomic radius of the rightmost elements being very small, and the atomic radius also decreases continuously from left to right.

Now, when HF ionizes in an aqueous solution, it forms Fluoride or F-ions. Fluoride ions are formed when fluorine takes up an electron and acquires a negative charge. Also, fluorine is the smallest of all the halogens as it lies on top of group 17 of the periodic table.

Fluorine atomic size order

So an F-ion is extremely unstable because of its extremely small size. Because of its size, the incoming electron faces a huge amount of repulsion from the other electrons, and therefore, fluorine will prefer to not form a fluoride ion at all.

Earlier, we said that for acid to be a strong acid, its conjugate base must be stable.

Here, since F-ion is the conjugate base, and it is very unstable, the HF molecule will prefer to stay in its un-ionized form.

Further, since the fluorine atom is very small, its bond with hydrogen is a very strong bond. 1s orbital of hydrogen and 2p orbital of fluorine share electrons to form the HF molecule.

Contrary to that, HCl has a 1s-3p overlap, HBr has a 1s-4p overlap, and so on. These bonds are all weaker than the 1s-2p strong overlap present in HF so the HF bond is very difficult to break.

Therefore, HF is a weak acid.

Also, the pH of 0.1M of HF is approximately 2.12. Compared to that, HCl has a pH of 1.08 for the same concentration of acid, and this indicates that HCl is a much stronger acid than HF. However, HF is also a very corrosive acid.


Factors Affecting Acidity of a Compound

1. The polarity of the H-X bond: When the H-X bond becomes more polar, the acidity of the compound increases.

The more electronegativity difference there is, the easier the H-X bond will be to break. So, for example, acidity order: HF>H2O>NH3>CH4

HF polarity

2. Size of the X atom: When the X atom is smaller in size, the extent of overlap of the orbitals of H and X is increased. So, the H-X bond becomes tougher to break apart.

As the size of X increases, the H-X bond becomes weaker and can be easily broken to give H+ions. For example, the acidity order:


3. Charge on the acid or base molecule:

As the charge on the molecule increases, the molecule becomes more basic and less acidic. For example, the acidity order:

H3PO4>H2PO4   ——->     HPO42-

4. Oxidation state of the central atom: When various acids consisting of the same central atom are compared, the central atom with the highest oxidation state will be the most acidic in nature.

These acids are formed by those elements which can show a range of different oxidation states. For example, acidity order:




In today’s lesson, we learned that how various compounds are classified as acids or bases according to various theories.

We focused primarily on HF and explored how it is a weak acid. Further, we learned about the pH of HF and about the various factors that affect the acidity of a compound.

I hope it was a fulfilling lesson.

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