Bronsted-Lowry Concept of Acids and Bases

 

Johannes Nicolaus Brønsted (1879-1947) & Thomas Martin Lowry (1874-1936)

J.N. Bronsted was a Danish scientist, while T.M. Lowry was a British chemist, both independently proposing a definition for acids and bases in 1923. Their collaborative work laid the foundation for the modern understanding of acids and bases in chemistry.

Johannes Nicolaus Bronsted (1879–1947) was a Danish physical chemist born in Varde, Denmark. In 1923, Bronsted introduced the Bronsted-Lowry acid-base theory alongside Thomas Martin Lowry.

Thomas Martin Lowry (1874–1936), a British physical chemist born in Low Moor, Bradford, England, independently proposed a similar theory around the same time as Bronsted. This theory expanded the concept of acids and bases beyond the limitations of the earlier Arrhenius definition, which defined acids as substances that release hydrogen ions (H⁺) in aqueous solutions and bases as substances that release hydroxide ions (OH⁻). The Bronsted-Lowry theory is more versatile, as it can be applied to non-aqueous systems and reactions that do not involve the transfer of hydroxide ions.

Bronsted and Lowry's contributions significantly influenced the field of acid-base chemistry, providing a more comprehensive and widely applicable framework for understanding the behavior of acids and bases in various chemical reactions.

According to this theory, an acid is a substance that can donate a proton, whereas a base can accept protons.

Acid: Proton donor                                             Base: Proton acceptor

2NH₃ ⇋ NH₄⁺ + NH₂^-, in this self-reaction NH₄⁺ is an acid and NH₂^- is a base.

2H₂O ⇋ H₃O⁺+ OH^-, here H₃O⁺ is an acid and OH^- is a base.

This theory also gives us a new term called conjugate acids and bases which are related through proton transfer process.

HCl + NH₃ ⇋ NH₄⁺+ Cl^-

In this reaction HCl and NH₃ is the acid and base while NH₄⁺ and Cl^- is the conjugate acid and base  

NH₃ + H⁺ ⇋ NH₄⁺

Here NH₄⁺ is the conjugate acid of NH₃ base as in the reverse reaction NH₄⁺ can donate H⁺ to form NH₃.

HCl ⇋ H⁺ + Cl^-

Here Cl^- is the conjugate base of HCl as in the reverse reaction it can have a tendency to accept the H⁺ ions.

Therefore the equilibrium between acid and bases with their conjugate counterpart can be written as follows.

Acid (A) + Base (B) ⇋ Conjugate Acid (CA) + Conjugate Base (CB)          

Here the properties of water are specially mentioned as it can act as a base toward strong acid and an acid toward strong base. Such properties of water are called amphoteric properties.

HCl+ H₂O ⇋ H₃O⁺ + Cl^-

NH₃+ H₂O ⇋ NH₄⁺ + Cl^-

In qualitative sense, we can also draw the same conclusion for H₂O with Arrhenius theory. Therefore the acidic behaviour of metal ions hydrated in the water can also be explained very easily.

[Fe(H₂O)₆]³⁺ + H₂O ⇋ [Fe(H₂O)(OH)]²⁺ + H₃O⁺

Therefore it can be seen the presence of metal ions makes the aqueous medium a little bit acidic. This phenomenon is also true for many organometallic hydrides which have their own acidic properties.

[H₂Fe(CO)₄] + H₂O ⇋ [HFe(CO)₄]^- + H₃O⁺

This theory also shed some light on autoionization of amphoteric compounds. This can also be explained as

2NH₃ ⇋ NH₄⁺ + NH₂^-

2H₂O ⇋ H₃O⁺+ OH^-

Acid-base properties of oxides can also be explained with this concept. In general, oxides of non-metal compounds are acidic in nature such as CO₂, P₂O₃, SO₂ etc whereas metal oxides are basic in nature. Although for metal oxides, acidic nature increases with the increase of oxidation state of the central metal ion.

CrO (basic oxide) Cr2O3 (amphoteric acid) CrO3 (acidic)

The Bronsted-Lowry concept is very useful in rationalizing acidic properties of hydrated metal ions, acidic properties of non-metallic oxides, acidic properties of organometallic hydrides, hydrolysis and amphoteric reactions.

SO₃ + 2H₂O → HSO₄^-+H₃O⁺ (Acidic properties of SO₃)

CaO + 2H₃O⁺ → Ca²⁺+ 3H₂O (Basic properties of CaO)

Al₂O₃ + 6H₃O⁺ + 3H₂O →2Al(H₂O)₆³⁺                Al₂O₃+7H₂O → 2Al(OH)₄^-+2H₃O⁺  

Since Al₂O₃ reacts with both acids and bases, this property is called its amphoteric property.

Merits of the Bronsted-Lowry Concept:

1) Substance-Centric Definition:

The Bronsted-Lowry concept defines acids and bases based on the properties of the substances themselves, rather than relying on the ions generated during the ionization in the solvent. This approach provides a more direct characterization of acidic and basic properties.

2) Hydrolysis of Salts:

To elucidate the hydrolysis of a salt, it considers the relative proton donor capacity and proton acceptor capacity of the solvated cation and anion, respectively, compared to that of bulk H₂O. If the solvated cation has a higher proton donor capacity than H₂O, the solution tends to be acidic. Conversely, if the solvated anion exhibits a higher proton acceptor capacity than H₂O, the solution tends to be basic. For instance, in FeCl₃ solution, where Cl^- is the conjugate base of the strong acid HCl, the predominant proton transfer equilibrium is:

In FeCl₃ solution, we get Fe(H₂O)₆³⁺ and hydrated Cl^- ions. Cl^- being the conjugate base of the strong acid, HCl is much weaker than H₂O as a base. Hence the predominant proton transfer equilibrium is:

Fe(H₂O)₆³⁺ + H₂O ⇋ Fe (H₂O)₅(OH)²⁺ + H₃O⁺

And the solution becomes acidic.

A similar solution arises for NH₄Cl.

In Na₂CO₃ solution, Na(H₂O)_x⁺ is almost comparable with the bulk H2O as a proton donor and the predominant proton transfer equilibrium is: CO₃^2- + H₂O ⇋ HCO₃^- + OH^-.

In CH₃CO₂NH₄ solution, NH₄⁺ is a stronger proton donor than H₂O and CH₃CO₂^- is also a stronger proton acceptor than H₂O. Hence the predominant proton transfer equilibrium is:

NH₄⁺ + CH₃CO₂^- ⇋ NH₃ + CH₃CO₂H.

3) The concept can explain the amphoteric reactions.

4) The concept can explain the acid base behavior of organometallic hydrides.

 

Limitations of this theory

Exclusively Proton-Centric Definition:

The Bronsted-Lowry theory is limited to proton transfer reactions. It defines acids as substances that donate protons (H⁺) and bases as substances that accept protons. While this definition is applicable to a wide range of reactions, there are some instances where other types of electron-pair interactions play a crucial role in acid-base behavior.

Water as the Solvent:

The theory was initially formulated based on reactions in aqueous solutions. While it can be extended to non-aqueous systems, the original framework was developed primarily for reactions occurring in water. In certain non-aqueous solvents, the behavior of acids and bases may deviate from the predictions of the Bronsted-Lowry theory.

Does Not Account for Lewis Acids and Bases:

The Bronsted-Lowry theory does not encompass all acid-base reactions. It does not address reactions where the acid-base interaction involves the donation or acceptance of an electron pair rather than a proton. This limitation led to the development of the Lewis acid-base theory by Gilbert N. Lewis, which provides a broader definition of acids and bases.

Complex Reaction Mechanisms:

In some reactions, particularly those involving complex reaction mechanisms, it may be challenging to assign clear roles of acids and bases based solely on proton transfer. In these cases, a more comprehensive understanding of reaction mechanisms beyond acid-base interactions may be necessary.

Equilibrium Constant Variation:

The equilibrium constants (K) for some reactions involving Bronsted-Lowry acids and bases may not always correlate with the strength of the acids or bases involved. Factors such as the solvent and other reaction conditions can influence the equilibrium position, making it less straightforward to predict the extent of reaction based solely on the acidity or basicity of the species involved.

Despite these limitations, the Bronsted-Lowry theory remains a valuable and widely used concept in chemistry. It provides a useful framework for understanding a broad range of acid-base reactions, especially in aqueous solutions, and serves as a foundation for more advanced theories like the Lewis acid-base theory.

Reference

1) Concise inorganic chemistry by J. D. Lee.

2) Inorganic Chemistry by James E. Huheey, Ellen A Keither, Richard L. Keither, Okhil K. Medhi.