Arrhenius & Brønsted-Lowry Acid/Base Models

This is part of the HSC chemistry course under the topic of Properties of Acids and Bases

HSC Chemistry Syllabus

  • Explore the changes in definitions and models of an acid and a base over time to explain the limitations of each model, including but not limited to:
    • Arrhenius’ theory
    • Brønsted–Lowry theory (ACSCH064, ACSCH067)

What are Acids and Bases?

This video explores the definition of acids and bases according to various theories/models including

  • Lavoisier's oxygen theory of acids
  • Davy's hydrogen theory of acids
  • Arrhenius' theory of acids and bases
  • Brønsted-Lowry theory of acids and bases

Oxygen Theory of Acids (1776)

  • French chemist Antoine Lavoisier showed that many non-metal compounds containing oxygen produced acidic solutions when dissolved in water. For example, oxides of carbon, sulfur and phosphorus.

 $$CO_2(aq)+H_2O(l) \rightleftharpoons H_2CO_3(aq)$$


  • Lavoisier hypothesised that the presence of oxygen atoms in these compounds gave them their acidic properties.
  • Limitation: many experiments showed acidic properties in substances which do not contain oxygen e.g. HCl


Hydrogen Theory of Acids (1810)

  • Davy electrolysed samples of hydrochloric acid and showed that it produced hydrogen gas and chlorine gas, but not oxygen gas. This experiment helped to disprove Lavoisier’s oxygen theory of acids.
  • Later experiments by other chemists showed that other acids e.g. hydrocyanic acid (HCN), also contained no oxygen but did contain hydrogen. As a result, Davy proposed that the presence of hydrogen in acids gave them their acidic properties.
  • Limitation: Davy’s theory did not explain why many compounds with hydrogen atoms were not acidic e.g. methane (CH4).


Extension of Davy’s Hydrogen Theory of Acids (1838)

  • German chemist Justus von Liebig stated that acids were substances that contained replaceable hydrogens. He reasoned this by saying that when acids react with metals, the metals replace the hydrogen atoms in the acid to form a salt. 

 $$2Na(s)+2HCl(aq) \rightarrow 2NaCl(aq)+H_2(g)$$


  • Limitation: Von Liebig’s ideas failed to account for the production of gases such as nitrogen dioxide (rather than hydrogen gas) when nitric acid reacts with metals.


Arrhenius’ Theory of Acids and Bases (1887)

Swedish scientist Svante Arrhenius describes:

  • Arrhenius acids are substances that dissociate to produce hydrogen ion in water.

  $$HCl(aq) \rightarrow H^+(aq) + Cl^-(aq)$$


  • Arrhenius bases are substances that dissociate to produce hydroxide ion in water.

 $$NaOH(aq) \rightarrow Na^+(aq) + OH^-(aq)$$


  • Arrhenius’ theory explains neutralisation between acids and bases. He proposed that when neutralisation occurs, the hydrogen ions (produced by acids), reacts with hydroxide ions (produced by bases), to form water


$$HCl(aq) + NaOH(aq) \rightarrow NaCl(aq) + H_2O(l)$$


Limitations of Arrhenius' theory of acids and bases:

  • Arrhenius could not explain why certain compounds do not contain hydroxide ions, despite displaying basic properties. For example, metal oxides and metal carbonates:


$$2HCl(aq) + CaCO_3(s) \rightarrow CaCl_2(aq) + H_2O(l) + CO_2(g)$$


Calcium carbonate (CaCO3) is basic in nature but it does not contain OH.


  • Arrhenius could not explain why some neutrally-charged salts e.g. ZnCl2 are acidic in solution whilst others are basic e.g. Na2S
  • Arrhenius could not explain acid-base reactions that do not occur in aqueous solutions. For example, the reaction between gaseous ammonia and hydrogen chloride is an example of acid-base reaction.

$$NH_3(g) + HCl(g) \leftrightharpoons NH_4Cl(s)$$


    Brønsted-Lowry Theory of Acids and Bases (1923)

    • Brønsted-Lowry acids are protons donors, any compound that can donate protons regardless of its state of matter or the nature of the solvent.
    • Brønsted-Lowry bases are protons acceptors, any compound that can accept protons regardless of its state of matter or nature of the solvent.

    proton vs hydrogen ion


      •  A proton is equivalent to a hydrogen ion (as described by the Arrhenius theory) because when a hydrogen-1 atom loses an electron, only a proton remains.


      $$HCl(aq) + H_2O(l) \rightarrow H_3O^+(aq) + Cl^-(aq)$$

       (Here HCl is a Brønsted-Lowry acid since it deprotonates, and water is considered a Brønsted-Lowry base since it is a proton acceptor)


      $$NH_3(aq) + H_2O(l) \leftrightharpoons NH_4^+(aq) + OH^-(aq)$$

       (Here ammonia is a Brønsted-Lowry base since it accepts a proton from water, and water is considered a Brønsted-Lowry acid since it donates a proton to ammonia)


      • Brønsted-Lowry theory assigned a role to the solvent. In addition to being a solvent, the theory outlines the important ionising role of water. Water can act has both an acid (proton donors) and a base (proton acceptors). 

      Brønsted-Lowry acid-base reaction

      • According to the Brønsted-Lowry theory, an acid-base reaction occurs when there is a transfer of proton(s) from the acid to the base. This can occur in non-aqueous states. 


      Acid-base pairs, conjugate acid and bases

      Consequently, the Brønsted-Lowry theory introduces the concept of acid-base conjugate pairs. When a Brønsted-Lowry acid donates its proton to the base, the resultant anion of the acid is deprotonated. This species can act as a Brønsted-Lowry base (conjugate base) by accepting a proton to reform the original acid molecule.

      • A conjugate base is the substance produced when a Brønsted-Lowry acid deprotonates.

       $$HF(aq) + H_2O(l) \leftrightharpoons H_3O^+(aq) + F^-(aq)$$ 

      When HF ionises in water, it forms its conjugate base, fluoride ion

      The acid-conjugate-base concept only applies to weak acids because the deprotonated anion of strong acids e.g. Cl has very little to no tendency of accepting a proton. In other words, if the Ka of an acid is too high or Kb of its conjugate base is too low, the equilibrium has little reversibility.


      • conjugate acid is the substance produced when a Brønsted-Lowry base accepts a proton.

      For example, protonation of ammonia by water:

       $$NH_3(aq) + H_2O(l) \leftrightharpoons NH_4^+(aq) + OH^-(aq)$$

      When ammonia ionises in water, it forms its conjugate acid - ammonium ion

      Arrhenius versus Brønsted-Lowry Theory

      Overall, the Brønsted-Lowry theory is better than the Arrhenius theory because it can explain the acid and base nature of more species. For example:

      • Bases that do not contain hydroxide ions e.g. ammonia, NH3. In Arrhenius’ theory, ammonia is not considered a base because it does not contain nor dissociate into hydroxide ions (OH).

      $$ NH_3(aq) + H_2O(l) \leftrightharpoons NH_4^+(aq) + OH^-(aq)$$


      • Acid and base reactions that occur in non-aqueous solution, namely the reaction between ammonia and hydroxide chloride gas. In this reaction, ammonia accepts a proton from hydrogen chloride (base) while hydrochloride donates a proton to ammonia (acid).

       $$NH_3(g) + HCl(g) \leftrightharpoons NH_4Cl(s)$$


      The Arrhenius theory does not explain the acid base nature of the two reactants above because they are in gaseous states (non-aqueous).


      Major Advantages of Brønsted-Lowry Theory

      The Brønsted-Lowry theory of acids and bases offers many advantages over the Arrhenius theory:


      1. It explains the basic property of substances that do not contain hydroxide ions.
      2. It expands the role of water in acid-base reactions as more than just a solvent.
      3. It can be expanded to include solvents other than water and reactions that occur in non-aqueous states.
      4. It links acids and bases into conjugate acid-base pairs. This concept
        • explains the acidity and basicity of acidic and basic salts respectively
        • explains the relative strengths of acids and bases according to the strength of their conjugate partner
      5. Explains the acidic and basic properties of amphiprotic substances


      Limitations of Brønsted-Lowry Theory

      • Does not explain the acidity of acidic oxides such as SO2 and SO3
      • Does not explain the basicity of basic oxides such as MgO and CaO
      • Does not explain reactions between acidic and basic oxides as they do not involve proton transfer. For example, $$SO_3(g) + CaO(s) \rightarrow CaSO_4(s)$$
      • The above limitations are explained in the Lewis Theory of acid and bases which states:
        • Acids are electron pair acceptors
        • Bases are electron pair donors