Advanced A level theoretical chemistry - acid-base equilibrium

Part 5.1 Acid–Base Theory – Lewis & Bronsted–Lowry Theories

[Author ©  Dr Phil Brown PhD: Doc Brown's exam revision notes suitable for A level chemistry students of advanced pre–university/college advanced level theoretical–physical chemistry courses:  acid–base equilibrium revision notes on Bronsted–Lowry and Lewis theories [updated April 29th 2026 *]

ALL my advanced A level theoretical chemistry revision study notes

Full index for Part 5 acid-base equilibria and calculations

5.0 Basic notes and equations on acids, bases, salts, uses of acid–base titrations

5.1 Equilibria:  Lewis and Bronsted–Lowry acid–base theories (sub–index for this page)

5.2 Self–ionisation of water and pH scale

5.3 Strong acids – examples and pH calculations

5.4 Weak acids – examples & pH, K_a and pK_a calculations

5.5 Strong bases – examples and pH calculations

5.6 Weak bases – examples and pH, K_b and pK_b calculations

5.1 Acid–base theory – Lewis and Bronsted–Lowry theories

Sub–index for this page section 5.1

5.1.0 Introduction

5.1.1 Lewis acid–base electron pair theory

5.1.2 Bronsted–Lowry acid–base theory

5.1.3 Examples of soluble substances giving aqueous solution acid–base interactions

5.1.4 Why can acids can be described as monobasic, dibasic or tribasic etc.?

5.1.5 Examples of water insoluble bases giving acid–base neutralization reactions

5.1.6 Examples of two solids reacting together in an acid–base reaction

5.1.0 Introduction

This page explains the Lewis theory of acids (electron pair acceptors) & bases (electron pair donors) and the Bronsted–Lowry theory of acids (proton donors) and bases (proton acceptors).

The terms conjugate acid, conjugate base and conjugate base are also explained via fully described acid–base reactions.

This sets the background for the rest of sections 5.1 to 5.6

 (I shouldn't use it, but M = old fashioned shorthand for mol dm^–3 !)

You should read 5.0 revision first, albeit an afterthought on my part.

• Basic ideas on acids, bases and their reactions, pH scale, using indicators and simple acid–base theory are described on the GCSE notes pages and are essential reading before tackling parts 5 and 6 of these more advanced notes, and much of it is not repeated here.

• These notes in Parts 5. and 6. involve a major upgrade from GCSE grade 9–1 notes on Acids, bases, salts, pH and neutralization, though these notes do describe the basic ideas on pH, examples of solution pH's, indicators and the reactions of acids with metals, soluble/insoluble oxides, hydroxides, carbonates, hydrogencarbonates and aqueous ammonia, salt preparations and introduction to pH titration curves, all this you should know.

5.1.1 Lewis acid–base electron pair theory

• A base is an electron pair donor and an acid is an electron pair acceptor.

• e.g. a non B–L, but a Lewis acid–base interaction is boron trifluoride (Lewis–acid, electron pair acceptor) reacting with ammonia (Lewis–base, electron pair donor).

• F₃B + :NH₃ ==> F₃B–NH₃

• Note: In organic chemistry mechanisms, nucleophiles are Lewis bases and electrophiles are Lewis acids and they may fit into the Bronsted–Lowry definition too e.g. protonation of alcohols and alkenes via acid.

• In Transition Metal chemistry, ligands like water, can donate a pair of non–bonding electrons (lone pair) into a vacant orbital of a central metal ion and so dative covalent (co–ordinate) bonds hold a complex together.

• The central metal ion with vacant bonding orbitals can act as a Lewis acid by accepting an electron pair to form a dative covalent bond.

• Ligands act as Lewis bases by electron pair donation to form the metal–ligand co–ordinate bond.

• See Transition Metal complexes and ligands for more details

• –

5.1.2 Bronsted–Lowry acid–base theory

• An acid is a proton donor and a base is a proton acceptor.

• Bronsted–Lowry acids and bases are a 'sub–set' of the general Lewis acid–base theory, namely acids are electron pair acceptors and bases are electron pair donors.

• All bases X:, will have a lone pair of non–bonding electrons that will except the electron deficient proton H⁺ to form a covalent X–H bond.

  • In general, a Lewis acid – Lewis base interaction involves the formation of a single dative covalent/co–ordinated bond where the bonding pair of electrons is donated by the base to the electron pair accepting acid.

  • The Bronsted–Lowry theory concentrates on proton donation and acceptance.

• The oxonium ion, H₃O⁺_(aq) (or more simply, the aqueous hydrogen ion, H⁺) is formed by any acidic substance in water.

  • Increase in H⁺ concentration decreases the pH of a solution. (see section 5.2)

• The hydroxide ion, OH^–_(aq), is formed by any soluble base forming an alkaline solution.

  • Increase in OH^– concentration increases the pH of a solution. (see section 5.2)

• Incidentally water is a neutral oxide because its pH is 7, logistically the oxonium/hydrated proton ion concentration equals the hydroxide ion concentration ...

  •  [H₃O⁺_(aq)] = [OH^–_(aq)] via the tiny fraction of water molecules undergoing dissociation or self–ionisation because of the reaction

    • 2H₂O_(l) [H₃O⁺_(aq)] + [OH^–_(aq)]

    • water <=> oxonium/hydrogen ion + hydroxide ion

  • BUT, in this reaction, water acts as both acid and base i.e. one water molecule (acid) donates a proton to another water molecule which becomes an oxonium ion (hydrated proton) and another water molecule (base) simultaneously accepts a proton!

  • Therefore water is an amphoteric oxide i.e. it reacts as both a proton acceptor and a proton donator.

    • e.g. water acting as a base – proton acceptor with a stronger acid like the hydrogen chloride gas

    •  HCl_(g) + H₂O_(l) ==> H₃O⁺_(aq) + Cl^–_(aq)

    • This is how hydrochloric acid is formed which you write simply as HCl.

    • e.g. water acting as an acid – proton donor with a weak BUT stronger base like ammonia gas

    • NH_3(aq) + H₂O_(l) NH₄⁺_(aq) + OH^–_(aq)

    • This is why ammonium solution is alkaline – sometimes wrongly called 'ammonium hydroxide' instead of aqueous ammonia.

  • More details on these reactions are given in subsequent sections on this web page.

  • –

5.1.3 Examples of soluble substances giving aqueous solution acid–base interactions

• 5.1.3a: Sulfuric acid:

  • H₂SO_4(l) + 2H₂O_(l) ==> 2H₃O⁺_(aq) + SO₄^2–_(aq)

  • Sulfuric acid, H₂SO₄, is the acidic proton donor and H₂O is the proton accepting base, however, the 2nd ionisation is weak, so the equation represents the maximum possible proton donation with a base.

  • Note the products are also acids and bases:

  • H₃O⁺ is the conjugate acid of the base H₂O

  • SO₄^2– is the conjugate base of the acid H₂SO₄

  • The conjugate acid and original base or the conjugate base and the original acid are known as a conjugate pair and are related by proton transfer.

  • –

• 5.1.3b: Hydrogen chloride gas:

  • HCl_(g) + H₂O_(l) ==> H₃O⁺_(aq) + Cl^–_(aq)

  • HCl is the acid and Cl^– is the conjugate base.

  • H₂O is the base and H₃O⁺ is the conjugate acid.

  • The resulting solution is called hydrochloric acid.

  • –

• 5.1.3c: Ammonia:

  • NH_3(aq) + H₂O_(l) NH₄⁺_(aq) + OH^–_(aq)

  • Ammonia is the base and the ammonium ion, NH₄⁺, is its conjugate acid,

  • and water is the acid and the hydroxide ion is its conjugate base.

  • –

• 5.1.3d: The hydrogen carbonate ion, HCO₃^–,

  • This can act as an acid with a base or act as a base with an acid, such behaviour is described as amphoteric.

  • HCO₃^– + H₃O⁺_(aq) ==> 2H₂O_(l) + CO_2(aq)

    • HCO₃^– acting as a base, accepting a proton from an acid.

  • HCO₃^– + OH^–_(aq) ==> H₂O_(l) + CO₃^2–_(aq)

    • HCO₃^– acting as an acid, donating a proton to the hydroxide ion base.

• 5.1.3e: Neutralisation

  • Since any soluble base gives hydroxide ions in aqueous and any soluble acid gives oxonium/hydrogen ions, they combine to form water.

  • The  ionic equation for these neutralisations is:

  • H₃O⁺_(aq) + OH^–_(aq) ==> 2H₂O_(l)

  • or more simply: H⁺_(aq) + OH^–_(aq) ==> H₂O_(l)

  • More reactions of H₃O⁺/H⁺ are given in 5.1.4

  • –

5.1.4 Why can acids can be described as monobasic (monoprotic), dibasic (diprotic) or tribasic (triprotic)?

• These terms depend on the maximum number of protons that are available for transfer from an acid in an acid–base reaction.

• The terms also applies to the maximum number of protons the final conjugate base can accept.

• monobasic acids can donate one proton per 'molecule' e.g.

  • hydrochloric HCl, nitric HNO₃, ethanoic CH₃COOH (the alkyl H's are not acidic),

  • e.g. with sodium hydroxide, giving the salts of formulae

  • NaCl, NaNO₃ and CH₃COONa with sodium hydroxide respectively.

  • The conjugate bases are Cl^–, NO₃^– and CH₃COO^– respectively.

  • –

• dibasic acids can donate two protons per 'molecule'

  • e.g. with sodium hydroxide

  • sulfuric H₂SO₄ ==> NaHSO₄ ==> Na₂SO₄

  • and the conjugate dibasic base is SO₄^2–

  • ethanedioic (COOH)₂ ==> HOOC–COONa ==> NaOOC–COONa or (COONa)₂

  • and the three isomeric benzene–x,y–dicarboxylic acids (x,y = 1,1 and 1,2 and 1,3)

  • giving three possible isomeric C₆H₄(COOH)₂, so all three will all give two possible salts – half neutralised and fully neutralised.

  • –

• tribasic acids can donate up to three protons per 'molecule'

  • and will give three possible salts with sodium hydroxide e.g.

  • boric acid H₃BO₃, phosphoric(V) H₃PO₄

  • The tribasic conjugate bases are theoretically BO₃^3– and PO₄^3–

  • Na₃BO₃ and Na₃PO₄, representing full neutralisation and the 'intermediates' being 'theoretically': NaH₂BO₃ ==> Na₂HBO₃ and NaH₂PO₄. ==> Na₂HPO₄

  • citric acid , the middle–left hydrogen of the HO–C (alcohol) is not acidic in water,

  • the trisodium salt is formed with excess sodium hydroxide via the monosodium and disodium salts.

  • –

5.1.5 Examples of water insoluble bases giving acid–base neutralization reactions.

• 5.1.5a: Water insoluble copper(II) oxide dissolving in dil. sulfuric acid.

  • CuO_(s) + H₂SO_4(aq) ==> CuSO_4(aq) + H₂O_(l)

  • but omitting any non–changing/reacting 'spectator' ions the actual ionic reaction is

  • CuO_(s) + 2H₃O⁺_(aq) ==> Cu²⁺_(aq) + 3H₂O_(l)

  • or more simply: CuO_(s) + 2H⁺_(aq) ==> Cu²⁺_(aq) + H₂O_(l)

  • where CuO is the insoluble base and H₃O⁺ is the acid. Effectively, the oxide ion, O^2–, acting as a base, gains two protons to form water.

  • Incidentally, a group 6, connection (O and S), copper(II) sulphide reacts with acids to form the copper(II) salt and hydrogen sulphide gas (hydrogen sulphide, rotten egg smell and very harmful, not just to our aesthetics!).

  • e.g. Copper(II) sulphide will dissolve in dil. hydrochloric acid to form a solution of copper(II) chloride and release hydrogen sulphide.

  • CuS_(s) + 2HCl_(aq) ==> CuCl_2(aq) + H₂S_(g)

  • CuS_(s) + 2H⁺_(aq) ==> Cu²⁺_(aq) + H₂S_(g)

  • where CuS is the insoluble base and HCl/H⁺/H₃O⁺ is the acid. Effectively, the sulphide ion, S^2–, acting as a base, gains two protons to form hydrogen sulfide.

  • –

• 5.1.5b: Water insoluble calcium carbonate dissolving in hydrochloric acid.

  • CaCO_3(s) + 2HCl_(aq) ==> CaCl_2(aq) + H₂O_(l) + CO_2(g)

  • full ionic equation: CaCO_3(s) + 2H₃O⁺_(aq) ==> Ca²⁺_(aq) + 3H₂O_(l) + CO_2(g)

  • or more simply: CaCO_3(s) + 2H⁺_(aq) ==> Ca²⁺_(aq) + H₂O_(l) + CO_2(g)

  • where CaCO₃ is the insoluble base and HCl/H₃O⁺ is the acid. Again, effectively, the carbonate ion, CO₃^2–, acting as a base, gains two protons to form water and carbon dioxide.

  • –

5.1.6 Examples of two solids reacting together in an acid–base reaction.

• 5.1.6a: When solid ammonium salts are heated with solid calcium hydroxide (slaked lime) ammonia gas is produced.

  • 2NH₄Cl_(s) + Ca(OH)_2(s) ==> CaCl_2(s) + 2H₂O_(l) + 2NH_3(g)

  • ionically: NH₄⁺_(s) + OH^–_(s) ==> H₂O_(l) + NH_3(g)

  • The ammonium ion is the acid and the hydroxide ion the base.

  • –

WHAT NEXT?

Advanced Equilibrium Chemistry Notes Part 1. Equilibrium, Le Chatelier's Principle–rules * Part 2. K_c and K_p equilibrium expressions and calculations * Part 3. Equilibria and industrial processes * Part 4 Partition between two phases, solubility product K_sp, common ion effect, ion–exchange systems * Part 5. pH, weak–strong acid–base theory and calculations * Part 6. Salt hydrolysis, acid–base titrations–indicators, pH curves and buffers * Part 7. Redox equilibria, half–cell electrode potentials, electrolysis and electrochemical series * Part 8. Phase equilibria–vapour pressure, boiling point and intermolecular forces watch out for sub–indexes to multiple sections or pages

Chemical Equilibrium Notes Parts 5 & 6 Index