A level Acid-base theory, Lewis acids and bases, Bronsted-Lowry proton theory GCE AS A2 chemistry revision notes KS5

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Theoretical-Physical Advanced Level Chemistry - Equilibria - Chemical Equilibrium RevisionNotes PART 5.1

5.1 Acid-Base Theory - Lewis and Bronsted-Lowry Theories

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

acceptors). The terms conjugate acid, conjugate base and conjugate base are also explainedvia fully described acid-base reactions.

GCSE/IGCSE reversible reactions-equilibrium notes * GCSE/IGCSE notes on acids

and bases

Equilibria Part 5 sub-index: 5.1 Lewis and Bronsted-Lowry acid-base theories * 5.2 self-ionisation of water and pH scale * 5.3 strong acids-examples-calculations * 5.4 weak acids-examples & pH-Ka-pKa calculations * 5.5 strong bases-examples-pH calculations * 5.6 weak

bases- examples & pH-Kb-pKb calculations * 5.7 A level notes on Acids, Bases, Salts, uses ofacid-base titrations - upgrade from GCSE!

Advanced Equilibrium Chemistry Notes Part 1. Equilibrium, Le Chatelier's Principle-rules * Part2. Kc and Kp equilibrium expressions and calculations * Part 3. Equilibria and industrial

processes * 4. Partition, solubility product and ion-exchange * 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 andelectrochemical series * Part 8. Phase equilibria-vapour pressure, boiling point and

intermolecular forces

5.1 Acid-base theory - Lewis and Bronsted-Lowry

Basic ideas on acids, bases and their reactions, pH scale, using indicators and simple

A level Acid-base theory, Lewis acids and bases, Bronsted-Lowry proton theory GCE AS A2 chemistry revision notes KS5

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acid-base theory are described on the GCSE notes pages and are essential readingbefore tackling parts 5 and 6 of these more advanced notes, and much of it is notrepeated here.

An upgraded from GCSE Acids, bases, salts, pH and neutralization describes thebasic ideas on pH, examples of solution pH's, indicators and the reactions of (ii) acidswith metals, soluble/insoluble oxides, hydroxides, carbonates, hydrogencarbonates andaqueous ammonia, salt preparations and introduction to pH titration curves.

Lewis acids and bases and the Bronsted-Lowry theory of acids and bases

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 basesand 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 ionand so dative covalent (co-ordinate) bonds hold a complex together.

The central metal ion with vacant bonding orbitals can act as a Lewis acidby 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 & ligands for more details

5.1.1: An acid is a proton donor and a base is a proton acceptor - Bronsted-Lowry acid-base theory

Bronsted-Lowry acids and bases are a 'sub-set' of the general Lewis acid-basetheory, namely acids are electron pair acceptors and bases are electron pairdonors.

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

In general, a Lewis acid - Lewis base interaction involves theformation of a single dative covalent/co-ordinated bond where thebonding pair of electrons is donated by the base to the electron pairaccepting acid.

The Bronsted-Lowry theory concentrates on proton donation andacceptance.

The oxonium ion, H3O+(aq) (or more simply, the aqueous hydrogen ion, H+) is

formed by any acidic substance in water.

+

A level Acid-base theory, Lewis acids and bases, Bronsted-Lowry proton theory GCE AS A2 chemistry revision notes KS5

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Increase in H concentration decreases the pH of a solution. (see section5.2)

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. (seesection 5.2)

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

[H3O+(aq)] = [OH–

(aq)] via the tiny fraction of water molecules undergoingdissociation 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 watermolecule (acid) donates a proton to another water molecule whichbecomes 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 protonacceptor and a proton donator.

e.g. water acting as a base - proton acceptor with a stronger acidlike the hydrogen chloride gas

HCl(g) + H2O(l) ==> H3O+(aq) + Cl–(aq)

This is how hydrochloric acid is formed which you writesimply as HCl.

e.g. water acting as an acid - proton donor with a weak BUTstronger base like ammonia gas

NH3(aq) + H2O(l) NH4+

(aq) + OH–(aq)

This is why ammonium solution is alkaline - sometimeswrongly called 'ammonium hydroxide' instead of aqueousammonia.

More details on these reactions are given in subsequent sections on thisweb page.

5.1.2: Examples of soluble substances giving aqueous solutionacid-base interactions

5.1.2a: Conc. sulphuric acid: H2SO4(l) + 2H2O(l) ==> 2H3O+(aq) + SO4

2–(aq)

Sulphuric acid, H2SO4, is the acidic proton donor and H2O is the protonaccepting base.

A level Acid-base theory, Lewis acids and bases, Bronsted-Lowry proton theory GCE AS A2 chemistry revision notes KS5

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Note the products are also acids and bases:

H3O+ is the conjugate acid of the base H2O

SO42– is the conjugate base of the acid H2SO4

The conjugate acid and original base or the conjugate base andthe original acid are known as a conjugate pair and are relatedby proton transfer.

5.1.2b: 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.

5.1.2c: Ammonia: NH3(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.

5.1.2d: The hydrogen carbonate ion, HCO3–, 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) + CO2(aq)

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

HCO3– + OH–

(aq) ==> H2O(l) + CO32–

(aq)

HCO3– acting as an acid, donating a proton to the hydroxide ion

base.

5.1.2e: Since any soluble base gives hydroxide ions in aqueous and any solubleacid gives oxonium/hydrogen ions, they combine to form water. The ionicequation 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

5.1.3: Acids can be described as monobasic, dibasic or tribasic etc. depending onthe maximum number of protons that are available for transfer in an acid-base reaction.The terms mono/di/triprotic are used to mean the same thing, the term then applies tothe maximum number of protons the final conjugate base can accept.

monobasic acids e.g.

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

A level Acid-base theory, Lewis acids and bases, Bronsted-Lowry proton theory GCE AS A2 chemistry revision notes KS5

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acidic),

giving the salts e.g. NaCl, NaNO3 and CH3COONa with sodium hydroxiderespectively.

dibasic acids e.g. with sodium hydroxide

sulphuric H2SO4 ==> NaHSO4 ==> Na2SO4

ethanedioic (COOH)2 ==> HOOC-COONa ==> NaOOC-COONa or(COONa)2

and the three isomeric benzene-x,y-dicarboxylic acids (x,y = 1,1 1,2 and1,3) C6H4(COOH)2,

tribasic acids e.g.

boric acid H3BO3, phosphoric(V) H3PO4,

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

the trisodium salt is formed with excess sodiumhydroxide via the monosodium and disodium salts.

5.1.4: Examples of water insoluble bases giving acid-base neutralizationreactions.

5.1.4a: Water insoluble copper(II) oxide dissolving in dil. sulphuric acid.

CuO(s) + H2SO4(aq) ==> CuSO4(aq) + H2O(l)

but omitting any non-changing/reacting 'spectator' ions the actual ionicreaction 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, O2–, acting as a base, gains two protons to formwater.

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

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

A level Acid-base theory, Lewis acids and bases, Bronsted-Lowry proton theory GCE AS A2 chemistry revision notes KS5

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CuS(s) + 2HCl(aq) ==> CuCl2(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, S2–, acting as a base, gains twoprotons to form hydrogen sulfide.

5.1.4b: Water insoluble calcium carbonate dissolving in hydrochloric acid.

CaCO3(s) + 2HCl(aq) ==> CaCl2(aq) + H2O(l) + CO2(g)

full ionic equation: CaCO3(s) + 2H3O+(aq) ==> Ca2+

(aq) + 3H2O(l) +CO2(g)

or more simply: CaCO3(s) + 2H+(aq) ==> Ca2+

(aq) + H2O(l) + CO2(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 protonsto form water and carbon dioxide.

5.1.5: Examples of two solids reacting together in an acid-base reaction.

5.1.5a: When solid ammonium salts are heated with solid calcium hydroxide(slaked lime) ammonia gas is produced.

2NH4Cl(s) + Ca(OH)2(s) ==> CaCl2(s) + 2H2O(l) + 2NH3(g)

ionically: NH4+

(s) + OH–(s) ==> H2O(l) + NH3(g)

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

-

Equilibria Part 5 sub-index: 5.1 Lewis and Bronsted-Lowry acid-base theories * 5.2 self-ionisation of water and pH scale * 5.3 strong acids-examples-calculations * 5.4 weak acids-examples & pH-Ka-pKa calculations * 5.5 strong bases-examples-pH calculations * 5.6 weak

bases- examples & pH-Kb-pKb calculations * 5.7 A level notes on Acids, Bases, Salts, uses ofacid-base titrations - upgrade from GCSE!

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A level Acid-base theory, Lewis acids and bases, Bronsted-Lowry proton theory GCE AS A2 chemistry revision notes KS5

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