Brown's Chemistry Advanced A Level Notes - Theoretical–Physical
Chemistry – Equilibria – Chemical Equilibrium Revision Notes PART 4
Part 5.1 Acid–Base Theory
– Lewis & Bronsted–Lowry
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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.
(I shouldn't use it, but
M = old fashioned shorthand for mol dm–3 !)
Lewis and Bronsted-Lowry acid-base theories
Self-ionisation of water and pH scale
Strong acids - examples and pH calculations
Weak acids - examples & pH, Ka and pKa calculations
Strong bases - examples and pH calculations
Weak bases - examples and pH, Kb and pKb calculations
Basic notes and equations on acids, bases, salts,
acid-base titrations - upgrade from GCSE!
5.1 Acid–base theory – Lewis and
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.
Lewis acids and bases and the Bronsted–Lowry theory of acids and bases
Lewis acid–base electron
A base is an electron pair donor and an
acid is an electron pair
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).
+ :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
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.
5.1.1: An acid is a
proton donor and a base is a
proton acceptor – Bronsted–Lowry acid–base theory
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.
Bronsted–Lowry theory concentrates on proton donation and
ion, H3O+(aq) (or more simply,
the aqueous hydrogen ion, H+) is formed by any
acidic substance in water.
ion, OH–(aq), is formed by any soluble
base forming an alkaline solution.
water is a neutral oxide because its pH is 7, logistically the
oxonium/hydrated proton ion concentration equals the hydroxide ion
= [OH–(aq)] via the tiny fraction of water
molecules undergoing dissociation or self–ionisation because of the
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!
water is an amphoteric oxide i.e. it reacts as both a proton
acceptor and a proton donator.
More details on
these reactions are given in subsequent sections on this web page.
5.1.2: Examples of
soluble substances giving aqueous solution acid–base interactions
5.1.2a: Conc. sulphuric acid:
+ 2H2O(l) ==> 2H3O+(aq)
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
products are also acids and bases:
is the conjugate acid of the base H2O
is the conjugate base of the acid H2SO4
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.
chloride gas: HCl(g) + H2O(l) ==> H3O+(aq) + Cl–(aq)
HCl is the
acid and Cl– is the conjugate base.
is the base and H3O+ is the conjugate acid.
resulting solution is called hydrochloric acid.
NH4+(aq) + OH–(aq)
the base and the ammonium ion, NH4+, is
its conjugate acid,
and water is
the acid and the hydroxide ion is its conjugate base.
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.
5.1.2e: 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
+ OH–(aq) ==>
+ OH–(aq) ==>
reactions of H3O+/H+
are given in 5.1.4
5.1.3: Acids can be described as monobasic,
dibasic or tribasic etc...
.... depending on the 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 to the maximum number of protons the final conjugate base
hydrochloric HCl, nitric
HNO3, ethanoic CH3COOH
(the alkyl H's are not acidic),
giving the salts e.g.
NaCl, NaNO3 and CH3COONa with sodium hydroxide
e.g. with sodium hydroxide
==> NaHSO4 ==> Na2SO4
==> HOOC–COONa ==> NaOOC–COONa or (COONa)2
and the three isomeric benzene–x,y–dicarboxylic acids
(x,y = 1,1 1,2 and 1,3) C6H4(COOH)2,
will all give two possible salts.
acids will give three possible salts with sodium hydroxide e.g.
boric acid H3BO3, phosphoric(V)
, the middle–left hydrogen
of the HO–C (alcohol) is not acidic in water,
the trisodium salt
formed with excess sodium hydroxide via the monosodium and disodium
5.1.4: Examples of
insoluble bases giving acid–base neutralization reactions.
5.1.5: Examples of
two solids reacting together in an acid–base reaction.
Advanced Equilibrium Chemistry Notes Part 1. Equilibrium,
Le Chatelier's Principle–rules
* Part 2. Kc and Kp equilibrium expressions and
calculations * Part 3.
Equilibria and industrial processes * Part 4
Partition between two
phases, solubility product Ksp, 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
pressure, boiling point and intermolecular forces watch out for sub-indexes
to multiple sections or pages
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Chemical Equilibrium Notes Parts 5 & 6 Index