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Brown's Chemistry Theoretical
Chemistry - Equilibria - Chemical Equilibrium 5.1
5.1 Acid-Base Theory - Lewis and Bronsted-Lowry
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KS4 Science GCSE/IGCSE notes acid-base
theory
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 of
acid-base titrations - upgrade from GCSE!
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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 * 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 and electrochemical 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 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.
-
An upgraded from GCSE
Acids, bases, salts, pH and
neutralization describes the basic ideas on pH, examples of
solution pH's, indicators and the reactions of (ii) acids with
metals, soluble/insoluble oxides, hydroxides, carbonates,
hydrogencarbonates and aqueous ammonia, salt preparations and
introduction to pH titration curves.
-
Lewis acids and bases and the Bronsted-Lowry theory of acids and bases
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Lewis acid-base 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.
-
Ligands act as Lewis
bases by electron pair donation to form the metal-ligand bond.
-
5.1.1: 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.
-
The hydroxide
ion, OH-(aq), is formed by any soluble
base forming an alkaline solution.
-
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
-
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.
-
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
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5.1.2a: Conc. sulphuric acid: H2SO4(l)
+ 2H2O(l) ==> 2H3O+(aq)
+ SO42-(aq)
-
Sulphuric
acid, H2SO4,
is the acidic proton donor and H2O is the proton
accepting base.
-
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 and the
original acid are known as a conjugate pair and are
related by 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.
-
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
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 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
can accept.
-
monobasic
acids e.g.
-
dibasic acids
e.g.
-
sulphuric H2SO4,
ethanedioic (COOH)2, and the three isomeric
-
benzene-x,y-dicarboxylic acids
(x,y = 1,1 1,2 and 1,3) 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.
-
5.1.4: Examples of
water
insoluble bases giving acid-base neutralization reactions.
-
5.1.5: Examples of
two solids reacting together in an acid-base reaction.
-
-

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