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Part 6.
The Chemistry of Carboxylic Acids and their Derivatives
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Brown's Chemistry Advanced Level Pre-University Chemistry Revision Study
Notes for UK KS5 A/AS GCE IB advanced level organic chemistry students US
K12 grade 11 grade 12 organic chemistry theory of structure and strength of
weal organic carboxylic acids salt formation carboxylates
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Part 6.4
The weakly acidic nature and general
reactions of carboxylic acids acting as acids (carboxylate salts) and factors affecting the strength
of carboxylic acids
Sub-index for this page
6.4.1
Carboxylic acids are
weak acids - ionisation and the structure of the carboxylate ion
6.4.2
Structure of carboxylic
acids and the
ionisation constant Ka (dissociation constant)
6.4.3
Observations - comparing the reactivity
of weak carboxylic acids and strong mineral acids
6.4.4
Reactions of carboxylic acids with
metals
6.4.5
Reactions of carboxylic acids with
oxides & hydroxides (soluble/insoluble bases)
6.4.6
Reactions of carboxylic acids with
hydrogencarbonates and carbonates
6.4.7
Reaction of carboxylic acids with
ammonia
6.4.1 Carboxylic acids
are weak acids - ionisation and the structure of the carboxylate ion
The general structure of a
carboxylic acid and the corresponding carboxylate ion.
and
The displayed general formula for a monocarboxylic acid. R = H, alkyl (e.g.
CH3CH2) or aryl (e.g. C6H5) and
the corresponding carboxylate anion.
Note (i) the shorter C=O double bond
of 0.123 nm in the unionised acid, (ii) the longer C-O single bond of 0.133 nm
in the unionised acid and (iii) the
equal length C-O bonds of intermediate bond length 0.128 nm in the delocalised bond system of the
carboxylate ion from the ionised acid.
The 'traditional' displayed formula for the carboxylate anion (this negative
ion is the conjugate base of the acid).
and
are known as resonance hybrids.
This is a way of describing bonding in
structures that can be better described by looking at multiple possible
contributing structures - imagine an oscillation between them (blue arrows
show the theoretical electron shifts).
The result is that the theoretical pi bond pair of
the C=O group is actually a delocalised system across the two C-O bonds.
At a higher level you can think of it as sp2 hybridisation
for the trigonal planar -C< bond arrangement and the delocalised '0.5'
bond of the O-C-O formed from the overlap of the pz atomic
orbitals.
and
The true structure of the carboxylate ion in which both C-O bonds are both equal
length and both equivalent to a bond order of 1.5.
The R-C-O and O-C-O bond angles
are 120o - a trigonal planar sigma bond arrangement, plus the
delocalised pi bond system in which two electrons reside in the pi
orbitals above and below the plane of the O-C-O bond system. Explanation:
bond order = the
equivalent pairs of bonding electrons in a covalent bond
e.g. single C-H, O-H and C-O bonds have a bond order of 1, C=O
usually has a bond order of 2 (as in the acid), but here it is 1.5 in
the carboxylate ion. Note again the
correspond bond lengths of C=O 0.123 nm (bond order 2 in acid), C-O bond of 0.133 nm
(bond order 1 in acid) length 0.128 nm (bond order 1.5) in the delocalised bond system of the
carboxylate ion.
The result of the possible 'resonance hybrids' is
that the theoretical pi bond pair of electrons from the C=O group
actually form a delocalised system across the two C-O bonds in the
carboxylate ion.
This means both of the C-O
bonds is identical and equivalent to 0.5 of a pi bond plus 1.0 of a sigma bond,
hence the bonder order of 1.5 in the carboxylate ion (a sort of half-way
bond length).
and
This shows the ionisation of a carboxylic acid and the usual 'shorthand'
formula in common use.
Expressing the
ionisation, both chemically and mathematically
In aqueous solution, carboxylic acids are weak acids,
that is only ionised or dissociated to a small extent.
RCOOH(aq)
+ H2O(l)
H3O+(aq) + RCOO-(aq)
RCOOH(aq)
H+(aq) + RCOO-(aq)
(useful when doing pH and Ka calculations)
The ionisation constant (dissociation constant, acidity
constant), that is
the equilibrium constant (Ka)
for the above equilibrium, is given by the expression:
|
Ka =
|
[H+(aq)]
[RCOO–(aq)]
|
|
––––––––––––––––––
mol dm-3 |
|
[RCOOH(aq)] |
Like pH, the range of Ka is so
great, it is often quoted on the logarithmic scale where
pKa
= -log10(Ka/mol dm-3)
The stronger the
acid, the greater the Ka and the lower the pKa values.
In the data table next in section 6.4.2
I've quoted both values.
See equilibrium calculations section 5c
for calculations involving pH, Ka
and pKa
Dissociation of dicarboxylic
and tricarboxylic acids in aqueous solution
The ionization of a dicarboxylic acid requires two
equations e.g. in 'shorthand' for propanedioic acid
(i)
HOOC-CH2-COOH(aq)
HOOC-CH2-COO-(aq) + H+(aq)
 (ii)
HOOC-CH2-COO-(aq)
-OOC-CH2-COO-(aq) + H+(aq)
For a tricarboxylic acid like citric acid, there will be
three ionisation equations.
HOOCCH2C(OH)(COOH)CH2COOH(aq)
===> HOOCCH2C(OH)(COO-)CH2COOH(aq)
+ H+(aq)
HOOCCH2C(OH)(COO-)CH2COOH(aq)
===> HOOCCH2C(OH)(COO-)CH2COO-(aq)
+ H+(aq)
HOOCCH2C(OH)(COO-)CH2COO-(aq)
===> -OOCCH2C(OH)(COO-)CH2COO-(aq)
+ H+(aq)
These acids are referred to as dibasic or
tribasic acids and also diprotic or triprotic acids.
Why are carboxylic acids stronger
acids than alcohols?
CH3COOH pKa = 4.76 (weak
acid) and CH3CH2OH pKa ~16
(extremely weak acid).
The are
two principal reasons (i) The
carbonyl C=O group pulls the electron clouds away from the O-H group
making the hydrogen atom more δ+, making the H+ ion easier to remove by
a lone pair donating base e.g. water :OH2.
Look at the electron shift above
towards the more electronegative oxygen δ+C=Oδ
(the top part of the diagram).
(ii) The delocalisation of the
negative charge between the two oxygen atoms of the carboxylate ion,
lowers the potential energy of the system, giving extra stability to the
anion (diagram above).
You can think of each oxygen as
carrying half a single negative charge.
The negative charge on the anion from an alcohol,
an alkoxide ion e.g.
CH3CH2O- from CH3CH2OH,
cannot be delocalised, so no extra stability.
The alkoxide ion is a
stronger conjugate base than the carboxylate ion.
In fact in water, e.g. the
ethoxide ion, rapidly hydrolyses yielding the original alcohol,
ethanol, and the hydroxide ion, so forming an alkaline solution.
CH3CH2O-(aq)
+ H2O(l) ===> CH3CH2OH(aq)
+ OH-(aq)
TOP OF PAGE and sub-index
6.4.2 Structure of
carboxylic acids and the value of the
ionisation constant Ka (dissociation constant)
Although they are
generally weak acids, the range of ionisation is quite wide and greatly
affected by substituent groups in the 'alky' or 'aryl' hydrocarbon part
of the molecule, so, like with pH, a logarithmic (base 10) scale is used ...
... for examples see the relative strengths of carboxylic
acids in the data table below.
If the R group contains electronegative, electron
cloud withdrawing atoms, the electron shift affects the δ-O-Hδ+
bond (-I induction effect) and promotes the donation of a proton to a base
(left of bottom diagram).
On the other hand, other groups like alkyl groups, can
have a +I inductive electron shift effect and make the proton less available
to a base (right of bottom diagram).
Reminders of mathematical
connections:
Ka =
[H+(aq)] [RCOO-(aq)] /
[RCOOH(aq)]
(R = H, alkyl, aryl, the
latter may be substituted with other atoms)
pKa = - log10(Ka/moldm-3)
and Ka = 10-pKa
Like the pH scale, a base 10
logarithmic scale means each unit
of pH is equal to a factor of 10 in the
hydrogen ion concentration [H+(aq)] and
each unit of pKa
represents a factor of 10 in the value of the ionization constant
(dissociation constant) in aqueous solution.
Note: For dicarboxylic acids or
tricarboxylic acids, the ionisation decreases from the 1st, 2nd
ionisation etc.
so that:
Ka1 > Ka2
etc. and pKa1
< pKa2 etc.
| Linear monocarboxylic acids |
Abbreviated formula |
pKa |
Ka/mol
dm-3 |
Comments |
| Methanoic acid |
HCOOH |
3.75 |
1.78 x 10-4 |
|
| Ethanoic acid |
CH3COOH |
4.76 |
1.74 x 10-5 |
Vinegar has a pH of ~3 |
| Propanoic acid |
CH3CH2COOH |
4.87 |
1.35 x 10-5 |
After ethanoic, most have a
similar pKa |
| Butanoic acid |
CH3(CH2)2COOH |
4.82 |
1.51 x 10-5 |
|
| Pentanoic acid |
CH3(CH2)3COOH |
4.86 |
1.38 x 10-5 |
|
| Chloroethanoic acid |
ClCH2COOH |
2.86 |
1.38 x 10-3 |
Three substituted
ethanoic acids - these show the increasing effect of one to three
electronegative atoms on the extent of dissociation into ions.
You get similar effects with F, Br, I, OH and NO2
substituent groups. |
| Dichloroethanoic acid |
Cl2CHCOOH |
1.29 |
5.13 x 10-2 |
| Trichloroethanoic acid |
Cl3CCOOH |
0.65 |
2.24 x 10-1 |
| |
|
|
|
|
| Dicarboxylic
aliphatic acids |
Abbreviated formula |
pKa |
Ka/mol
dm-3 |
Comments |
| Ethanedioic
acid |
HOOCCOOH |
pKa1 = 1.23 pKa2 =
4.28 |
Ka1
= 5.89 x 10-2 Ka2
= 5.25 x 10-5 |
Two ionisations
possible to yield hydrated protons H3O+. Note the order
Ka1 > Ka2
and
pKa1 < pKa2
for dibasic acids. |
| Propanedioic
acid |
HOOCCH2COOH |
pKa1 = 2.83 pKa2 =
5.69 |
Ka1
= 1.48 x 10-3 Ka2
= 2.04 x 10-6 |
|
| Butanedioic acid |
HOOCCH2CH2COOH |
Ka1
= 4.22 Ka2 = 5.64 |
Ka1
= 6.03 x 10-5 Ka2
= 2.29 x 10-6 |
|
| |
|
|
|
|
| Aromatic acids |
Abbreviated formula |
pKa |
Ka/mol
dm-3 |
Comments |
| Benzoic acid |
C6H5COOH |
4.20 |
6.31 x 10-5 |
|
| 2-chlorobenzoic
acid |
ClC6H4COOH |
2.94 |
1.48 x 10-3 |
The
presence of the electronegative chlorine atom, increases the
strength of the acid compared to benzene. |
| 3-chlorobenzoic acid |
ClC6H4COOH |
3.83 |
1.48 x 10-4 |
| 4-chlorobenzoic acid |
ClC6H4COOH |
3.99 |
1.02 x 10-4 |
| 2-nitrobenzoic acid |
O2NC6H4COOH |
2.17 |
6.76 x 10-3 |
The
presence of the electronegative nitro group -NO2,
increases the strength of the acid compared to benzene, more so than
the a single chlorine atom. |
| 3-nitrobenzoic acid |
O2NC6H4COOH |
3.45 |
3.55 x 10-4 |
| 4-nitrobenzoic acid |
O2NC6H4COOH |
3.43 |
3.72 x 10-4 |
| benzene-1,2-dicarboxylic acid |
C6H4(COOH)2 |
pKa1 = 2.98 pKa2 =
5.41 |
Ka1
= 1.05 x 10-3 Ka2
= 3.89 x 10-6 |
Two ionisations
possible to yield hydrated protons H3O+. Note the order
Ka1 > Ka2
and
pKa1 < pKa2
for dibasic acids. |
| benzene-1,3-dicarboxylic acid |
C6H4(COOH)2 |
pKa1 = 3.46 pKa2 =
4.60 |
Ka1
= 3.47 x 10-4 Ka2
= 2.51 x 10-5 |
|
| benzene-1,4-dicarboxylic acid |
C6H4(COOH)2 |
pKa1 = 3.51 pKa2 =
4.82 |
Ka1
= 3.09 x 10-4 Ka2
= 1.51 x 10-5 |
|
| |
|
|
|
|
Structure notes:
2-, 3- and
4-chlorobenzoic acid
2-, 3-
and
4-nitrobenzoic acid
structures of the aromatic dicarboxylic acids
TOP OF PAGE and sub-index
6.4.3 Observations - comparing the reactivity
of weak carboxylic acids and strong mineral acids
Carboxylic acids are weak acids
Typically weak acid
solutions have a pH of around 2 to 6 (yellow–orange–pink with universal indicator),
which is somewhat higher than strong acid solutions with a pH of 0 to 1
(always <2).
They are called
weak acids because only a
few % of the molecules in aqueous ionise to release protons (hydrogen
ions, H+).
It is the generation of hydrogen ions that
makes aqueous solutions of carboxylic acids acidic.
e.g. for ethanoic acid (vinegar) around 98%
remains unionised i.e. as the original neutral molecule and only ~2%
form ethanoate ions and hydrogen ions..
CH3COOH(aq)
CH3COO–(aq)
+ H+(aq)
Propanoic acid and butanoic acid are equally weak
carboxylic acids.
CH3CH2COOH(aq)
CH3CH2COO–(aq)
+ H+(aq)
CH3CH2CH2COOH(aq)
CH3CH2CH2COO–(aq)
+ H+(aq)
This is a reversible reaction
with only 2% of the weak acid ionised on the right–hand side of the equilibrium.
At similar aqueous solution concentrations, strong
mineral acids like hydrochloric acid, sulfuric acid and nitric acid have
a low pH of -1, 0 or 1, because they are fully ionised
(for H2SO4 this only applies to the 1st ionisation).
So when you dissolve gaseous hydrogen chloride or
liquid sulfuric acid in water, the ionisation is ~100%
e.g.
HCl(g)
+ aq
===> H+(aq) + Cl-(aq)
(the Ka ionisation constant for
hydrochloric acid is 108 mol/dm3, a sharp
contrast with the Ka values of weak organic carboxylic acids)
H2SO4(l)
+ aq
===> H+(aq) + HSO4-(aq)
(the Ka for the 1st ionisation of
sulfuric acid is ~1016 mol/dm3)
and similarly for nitric acid
HNO3(g)
+ aq
===> H+(aq) + NO3-(aq)
(the Ka ionisation constant for
nitric acid is 1015 mol/dm3)
Even so, weak acids like ethanoic acid will turn blue litmus pink
and universal indicator gives a red colour in aqueous solution and
usually liberate carbon dioxide from carbonates (and CO2 test) -
simple tests for an acidic substance.
Carboxylic acids react with (i) metals, (ii)
oxides and hydroxides, (iii) hydrogencarbonates and carbonates and (iv)
ammonia.
(see sections
6.4.4,
6.4.5, 6.4.6 and
6.4.7).
Comparing
the pH and rates of reaction of weak and strong acids of equal molarity
e.g. 1.0 mol dm-3 solutions of ethanoic
acid and hydrochloric acid.
For equimolar solutions, the solution of the
strong acid will have a much great concentration of hydrogen ions, so
its pH will be much lower - use an accurately buffer calibrated pH meter
The
pH of solutions of equal
concentration e.g. of molarity 1.0 mol/dm3
The pH of a strong acid might be pH 0-1
(hydrochloric, sulfuric or nitric acids).
The pH of a weak acid might be typically
pH 3-6 (vinegar ~pH 3, carbonic acid pH 4-5).
The
rate of reaction with metals.
(1 molar means 1.0 mol/dm3, sometimes written as 1M for
shorthand))
If you put magnesium ribbon into 1 molar
solutions of hydrochloric acid (strong, high % ionisation so high H+(aq)
concentration) and 1 molar solution of ethanoic acid (weak, low percentage ionization so much lower H+(aq)
concentration), you can
see the difference in the fast and slow 'fizzing' rates!
You can repeat the experiment
using calcium carbonate (limestone granules) instead of magnesium ribbon and
collect carbon dioxide gas.
You can do simple
rate of reaction experiments
comparing how
fast the gas is evolved from the reaction mixture.
The above links takes you to a
page where the experimental procedures are described, little point in
repeating them here.
The basic experimental procedure
is shown in the diagram above and a graph of typical results below.
The above graph shows the sort
of results you might expect by adding the same masses of magnesium ribbon or
calcium carbonate granules to the same volume of ethanoic acid, CH3COOH,
or hydrochloric acid, HCl, of equal concentration e.g. both acids
with a concentration of 1.0 mol/dm3.
Remember that its
the hydrogen ion, the H+(aq) ion (H3O+), is the active chemical species in acid
solutions NOT a 'HCl' or a 'H2SO4' or a 'CH3COOH'
molecule.
Electrolysis observations
 Since stronger/weak
acid solutions (or alkalis) contain more/less hydrogen ions, they are better/poorer conductors of electricity.
e.g. If you carry out
electrolysis experiments with the same molarity solutions of hydrochloric
acid and ethanoic acid, you get a greater
rate of hydrogen collected at the negative cathode from the hydrochloric acid
compared to the ethanoic acid.
2H+(aq) +
2e- ===> H2(g)
The apparatus is the same as for
electrolysing water acidified with sulfuric acid.
You must use solutions
of the same concentration and electrolysed them for the same length time at
the same voltage (potential difference, p.d.) before
measuring the gas volumes of hydrogen formed. (Electrolysis
methods).
From the strong acid solution, you should
get a greater volume of hydrogen in the same time.
The greater concentration of
ions in the strong acid solution reduces the electrical resistance and more
current flows via the greater number of ions present to carry it, hence more
hydrogen ions are reduced at the cathode to form hydrogen.
For a more detailed discussion of these points see
The theory
of acids and bases AND make sure you know the difference
between 'strength' and concentration !!!
'Strength' is about how ionised is the acid
i.e. how big is Ka (or Kb for bases)
'Concentration' is about how many particles
of the solute per unit volume e.g. molarity in mol dm-3.
Despite being a weak acid, carboxylic
acids like ethanoic acid behave like any other acid and react with
metals, alkalis and carbonate to form salts and fizzing here and there!
TOP OF PAGE and sub-index
6.4.4 The reaction of carboxylic acids
with metals
The salt names depends on the name of the acid,
but the end of the name is ... oate
(carboxylate salt)
So aqueous solutions
of methanoic acid form methanoate salts, ethanoic acid gives ethanoate salts, propanoic
acid gives propanoate salts and butanoic acid gives butanoate
salts on neutralisation etc.
The salts can be crystallised from the
solution by evaporation.
Metals dissolve
in aqueous carboxylic acid solutions to form a salts and
hydrogen e.g.
(i) ethanoic acid + magnesium ==>
magnesium ethanoate + hydrogen
2CH3COOH(aq) + Mg(s) ===> (CH3COO)2Mg(aq) + H2(g)
2CH3COOH(aq) + Mg(s) ===>
2CH3COO-(aq) + Mg2+(aq) + H2(g)
(ii) butanoic acid + zinc ====>
zinc butanoate + hydrogen
2CH3CH2CH2COOH(aq) + Zn(s) ===> (CH3CH2CH2COO)2Zn(aq) + H2(g)
2CH3CH2CH2COOH(aq) +
Zn(s) ===> 2CH3CH2CH2COO-(aq)
+ Zn2+(aq) + H2 (g)
-
Not on the syllabus, but an interesting
tragic story of 'old acetic acid' and lack of appreciation of
chemical hazards.
Ethanoic acid very slowly reacts with lead to
form lead(II) ethanoate (old name lead acetate), once called 'sugar
of lead'
2CH3COOH(aq) + Pb(s) ===> (CH3COO)2Pb(aq) + H2(g)
The salt formed was called 'sugar of lead' because
it had a sweet taste, but, ironically, its a deadly nerve toxin !
Cider makers in the past had dipped
rods of lead into cider to neutralise any acetic acid that had
formed and sweeten the beverage.
Unfortunately, lead is one of many heavy
metals and that are highly toxic and lead compounds affect the brain
and nervous systems and can be fatal.
Cases of lead poisoning have occurred through
millennia, including the Romans, by using lead piping, lead pots in food
preparation or concentrating liquids in cooking.
So, any cider left over that goes sour,
dispose of it or, even better, let it turn completely into cider vinegar for the
kitchen!
TOP OF PAGE and sub-index
6.4.5 Reaction of
carboxylic acids with oxides & hydroxides (soluble or insoluble bases)
The salts formed in these
reactions are known as 'carboxylates', but the salt name ends in
...oate (derived from the name of the carboxylic acid)
(a) Alkalis (soluble bases)
are neutralised by carboxylic acids to form a carboxylic acid salt
and water e.g.
The ionic
equation is: RCOOH(aq)
+ OH-(aq) ===> RCOO-(aq)
+ H2O(l)
R = H, alkyl (e.g. CH3CH2)
or aryl (e.g. C6H5)
(i) ethanoic acid + sodium hydroxide
===> sodium ethanoate + water
CH3COOH(aq) + NaOH(aq) ===> CH3COONa(aq) + H2O(l)
CH3COOH(aq) + OH-(aq) ===> CH3COO-(aq)
+ H2O(l)
The pH of the neutralised solution is ~9 (the
salt of a weak acid and a strong base).
(ii) propanoic acid + potassium hydroxide
===> potassium propanoate + water
CH3CH2COOH(aq) + KOH(aq) ===> CH3CH2COOK(aq)
+ H2O(l)
(iii) butanoic acid + sodium hydroxide
===> sodium butanoate + water
CH3CH2CH2COOH(aq) + NaOH(aq) ===> CH3CH2CH2COONa(aq)
+ H2O(l)
(iv) benzoic acid + sodium hydroxide
===> sodium benzoate + water
(aq) + NaOH(aq) ===>
(aq) + H2O(l)
(v) Making soluble aspirin
from:
(aq)
+ NaOH(aq) ===>
(aq)
+ H2O(l)
If you gradually add alkali to the carboxylic
acid, you get a characteristic curve of pH changes.
Curve
(2) +
(4):
Adding a strong base to weak acid, end point
(i3), at pH ~9.
pH change at
end–point reasonable sharp e.g. you can titrate weak organic
acids like ethanoic acid with standard sodium hydroxide solution
NaOH is a strong soluble base.
Suitable
indicators:
phenolphthalein (pKind 9.3,
range 8.3–10.0), end-point is the first permanent pink.
Thymol blue (base, pKind 8.9,
range 8.0–9.6)
This neutralisation reaction can be used for
the quantitative analysis of carboxylic acids.
You can even analyse sparingly soluble carboxylic
acids by dissolving it in aqueous ethanol solvent.
For more details see:
Volumetric titration procedures and
calculations for acid-alkali titrations
Advanced level
acid-base titration calculation questions
Advanced level theory of pH curves
and indicator choice for acid - alkali titrations
For a dibasic acid, there are two stages of
neutralisation and two possible salts can be crystallised from solution,
depending on the ratio of alkali added to the carboxylic acid e.g. for
...
(iv) Ethanedioic acid
HOOC-COOH(aq) + NaOH(aq) ===>
HOOC-COO-Na+(aq) + H2O(l)
HCOO-COO-Na+(aq) + NaOH(aq)
===> Na+-OOC-COO-Na+(aq)
+ H2O(l)
Overall
HOOC-COOH(aq) + 2NaOH(aq) ===>
Na+-OOC-COO-Na+(aq) +
2H2O(l)
(v) Butanedioic acid
HOOCCH2CH2COOH(aq) + NaOH(aq)
===> HOOCCH2CH2COO-Na+(aq)
+ H2O(l)
HCOOCH2CH2COO-Na+(aq)
+ NaOH(aq) ===> Na+-OOCCH2CH2COO-Na+(aq)
+ H2O(l)
Overall
HOOCCH2CH2COOH(aq) + 2NaOH(aq)
===> Na+-OOCCH2CH2COO-Na+(aq)
+ 2H2O(l)
Dibasic carboxylic acids (dicarboxylic acids) can
also be titrated with standard sodium hydroxide solution.
The more complicated pH curve, adding alkali to a dicarboxylic acid
There are two
inflexion points on the pH curve corresponding to the half and
full neutralisation of the dibasic/diprotic acid.
The equations for the two step
neutralisation of ethanedioic acid are given below, including the ionic
equations.
HOOC–COOH(aq)
+ NaOH(aq) ==> HCOO–COO–Na+(aq)
+ H2O(l)
ionically:
HOOC–COOH(aq) + OH–(aq) ==>
HCOO–COO–(aq) + H2O(l)
HCOO–COO–Na+(aq)
+ NaOH(aq) ==> Na+–OOC–COO–Na+(aq)
+ H2O(l)
ionically:
HCOO–COO–(aq) + OH–(aq)
==> –OOC–COO–(aq) + H2O(l)
To detect the
2nd end–point, and hence the acid quantitatively, phenolphthalein
indicator (pKind 9.3,
range 8.3–10.0) is used, since it is essentially a weak
acid–strong base titration.
Other acids
like propanedioic acid (malonic acid) and butanedioic acid (succinic
acid)
behave, and be titrated, in the same way.
You
can titrate citric acid in fruit juice with phenolphthalein indicator.
The end-point is the first permanent pink.
The overall neutralisation equation is:
HOOCCH2C(OH)(COOH)CH2COOH(aq)
+ 3OH-(aq) ===> -OOCCH2C(OH)(COO-)CH2COO-(aq)
+ 3H2O(l)
For more details see:
Volumetric titration procedures and
calculations for acid-alkali titrations
Advanced level
acid-base titration calculation questions
Advanced level theory of pH curves
and indicator choice for acid - alkali titrations
(b) Insoluble bases dissolve
in carboxylic acids to give a
salt and water e.g.
(i) zinc oxide + ethanoic acid ====>
zinc ethanoate + water
2CH3COOH(aq) + ZnO(s) ===> (CH3COO)2Zn(aq) + H2O(l)
2CH3COOH(aq) + ZnO(s) ===>
2CH3COO-(aq) + Zn2+(aq) + H2O(l)
(ii) ethanoic acid + calcium hydroxide ====>
calcium ethanoate + water
2CH3COOH(aq) + Ca(OH)2(s) ===> (CH3COO)2Ca(aq) + 2H2O(l)
2CH3COOH(aq)
+ Ca(OH)2(s) ===> 2CH3COO-(aq)
+ Ca2+(aq) + 2H2O(l)
(iii) propanoic acid + magnesium hydroxide ====>
magnesium propanoate + water
2CH3CH2COOH(aq) + Mg(OH)2(s) ===> (CH3CH2COO)2Mg(aq) + 2H2O(l)
2CH3CH2COOH(aq)
+ Mg(OH)2(s) ===> 2CH3CH2COO-(aq)
+ Mg2+(aq) + 2H2O(l)
(iv) butanoic acid + magnesium hydroxide ====> magnesium
butanoate + water
2CH3CH2CH2COOH(aq) + Mg(OH)2(s) ===>
(CH3CH2CH2COO)2Mg(aq) + 2H2O(l)
2CH3CH2CH2COOH(aq)
+ Mg(OH)2(s) ===> 2CH3CH2CH2COO-(aq)
+ Mg2+(aq) + 2H2O(l)
The colourless salts can be crystallised out by
carefully evaporating most of the water off.
TOP OF PAGE and sub-index
6.4.6 The
reaction of carboxylic acids with carbonates and hydrogencarbonates
The salts formed in these
reactions are known as 'carboxylates', but the salt name ends in
...oate (derived from the name of the carboxylic acid)
Carbonate and hydrogencarbonate
bases
produce a carboxylic acid salt, water and carbon dioxide gas e.g.
(i) ethanoic acid + sodium hydrogen
carbonate ==> sodium ethanoate + water + carbon dioxide
CH3COOH(aq) + NaHCO3(s)
====> CH3COONa(aq) + H2O(l) + CO2(g)
CH3COOH(aq) + NaHCO3(s)
====> CH3COO-(aq) + Na+(aq) + H2O(l) + CO2(g)
(ii) ethanoic acid + sodium carbonate
====> sodium ethanoate + water + carbon dioxide
2CH3COOH(aq) +
Na2CO3(s) ====> 2CH3COONa(aq)
+ H2O(l) + CO2(g)
2CH3COOH(aq) + Na2CO3(s)
====> 2CH3COO-(aq) +
2Na+(aq) + H2O(l) + CO2(g)
Sodium carbonate is quite soluble in water, but
I've just assumed your adding the acid solution to the solid.
(iii) propanoic acid + potassium carbonate
====> potassium propanoate + water + carbon dioxide
2CH3CH2COOH(aq)
+ K2CO3(s) ====> 2CH3CH2COOK(aq)
+ H2O(l) + CO2(g)
2CH3CH2COOH(aq)
+ K2CO3(s)
====> 2CH3CH2COO-(aq)
+ 2K+(aq) + H2O(l) + CO2(g)
(iv) ethanoic acid + magnesium carbonate
==> magnesium ethanoate + water + carbon dioxide
2CH3COOH(aq) + MgCO3(s) ====> (CH3COO)2Mg(aq) + H2O(l) + CO2(g)
2CH3COOH(aq) +
MgCO3(s) ===>
2CH3COO-(aq) + Mg2+(aq) + H2O(l) + CO2(g)
(v) butanoic acid + calcium carbonate
==> calcium butanoate + water + carbon dioxide
2CH3CH2CH2COOH(aq) + CaCO3(s) ====> (CH3CH2CH2COO)2Ca(aq) + H2O(l) + CO2(g)
2CH3CH2CH2COOH(aq) + CaCO3(s) ===>
2CH3CH2CH2COO-(aq) + Ca2+(aq) + H2O(l) + CO2(g)
The colourless salts can be crystallised out by
carefully evaporating most of the water off.
TOP OF PAGE and sub-index
6.4.7 The reaction
of carboxylic acids with aqueous ammonia
The salts formed in these
reactions are known as 'carboxylates', but the salt name ends in
...oate (derived from the name of the carboxylic acid)
Aqueous ammonia solution
forms ammonium salts e.g.
methanoic acid + ammonia ====>
ammonium methanoate
HCOOH(aq) + NH3(aq) ===> HCOONH4(aq)
ethanoic acid + ammonia ==>
ammonium ethanoate
CH3COOH(aq) + NH3(aq) ===> CH3COONH4(aq)
pH of the neutralised salt solution is ~7
Here we are adding a weak base to a weak acid
(or vice versa!)
The pH curve for adding a weak acid to a weak
base
Curves (2)
+ (3)
adding weak acid to weak base
The pH curve for adding a weak base to a weak
acid
Curves (2)
+ (3)
adding weak base to weak acid
In both cases, the point of inflexion in the pH curve is not
sufficient to give a sharp end-point - not a viable quantitative
titration.
For details see
Advanced level theory of pH curves
and indicator choice for acid - alkali titrations
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