Acids and Bases
Basics
1.
2.
3.
Acids taste sour
Acids turn litmus red
ACID + REACTIVE METAL  SALT + HYDROGEN
2HCl + Mg  MgCl2 + H2
4.
ACID + BASE  SALT + WATER
H2SO4 + 2NaOH  Na2SO4 + 2H2O
5.
ACID + METAL CARBONATES
 SALT + WATER+ CARBON DIOXIDE
2HNO3 + CaCO3  Ca(NO3)2 + H2O + CO2
Arrhenius
ACID: substance which produces hydrated oxonium (hydrogen) ions in aqueous
solution.
HCl + H2O  H3O + Cl
+
-
Monoprotic: HCl, HNO3, CH3COOH
Diprotic:
H2SO4, H2CO3
Triprotic:
H3PO4
ALKALIS: substance which produces hydrated hydroxide ions in aqueous
solution.
NaOH(s)  Na+(aq) + OH-(aq)
Strong acid or strong base: These are highly ionised in solution
(high Ka and Kb value)
e.g. hydrochloric acid, sulphuric acid, nitric acid
all soluble metal hydroxides NaOH, KOH
Weak acid or weak base: These are slightly ionised in solution
(low Ka and Kb value)
e.g. all organic acids, sulphurous acid
base : ammonia.
BRONSTED – LOWRY ( learn these as definitions.)
ACID IS A PROTON (H+) DONOR
BASE IS A PROTON (H+) ACCEPTOR
HCl + NH3  NH4+ + Clacid
base
:
A reaction in which a proton (H+) is transferred is called PROTOLYSIS
This reaction is reversible so: NH4+ + Cl-  HCl + NH3
acid base
Every base has a conjugate acid and every acid a conjugate base.
HCl + NH3 ⇄ NH4+ + Clacid
base
acid
base
Note: the conjugate acid always has an
extra H+ compared to its conjugate
base.
e.g.
Cl- is the conjugate base of the
acid HCl
NH4+ is the conjugate acid of
the base NH3
the conjugate acid of OH- is H2O
the conjugate base of OH- is O2-
AMPHOLYTES
Some substances can act as acids or base and these are known as ampholytes or
ampholytic.
Water is a common example:
HCl + H2O  H3O+ + Clbase
NH3 + H2O ⇄ NH4+ + OHacid
Other commonly asked examples are the HCO3- and HSO4- ions as they can donate or
gain a proton (H+).
1.
a) Identify the conjugate acid-base pairs in the following reversible reactions:
HSO4- + OH- ⇄ SO42- + H2O
CH3COOH + KOH ⇄ CH3COOK + H2O
b) The conjugate acid of H2O is ...........................
c) The conjugate base of H2PO4- is ......................................
d) Why can CO2 not be an acid? ..................................................................
e) What is the conjugate acid of SO42-? ..................................
Ionic Product of a water, Kw
Water undergoes autoprotolysis. This can be written as either
H2O ⇄ H+ + OH-
or
H2O(l) + H2O(l) ⇄ H3O+(aq) + OH-(aq)
In any aqueous solution at 25 oC
Kw = 10-14 = [H+].[ OH-]
In a neutral solution [H+] = [ OH-] = 10-7 mol.dm-3
This means that ALL aqueous solutions have both H+ ions and OH- ions in them.
If [H+] > [ OH-] then the solution is acidic.
If [H3O+]<[ OH-] then the solution is alkaline.
If [H3O+] = [ OH-] then the solution is neutral.
e.g.1
What is the concentration of hydroxide ions in an aqueous solution of a salt
where the concentration of hydrogen ions is 3 x 10-4 mol.dm-3 at 25 oC?
e.g.2
What is the concentration of hydrogen ions in a 0,02 mol.dm-3 solution of
calcium hydroxide (Ca(OH)2) at 25oC ?
Ca(OH)2  Ca2+ + 2OH-
e.g.3
If a solution of sulphuric acid contains OH- ion with a concentration
2,4 x 10-14 mol.dm-3, what was the concentration of the acid. [Hint: remember
sulphuric acid is a diprotic acid].
pH
This indicates the acidity or alkalinity of a solution.
It is a logarithmic measure of the concentration of hydrogen (oxonium) ions.
pH = -log[H+]
pH < 7 is acidic, >7 alkaline, 7 neutral (at 25oC)
e.g
What is the pH of a 0,01 mol.dm-3 of HNO3?
HNO3  H+ + NO3pH = -log[H+] = -log(0,01) = 2
Note: we only work out pH for acids we assume of fully ionised.
i.e. strong ones
e.g.1
What is the pH of a solution of 0,004 mol.dm-3 hydrochloric acid?
e.g.2
What is the concentration of a H2SO4 solution with pH 0,22? [Hint: diprotic acid]
e.g.3
What is the pH of a solution of nitric acid if 16,4 g of it is mixed with water to make
400 cm3 of solution?
e.g.4
If all the following solutions are the same concentration, place in order of pH. Lowest
pH first:
ethanoic acid; sodium hydroxide; ammonium hydroxide; hydrochloric acid; sulphuric acid
TITRATIONS
A titration is used to determine the concentration of a solution using a standard
solution.
A standard solution is one where the concentration is know accurately
and remains constant for a long time.
A standard solution is made in a volumetric flask, usually by dissolving the required
mass of solid.
n = m/M
c = n/V
(V in dm-3)
c = m .
M.V
A pipette is used to measure out one of the solutions and this is
placed in a conical flask. An indicator is then added to this. The
other solution is then run from a burette into this mixture, with
stirring, until one drop causes the end-point to be reached and the
indicator changes to the required colour. This is usually repeated a
number of times and an average value taken. The unknown
concentration is then calculated.
e.g.
25,0 cm3 of a standard solution of 0,05 mol.dm-3 calcium hydroxide was placed
in a conical flask. Hydrochloric acid was run into this and the end-point was
reached after 14,8 cm3 was added. Calculate the concentration of the acid.
2HCl + Ca(OH)2  CaCl2 + 2H2O
Method 1
(questions are often asked so that
you have to use this method)
n(Ca(OH)2) = c.V = 0,05 x 0,025
= 0.00125mol
n(HCl) = 0,00125 x 2 (ratio)
= 0,0025mol
c(HCl) = n/V = 0,0025/0,0148
= 0,169 mol.dm-3
Method 2
ca.Va = a
cb.Vb
b
INDICATORS
An indicator is a dye which changes colour with a certain change in the
pH of its surroundings. It can be regarded as a weak acid, HIn, where the
In- is the anion of the dye.
The HIn and I- have different colours. When the concentration of one is
10 times the concentration of the other then its colour will be seen. This
requires a pH change of 2 so the colour normally changes over a range of
about 2 (see table below)
HIn ⇄
colour 1
H+ + Incolour 2
It is a reversible reaction.
When acid is added the [H+] increases forcing the equilibrium position to
the left and colour 1 dominates.
When as alkali is added [H+] decreases as the H+ reacts with OH- to form
water, and so equilibrium shifts to the right and colour 2 dominates.
CHOOSING AN INDICATOR
Different indicators change colour over different ranges. One chooses
ones indicator according to the type of titration (strong acid/ weak base
etc). You must learn the colours and types of titration for the 3 indicators
below:
indicator
methyl red
bromothymol
blue
phenolphthalein
pH range
4,2 – 6,2
6,0 – 7,6
in acid
red
yellow
in alkali
yellow
blue
titration type
strong acid/weak base
strong acid/ strong base
8,3 – 10,0
colourless
red
weak acid/ strong base
END POINT
The end-point of an acid-base titration results when the number of moles
of acid added to base, or visa-versa, is in the same ratio as that in the
chemical equation. This does not necessarily mean that the solution is
neutral. It will be in a strong acid/strong base titration but not in the other
two examples in the table.
At the end-point a drop of acid or alkali will cause a large swing in pH
and this causes the colour of the solution to change (see titration curves).
1.
2,4 g of NaOH was weighed out and dissolved in some water. This was then made up to
100 cm3 of volume with further water. 25,0 cm3 of this NaOH solution was then titrated
against oxalic acid and it took 13,8 cm3 of the acid solution to effect the correct colour
change in the indicator.
HOOCCOOH + 2NaOH  (OOCCOO)Na2 + 2H2O
a) What would be a good indicator for this titration and what colour would it be in
alkali?
b) Was the acid in the burette?
c) Calculate the concentration of the acid solution.
2.
If 23,2 cm3 of sulphuric acid was required to neutralise 25 cm3 of a 0,04 mol.dm-3
solution of sodium hydroxide solution, what was the concentration of the acid?
H2SO4 + 2NaOH  Na2SO4 + 2H2O
EXTENDED TOPICS
Ionization constant of an acid, Ka
HCl(aq) + H2O(l) ⇄ H3O+(aq) + Cl-(aq)
Square brackets [ ]
mean concentration in
mol.dm-3.
Ka = [H3O+].[ Cl-]
[HCl]
The Ka value indicates the strength of an acid (i.e. degree of ionisation)
e.g. Ka for CH3COOH is 1,8 x 10-5 at 25oC
Ionization constant of a base, Kb
NH3(aq) + H2O(l) ⇄ NH4+(aq) + OH-(aq)
Kb = [NH4+].[ OH-] = 1,8 x 10-5 ( NH3 weak base)
[NH3]
pH of alkalis
1.
14 g of sodium hydroxide were dissolved in water to make 400 cm3 of solution.
What is the pH?
2.
The pH of a solution a calcium hydroxide is 10.2. What mass of calcium hydroxide
was dissolved in 0,2 dm3 of solution?
Back Titration
1.
5,0 g of calcium hydroxide were added to 250 cm3 of a solution of hydrochloric acid
with a concentration of 0,6 mol.dm-3. The resulting solution was then neutralised by a
solution of potassium hydroxide of concentration 0,2 mol.dm-3. What was the volume
of potassium hydroxide used?
Ca(OH)2 + 2HCl  CaCl2 + 2H2O
and
KOH + HCl  KCl + H2O
Acids
Strong

HCl (hydrochloric acid)

HNO3 (nitric acid)

H2SO4 (sulphuric acid)

HBr (hydrobromic acid)

HI hydroiodic acid

HClO4 (perchloric acid)
Weak

CH3COOH (acetic acid)

HCOOH (formic acid)

HF (hydrofluoric acid)

HCN (hydrocyanic acid)

HNO2 (nitrous acid)

HSO4- (hydrogen sulphate ion)

(COOH)2 (oxalic acid)
Bases
Strong

NaOH sodium hydroxide

KOH potassium hydroxide

Ba(OH)2 barium hydroxide
Weak

NH3 ammonia

CH3NH2 methylamine

C5H5N pyridine

NH4OH ammonium hydroxide