GRADE 11 SCIENCE
ACIDS AND BASES
RECAP
Properties of Acids
- taste sour
- corrosive
- turn litmus red
- pH < 7
- neutralised by bases
- react with metals to produce hydrogen
Properties of Bases
- feel soapy
- bitter taste
- turn litmus blue
- pH > 7
- neutralised by acids
- corrosive
Common Acids
- carbonic acid - H2CO3
- hydrochloric acid – HCl
- nitric acid – HNO3
- phosphoric acid – H3PO4
- sulphuric acid – H2SO4
- ethanoic acid – CH3COOH
- oxalic acid – (COOH)2
Common Bases
- ammonia – NH3
-calcium hydroxide (slaked lime)– Ca(OH)2
-calcium oxide (quick lime) – CaO
- potassium carbonate (potash) – K2CO3
- sodium carbonate (washing soda) – Na2CO3
- sodium bicarbonate (baking soda) – NaHCO3
- sodium hydroxide (caustic soda) – NaOH
Arrhenius’ theory (First theory of acids and bases)
ACID: substance which produces hydrated oxonium (or hydrogen) ions in aqueous solution.
HCl + H2O  H3O+ + Cl-
(ionisation)
ALKALIS: substance which produces hydrated hydroxide ions in aqueous solution.
NaOH(s)  Na+(aq) + OH-(aq)
(dissociation)
- Strong acid or strong base: highly (fully) ionised in solution
e.g.
HCl, H2SO4, HNO3
H2SO4 + 2H2O  2H3O+ + SO42all soluble metal hydroxides e.g. NaOH, KOH, Ca(OH)2
Ca(OH)2  Ca2+(aq) + 2OH-(aq)
- Weak acid or weak base: slightly ionised in solution
e.g.
all organic acids e.g. ethanoic acid CH3COOH
CH3COOH(aq) + H2O(l) ⇌ CH3COO-(aq) + H3O+(aq)
base : ammonia NH3
NH3(aq) + H2O(l) ⇌ NH4+(aq) + OH-(aq)
BUT Arrhenius doesn’t explain how acids/bases still react when not in aqueous solution
Bronsted – Lowry Theory (40 yrs after Arrhenius)
ACID IS A PROTON (H+) DONOR
monoprotic – has 1 H+ to donate eg HCl
diprotic - has 2 H+ to donate eg H2SO4
triprotic - has 3 H+ to donate eg H3PO4
BASE IS A PROTON (H+) ACCEPTOR
(learn these as definitions of acid/base)
HCl + NH3  NH4+ + Clacid base
This is a PROTOLYSIS reaction because a
proton (H+) is transferred.
Conjugate acid – base pairs
The above reaction is reversible so: NH4+ + Cl-  HCl + NH3
acid
base
Every base has a conjugate acid and every acid a conjugate base.
Cl- is the conjugate base of the acid HCl
HCl + NH3 ⇄ NH4+ + Clacid base
acid
base
NH4+ is the conjugate acid of the base NH3
Note: the conjugate acid always has an extra H+ compared to its conjugate base.
the conjugate acid of OH- is H2O
e.g.
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 – it can donate or gain a proton
HCl + H2O  H3O+ + Clbase
NH3 + H2O ⇄ NH4+ + OHacid
Other commonly asked examples:
HCO3-
HSO4-
a) Identify the conjugate acid-base pairs in the following reversible reactions:
HSO4- + OH- ⇄ SO42- + H2O
CH3COOH + H2O ⇄ CH3COO- + H3O+
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-? _______________
Reactions
1. ACID + REACTIVE METAL  SALT + HYDROGEN
(Zn, Fe, Mg, Ca)
2HCl + Mg  MgCl2 + H2
2. ACID + BASE  SALT + WATER
(note: a base is a metal hydroxide or metal oxide)
H2SO4 + 2NaOH  Na2SO4 + 2H2O
3. ACID + METAL CARBONATES  SALT + WATER+ CARBON DIOXIDE (also for hydrogen carbonates)
2HNO3 + CaCO3  Ca(NO3)2 + H2O + CO2
Note:
-
the salt of hydrochloric acid is a chloride
the salt of nitric acid is a nitrate
the salt of sulphuric acid is a sulphate
the salt of phosphoric acid is a phosphate
Write the balanced chemical equation for the reaction between the following:
a)
aluminium and hydrochloric acid
b)
copper oxide and nitric acid
c)
sodium carbonate and sulphuric acid
d)
magnesium hydroxide and phosphoric acid
e)
sodium hydroxide and ethanoic acid
f)
potassium hydrogen carbonate and sulphuric acid
Neutralisation
The reaction between an acid and a base to form salt.
If the base is a solution of a hydroxide the actual reaction that happens is:
H+ + OH-  H2O
or
H3O+ + OH-  2H2O
If the base is a carbonate or hydrogen carbonate then the actual reaction is:
CO32- + 2H3O+  CO2 + 3H2O
HCO3- + H3O+  CO2 + 2H2O
Preparation of salts
Salts are ionic substances consisting of anions and cations – produced by reacting acids and bases.
SALT = CATION + ANION
From Base
From Acid
Write Equations for the preparation of the following:
a) sodium sulphate from an oxide
b) magnesium chloride from a hydroxide
c) zinc phosphate from a metal
d) ammonium nitrate
Ionization constant of an acid, Ka
HCl(aq) + H2O(l) ⇄ H3O+(aq) + Cl-(aq)
Ka = [H3O+].[ Cl-]
[HCl]
Square brackets [ ] = concentration in mol.dm-3
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 (low value therefore weak acid)
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]
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.
a) 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?
b) 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-
c) If a solution of sulphuric acid contains OH- ions 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 fully ionised acids and alkalis i.e. strong ones
pH of acids
pH = -log[H+] c = n/V n = m/M
a) What is the pH of a solution of 0,004 mol.dm-3 hydrochloric acid?
b) What is the concentration of a H2SO4 solution with pH 0,22? [Hint: diprotic acid]
c) What is the pH of a solution of nitric acid if 16,4 g of HNO3 is mixed with water to make 400cm3 of
solution?
d) 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.
pH of alkalis
When working out the pH of alkalis one has to us the ionic product of water at 25oC.
Kw = [H+].[OH-] = 10-14
Eg
or
pH + pOH = 14
What is the pH of a 2,0 x 10-3 mol.dm-3 solution of NaOH
[OH-] = 2 x 10-3 mol.dm-3
pOH = - log[OH-] = -log(2 x 10-3) = 2,70
pH = 14 – pOH = 14 – 2,70 = 11,30
a) What is the pH of a solution if 2,2 g of Ca(OH)2 is dissolved in water to make 500 cm3 of solution?
b) What is the concentration of a solution of KOH with a pH of 8,7?
c) What mass of sodium hydroxide is needed to make 250 cm3 of solution with a pH of 12,5?
pH of salts in solution
a) Salts derived from a strong acid and a strong base give neutral solutions.
b) When a salt derived from a strong base and a weak acid dissolves in water, the solution becomes
basic.
c) When a salt derived from a strong acid and a weak base dissolves in water, the solution becomes
acidic.
d) Salts derived from a weak base and a weak acid may be acidic or basic. To determine whether a
salt will give an acidic or basic solution the Ka and Kb values of the weak acid and weak base have
to be compared. If Ka > Kb it will give an acidic solution
Assign the the following salts as acidic, basic or neutral:
a) calcium chloride
b) ammonium nitrate
c) potassium sulphate
d) sodium ethanoate
e) ammonium ethanoate
pH indicators
An indicator is a dye which changes colour with a certain change in the pH of its surroundings.
Some indicators are natural
e.g Beetroot – red in acid – purple in alkali
Commonly used laboratory indicators – see titrations
Universal indicator is a mixture of dyes which changes colour as the pH changes.
Solution
strong acid
weak acid
neutral
weak base
strong base
pH
<3
3-6
7
8-11
>11
Colour
red
orange/yellow
green
blue
purple
Titrations
A titration is used to determine the concentration of a solution using a standard solution.
Definition: A standard solution is one where the concentration is known
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=
C=
n
V
m
MxV
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.
Selecting 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
8,3 – 10,0
in acid
red
yellow
colourless
in alkali
yellow
blue
red
titration type
strong acid/ weak base
strong acid/ strong base
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).
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 sometimes 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 x Va a
=
Cb x Vb b
a = coeffiecient of acid
b = coefficient of base
Exercise
2,4 g of NaOH was weighed out and dissolved in some water. This was then made up to 100 cm 3 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.