ACIDS, BASES AND SALTS
Objectives
At the end of the lesson, the students should be able to
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Define acids and bases
Mention the physical properties of acids and bases
Explain the chemical properties of acids of acids and bases
Perform simple calculations on pH and pOH
Define and classify salts
Describe the types of salts
Describe the methods of preparing salts
Explain salt hydrolysis
Key words
Acid
Base
Salt
Hydrolysis
WEEK 1, PERIOD 1: Definition of acids and bases
Acids and bases have been defined differently by different scientists.
Svante Arrhenius, a Swedish scientist defined acids as substances that produce
hydrogen ions as the only positive ions when dissolved in water and bases as
substances that produce hydroxide ions as the only negative ions in water.
Examples of Arrhenius’acid
H2SO4, HCl, HNO3, H2CO3, HClO4
Examples of Arrhenius’ base
NaOH, KOH, Ca(OH)2
Johannes Bronsted, a Danish scientist and Thomas Lowry, an American scientist had
same views for the definition of acids and bases. They individually and at different
locations defined acids as proton donors and bases as proton acceptors.
Examples of Bronsted- Lowry acid
All Arrhenius’ acids are Bronsted-Lowry acids. Other examples of Bronsted-Lowry
acid include NH3, H2O(l)
Examples of Bronsted- Lowry base
NH3, H2O(l)
Gilbert Newton Lewis, an American scientist defined acids as electron pair acceptors
and bases as electron pair donors.
Lewis’acids are characterized by having an orbital to accommodate electrons while
Lewis’ bases are characterized by having lone pair of electrons.
Examples of Lewis’ acid
BF3, AlCl3, BCl3
Examples of Lewis’ base
All Bronsted-Lowry bases are Lewis’s bases. Other examples of Lewis’ base include
NH3, PH3, PCl3
Types of acids
Acids are two types based on their source. These include inorganic or mineral acids
and organic acids.
Inorganic acids
Inorganic acids are prepared from naturally occurring compounds called minerals
such as nitrates, sulphates, chlorides. Hence they are referred to as mineral acids.
Examples of inorganic acids
Hydrochloric acid (HCl), trioxonitrate (V) acid (HNO3), tetraoxosulphate (VI) acid
(H2SO4), trioxocarbonate (IV) acid (H2CO3), trioxosulphate (IV) acid (H2SO4).
Organic acids
Organic acids are derived from plants and animals. Hence they are naturally
occurring acids. They are also known as alkanoic acids because they are derivatives
of the hydrocarbons called alkanes. They can also be referred to as carboxylic acids
because they possess the carboxyl group (-COOH) as functional group.
Examples of organic acids
Ethanoic acid (acetic acid), ethanedioic acid (oxalic acid), 2-hydroxypropanoic acid
(lactic acid), hexadecanoic acid (palmitic acid), octadecanoic acid (stearic acid), 2hydroxypropane-1, 2, 3 – trioic acid (citric acid)
Classification of acids
Acids are classified based on concentration, strength and basicity.
Based on concentration, acids are either concentrated or dilute.
Concentrated acids contain a high proportion of the acid in a specific volume of
water.
Dilute acids contain a less proportion of the acid in a given volume of water.
Based on strength, acids are either weak or strong.
Weak acids ionize slightly in water and strong acids ionize completely in water.
Examples of weak acids
Hydrofluoric acid (HF), dioxonitrate (III) acid (HNO2), trioxocarbonate (IV) acid
(H2CO3), trioxosulphate (IV) acid (H2SO3), trioxophosphate (III) acid (H3PO3),
tetraoxophosphate (V) acid (H3PO4), all organic acids such as ethanoic acid
(CH3COOH), ethanedioic acid (COOH)2
Examples of strong acids
Hydrochloric acid (HCl), hydrobromic acid (HBr), hydroiodic acid HI),
tetraoxochlorate (VII) acid (HClO4), trioxonitrate (V) acid (HNO3),
tetraoxosulphate(VI) acid (H2SO4)
Based on basicity, acids may be monobasic, dibasic or tribasic.
The basicity of an acid is the number of replaceable hydrogen ions in a molecule of
the acid.
Monobasic acids possess one replaceable hydrogen ion in a molecule of the acid.
Examples of monobasic acid
All hydrohalic acids such as HF, HCl, HBr and HI. Other monobasic acids include
HNO2, HNO3, HOCl, HOCl4, hexadecanoic acid acid CH3CH2(CH2)13COOH,
octadecanoic acid CH3CH2(CH2)15COOH
Dibasic acids possess two replaceable hydrogen ions in a molecule of the acid.
Examples of dibasic acid
Trioxosulphate (IV) acid (H2SO3), tetraoxosulphate (VI) acid (H2SO4),
trioxocarbonate (IV) acid (H2CO3), ethanedioic acid.
Tribasic acids possess three replaceable hydrogen ions in a molecule of the acid.
Examples of tribasic acid
Trioxophosphate (III) acid, tetraoxophosphate (V) acid, 2 – hydroxypropane - 1, 2, 3trioic acid.
Physical properties of acids
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Physical state
Inorganic acids are colourless liquids.
Organic acids are either colourless liquids or white solids. Oxalic acid is a white
crystalline solid.
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Corrosiveness
Concentrated solutions of acids are corrosive.
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Solubility in water
All inorganic acids and most organic acids are water – soluble.
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Taste
Acids taste sour
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Ability to conduct electricity
Acids conduct electricity in solution because of the ions produced when dissolved in
water hence they are electrolytes
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pH
The pH values of acids are within the range of 1 and 6. In other words, pH value of
acids is less than 7.
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Colour change with indicators
Acids turn blue litmus red and have no effect on red litmus
Acids turn methyl orange solution pink or red
Acids have no colour change with phenolphthalein. Acids remain colourless with
phenolphthalein.
Chemical properties of acids
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Reaction with bases (neutralization reaction)
Acids react with bases to form salt and water
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Reaction with metallic oxides (neutralization reaction)
Acids react with metallic oxides to form salt and water
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Reaction with reactive metals (displacement reaction)
Acids react with metals more reactive than hydrogen to produce salt and hydrogen
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Reaction with carbonates (metallic trioxocarbonate (IV) compounds) or
bicarbonates (metallic hydrogen trioxocarbonate (IV) compounds)
Acids react with carbonates or bicarbonates to produce salt, water and carbon (IV)
oxide gas.
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Reaction with sulphites (metallic trioxosulphate (IV) compounds) or
bisulphites (metallic hydrogen trioxosulphate (IV) compounds)
Acids react with sulphites or bisulphites to produce salt, water and sulphur (IV)
oxide gas.
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Reaction with metallic sulphides
Acids react with metallic sulphides to produce salt and hydrogen sulphide gas.
Laboratory preparation of acids
Acids are prepared in the laboratory by four methods:
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Direct combination (preparation of hydracids)
Dissolution of acid anhydride in water
Reaction of non-metals with strong oxidizing agent like concentrated HNO3
Reaction of salts of more volatile acids with less volatile acids
Uses of acids
Acids have quite useful applications
Types of bases
Bases are two types based on their solubility in water. These include soluble bases
and insoluble bases.
Soluble bases
Soluble bases are soluble in water. They are also referred to as alkalis.
All alkalis are bases but not all bases are alkalis.
Examples of soluble bases
Sodium hydroxide (NaOH), potassium hydroxide (KOH), aqueous ammonia (NH3).
Insoluble bases
Insoluble bases are insoluble in water.
Examples of insoluble bases
Magnesium hydroxide Mg(OH)2, aluminum hydroxide Al(OH)3, copper (II) oxide
(CuO), copper (II) hydroxide, zinc hydroxide, magnesium oxide, iron (II) hydroxide,
iron (III) hydroxide
Calcium hydroxide Ca(OH)2 is slightly water soluble
Classification of bases
Bases are classified based on concentration, strength and acidity.
Based on concentration, bases are either concentrated or dilute.
Concentrated bases contain a high proportion of the base in a specific volume of
water.
Dilute bases contain a less proportion of the base in a specific volume of water.
Based on strength, bases are either weak or strong.
Weak bases ionize slightly in water and strong bases ionize completely in water.
Examples of weak bases
Magnesium hydroxide, aluminum hydroxide, aqueous ammonia
Examples of strong bases
Sodium hydroxide, potassium hydroxide
Based on acidity, bases may be monobasic, dibasic or tribasic.
The acidity of a base is the number of ionizable hydroxide ions in a molecule of the
base.
Monoacidic bases possess one ionizable hydroxide ion in a molecule of the base.
Examples of monoacidic base
NaOH, KOH
Diacidic bases possess two ionizable hydroxide ions in a molecule of the base.
Examples of diacidic base
Mg(OH)2, Ca(OH)2, Fe(OH)2, Cu(OH)2, Zn(OH)2
Triacidic bases possess three ionizable hydroxide ions in a molecule of the base.
Examples of triacidic base
Al(OH3, Fe(OH)3
Physical properties of bases
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Physical state
Bases may be solids such as sodium hydroxide, potassium hydroxide or liquids such
as aqueous ammonia.
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Touch
Bases have a soapy feel
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Corrosiveness
Concentrated solutions of bases are corrosive.
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Solubility in water
Some bases are water-soluble while some are water-insoluble
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Taste
Bases taste bitter
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Ability to conduct electricity
Bases conduct electricity in solution because they produce ions when dissolved in
water, hence they are electrolytes.
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pH
The pH values of bases are within the range of 8 and 14. In other words, pH value of
bases is greater than 7.
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Colour change with indicators
Bases turn red litmus blue and have no effect on blue litmus
Bases turn methyl orange solution yellow or orange
Bases turn colourless phenolphthalein pink or red
Chemical properties of bases
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Reaction with acids (neutralization reaction)
Bases react with acids to form salt and water
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Reaction with ammonium salt
Alkalis except ammonia react with ammonium salts to produce salt, water and
ammonia gas.
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Thermal decomposition reaction
Bases with the exception of sodium hydroxide and potassium hydroxide decompose
on heating to produce their respective oxides and water.
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Reaction with carbon (IV) oxide
Bases react with carbon (IV) oxide to produce trioxocarbonate (IV) of the metal and
water
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Reaction of alkalis with some salts
Alkalis react with solution of some salts such as salts containing metals like
aluminum, lead, zinc, iron or copper to form insoluble metal hydroxides which can
be identified by their distinct colour.
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Amphoteric nature of some bases
Some bases such the hydroxides of aluminum, zinc and lead have dual properties as
they react with acids as well as bases to form salt and water.
Laboratory preparation of bases
Bases are prepared in the laboratory by a number of methods:
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Reaction of metals with oxygen
Dissolution of certain metals more reactive than hydrogen in cold water or
steam
Thermal decomposition of certain trioxocarbonate (IV) compounds
Thermal decomposition of certain trioxonitrate (V) compounds
Uses of bases
Bases have numerous applications
pH and pOH
Meaning of pH
p in pH stands for potenz in German which means strength or power in English.
H stands for hydrogen
Definition of pH
The pH of a solution is the logarithmic reciprocal of the hydroxonium ion
concentration or the negative logarithm to base ten of the hydroxonium ion
concentration.
Mathematical representation of pH
For simplicity, the hydrogen ion is used in place of the hydroxonium ion
or
Definition of pOH
The pOH of a solution is the logarithmic reciprocal of the hydroxide ion
concentration or the negative logarithm to base ten of the hydroxide ion
concentration.
Mathematical representation of pOH
or
Calculating the pH of pure water
Consider the ionization of water:
H2O(l)
H+(aq) +
OH-(aq)
Equilibrium constant (K) expression for the ionization of water:
Ionic product of water (kw)
[H2O(aq)]K =
K=
is the ionic product of water
It’s been determine that the ionic product of water (kw) at 25 oC (298 K) is 10-14
mol2dm-6
For pure water, [H+} = 10-7 mol dm-3 and [OH-] = 10-7 moldm-3
Calculating pH of pure water
pH = - Log1010-7
= -7(-Log1010)
= -7 (-1)
pH = 7
Calculating pOH of pure water
pOH = - Log1010-7
= -7(-Log1010)
= -7 (-1)
pOH = 7
Writing the pkw of pure water
The sum of the pH and pOH of pure water is called pkw and its value is 14
pkw = pH + pOH
pH + pOH = 14
pH scale
pH scale ranges from 1 to 14. pH values between 1 and 6 are acidic. pH value of 7 is
neutral and pH values between 8 and 14, basic.
Acidity increases with increase in hydroxonium ion concentration but decrease in
the hydroxide ion concentration
Therefore acidity increases with decrease in the value of pH.
Basicity increases with increase in hydroxide ion concentration but decrease in the
hydroxonium ion concentration
Therefore basicity increases with increase in the value of pH.
The figure below shows the summary
Measurement of pH
pH can be measured by any of three methods:
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The use of litmus
The use of universal indicator
The use of pH meter
Quiz
1.
2.
3.
4.
5.
6.
7.
What is the pH of a 0.01 mol dm-3 solution of hydrochloric acid?
What is the pOH of a 0.1 mol dm-3 solution of potassium hydroxide?
What is the pH of a 0.0001 mol dm-3 solution of tetraoxosulphate (VI) acid?
What is the pH of a 0.001 mol dm-3 solution of sodium hydroxide?
What is the pOH of a 0.01 mol dm-3 solution of hydrochloric acid?
What is the hydrogen ion concentration of a solution of pH 4.398?
A sample of orange juice is found to have a pH of 3.80. What is the hydroxide ion
concentration?
8. What is the pH of a 2.50 × 10-5 mol dm-3 solution of sodium hydroxide?
9. What is the pOH of a solution of 0.25 mol dm-3 of hydrochloric acid?
10. What is the pOH of a 0.05 mol dm-3 solution of calcium hydroxide?
Definition of salts
A salt is a compound formed by the replacement of the hydrogen ion in an acid with
a metallic ion or ammonium radical.
Types of salts
Salts are of six types and they include
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Normal salts
A normal salt is one formed by the complete replacement of ionizable hydrogen ion
of an acid by a metallic ion. Hence does not contain any replaceable hydrogen ion in
its molecule. Most are neutral to litmus such as NaCl(aq), KCl(aq),Na2SO4(aq), KNO3(aq),
but some are acidic or basic. For instance, NH4Cl(aq) is slightly acidic and
CH3COONa(aq) is slightly basic.
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Acid salts
An acid salt is one formed by the partial replacement of the ionizable hydrogen ion
in an acid with metallic ions. Hence it contains replaceable hydrogen ion in its
molecule.
NaHSO4(aq) and KHSO4(aq)are acidic in solution because they formed by the reaction
between a strong acid and a strong base and contain ionizable hydrogen ion but
some acid salts such as NaHCO3(aq), KHCO3(aq), Na2HPO4(aq) are slightly alkaline in
solution because they are formed by the reaction between a weak acid and a strong
base.
Acid salts ionize in solution to form hydroxonium ion and react with bases to form
normal salts.
Examples of acid salt
NaHSO4(aq), KHSO4(aq), NaHCO3(aq), KHCO3(aq), NaH2PO4(aq), Na2HPO4(aq)
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Basic salts
A basic salt is one formed by the partial replacement of the hydroxide ion in a base
with an anion.
Basic salts are not necessarily alkaline to litmus and reacts with acid to form normal
salt
Examples of basic salts
Zn(OH)Cl, Pb(OH)Cl, Pb(OH)NO3, Cu(OH)NO3, Cu(OH)Cl
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Double salts
A double salt is formed when stoichiometric (calculated) amounts of two simple
salts combine and exists in the solid state but dissociate into their constituent simple
ions when dissolved in water.
Examples of double salts
Potash alum K2SO4.Al2(SO4)3.24H2O, Mohr’s salt FeSO4(NH4)2SO4.6H2O, carnalite
KCl.MgCl2.6H2O
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Complex salts
A complex salt is a compound in which the central metal ion is bonded to a number
of neutral molecules, anions or cations by coordinate bonds. The coordinate bonds
are formed by the donation of electron by neutral molecule, anions or cations to the
metal ion.
Examples of complex salts
[Cu(NH3)4](OH)2, K3[Fe(CN)6],
K3[Fe(CN)6]

K+
+
Simple ion
[Fe(CN)6]3complex ion
Mixed salts
A mixed salt is one that contains more than one acidic or basic radical other than
hydrogen or hydroxide ions.
Examples of mixed salts
CaOCl2 (bleaching powder), KNaCO3, CaKPO4, NH4NaHPO4
Classification of salts
Salts are classified based on their solubility in water into soluble salts and insoluble
salts, as well as based on their stability to heat.
Soluble salts are water soluble.
Examples of soluble salts
All sodium containing salts
All potassium containing salts
All ammonium salts
All dioxonitrate (III) salts
All trioxonitrate (V) salts
All bicarbonate (metallic hydrogen trioxocarbonate (IV) salts
All ethanoate salts
All tetraoxosulphate (VI) salts except those containing calcium, barium, mercury (I),
lead or silver
All chlorides except those containing silver, lead or mercury (I)
Examples of insoluble salts
All sulphides
All carbonates except those containing sodium, potassium or the ammonium radical
All tetraoxophosphate (V) salts except those containing sodium, potassium or the
ammonium radical
All trioxophosphate (III) salts except those containing sodium, potassium or the
ammonium radical
All trioxosulphate (IV) salts except those containing sodium, potassium or the
ammonium radical
Based on stability to heat, salts are either stable to heat or unstable to heat.
Salts stable to heat do not decompose on heating
Examples of salts stable to heat
Na2CO3, K2CO3
Salts unstable to heat decompose on heating
Examples of salts unstable to heat
CaCO3, NaHCO3, all trioxonitrate (V) salts
Physical properties of salts
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Physical state
Salts are solids at room temperature
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Colour
Some salts appear white in colour. Examples include salts of sodium and potassium.
Some other salts have characteristic colour such as blue (copper containing salt),
green (iron (II) containing salt), brown (iron (III) containing salt), orange (dichromate
containing salt).
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Solubility in water
Some salts are water soluble while others are water insoluble. Some are said to be
slightly or sparingly water soluble.
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Ability to conduct electricity
Salts conduct electricity in the molten state or in solution because of the presence of
ions when in these states. Hence they are electrolytes.
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Melting temperature
The melting point of salt is quite high due to electrostatic force of attraction between
the ions that make up the salt.
Laboratory preparation of salts
Salts are prepared in the laboratory according to their solubility in water.
Methods of preparing water soluble salts
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Reaction of acids with bases
Reaction of acids with metallic oxides
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Reaction of acids with reactive metals
Reaction of acids with carbonates and bicarbonates
Methods of preparing water insoluble salts
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Direct combination of the constituent elements of the salt
Double decomposition reaction (precipitation reaction)
Definition of salt hydrolysis
Salt hydrolysis is the partial reversal of neutralization reaction when the metallic
ions or acid radicals of the salt react with water to form a weak base or weak acid.
Mechanism of salt hydrolysis
Salts formed by the reaction between strong acids and strong bases do not undergo
hydrolysis and their solution is neutral. This is because the acid and base that form
them ionize completely in water. Examples of such salts include NaCl, K2SO4
Salts formed by the reaction between strong acids and weak bases undergo
hydrolysis to produce an acidic solution. This is because the cation of the salt
undergoes hydrolysis by reacting with the hydroxide ion from water to form
solution of weak base with the formation of hydrogen ions. The hydrogen ions
formed is responsible for the acidity of the salt solution. This is known as cation
hydrolysis. Examples of such salts include NH4Cl, CaSO4, FeCl3.
Consider the hydrolysis of NH4Cl(aq),
In solution:
NH4Cl(aq)
H2O(l)
NH4+(aq)
+
Cl-(aq)
H+(aq) + OH-(aq)
The cation from the salt combines with the hydroxide ion from water:
NH4+(aq) + OH-(aq)
NH3(aq)
+
H2O(l)
(i)
H2O(l) ionizes:
H2O(l)
H+(aq) + OH-(aq)
(ii)
Combining both ionic equations (i) and (ii)
NH4+(aq) + OH-(aq) + H2O(l)
NH3(aq)
NH4+(aq)
NH3(aq)
+
+ H2O(l) + H+(aq) + OH-(aq)
H+(aq)
Responsible
for acidity
Salts formed by the reaction between weak acids and strong bases undergo
hydrolysis to produce a basic solution.. This is because the anion of the salt
undergoes hydrolysis by reacting with the hydrogen ion from water to form a
solution of weak acid with the formation of hydroxide ions. The hydroxide ions
formed is responsible for the basicity of the salt solution. This is known as anion
hydrolysis. Examples of such salts include Na2CO3, KCN, Na2S, CH3COONa.
Consider the hydrolysis of CH3COONa(aq),
In solution:
CH3COOH(aq)
H2O(l)
CH3COO-(aq)
+
H+(aq)
H+(aq) + OH-(aq)
The anion from the salt combines with the hydrogen ion from water:
CH3COO-(aq) + H+(aq)
CH3COOH(aq)
(iii)
H2O(l) ionizes:
H2O(l)
H+(aq) + OH-(aq)
(iv)
Combining both ionic equations (iii) and (iv)
CH3COO-(aq) + H+(aq) + H2O(l)
CH3COO-aq) + H2O(l)
CH3COOH(aq) + H+(aq) + OH-(aq)
CH3COOH(aq)
+
OH-(aq)
Responsible
for basicity
Salts formed by the reaction between weak acids and weak bases undergo
hydrolysis but whether the resulting salt solution is acidic or basic depends on the
relative strength of the acid and base that formed the salt. Examples of such salts
include CH3COONH4, (NH4)3PO4
.
Multiple choice Quiz
1. Which of the following statements is not a chemical property of an acid?
A. evolution of ammonia gas when heated with ammonium salts
B. evolution of CO2 gas when added to a trioxocarbonate (IV) salts
C. formation of salt and water with alkalis
D. formation of salt and hydrogen gas with reactive metals
2. A substance L reacts with NH4NO3(aq) to generate ammonia gas. L is likely to be
A. HCl
C. CH3COOH
B. NaOH
D. CaSO4
3. Which of the following chlorides is insolubke in water?
A. AgCl
C NH4Cl
B. KCl
D. ZnCl2
4. Which of the fooling hydrohalic acids is the weakest?
A. HBr
C. HF
B. HCl
D. HI
5. Which of the following pH values indicates that a solution is a strong base?
A. 1
C. 9
B. 5
D. 13
6. The hydrolysis of NH4Cl gives
A. an acidic solution
C. a buffer solution
B. an alkaline solution
D. a neutral solution
7. Which of the following acids would readily react with CaCO3 to liberate CO2?
A. CH3COOH
C. H2SO3
B. H2SO4
D. HNO3
8. Which of the following oxides is basic?
A. NO2
C. SO2
B. Al2O3
D. CaO
9. An example of an acid salt is
A. CH3COONa
C. NaHSO4
B. Mg(OH)Cl
D. (NH4)2SO4
10. The aqueous solution which has pH > 7 is
A. FeCl3(aq)
C. KNO3(aq)
B. CuSO4(aq)
D. Na2CO3(aq)