CONTENT

• Definition.
• Properties.
• Preparation and uses.
• pH.
• Bases
• Salts
• Efflorescence, deliquescence and hygroscopic

 

ACIDS

Definition: An acid is a substance that in aqueous solution produces hydroxonium ion (H3O+) or hydrogen ion (H+) as the only positive ion. Also, acids can be referred to as proton donors.

 

CLASSES OF ACIDS

There are two classes of acids:

(1). Organic acids occur as natural products in plants and animal material.

Organic acids	Source
Ethanoic acids	Vinegar
Lactic acids	Milk
Citric acids	Lime, Lemon
Amino acids	Proteins
Fatty acids	Fats and oils
Ascorbic acids (Vitamin C).	Oranges

 

 

 

(2). Inorganic acid: Inorganic acid can be prepared from mineral elements or inorganic matter.

Inorganic acid	Formula	Constituents
Hydrochloric acid	HCl	Hydrogen, Chlorine
Tetraoxosulphate(VI) acid	H2SO4	Hydrogen, Sulphur & Oxygen
Trioxonitrate(VI) acid	HNO3	Hydrogen, Nitrogen & Oxygen

An acid is also defined as a substance that produces a hydroxonium ion as the only positive ion when dissolved in water.

H+(aq)    +    H2O(l)       H3O+(aq)

Acid can be dilute or concentrated depending on the amount of water added. A dilute acid is acid produced when a large amount of water is added to a small amount of acid. A concentrated acid is an acid produced when only a small amount of water is added to a relatively large amount of acid.

 

STRENGTH OF AN ACID

The strength of an acid can either be weak or strong.

(1) Strong acids are acids that ionize completely in aqueous solution, and such acid solutions have a high concentration of H+. Examples are HCl, H2SO4 and HNO3.

H2SO4       2H+ + SO42-

HNO3     H+ + NO3–

HCl                  H+ + Cl–

 

(2) Weak acids: are acids which ionize or dissociate slightly or partially in aqueous solution and such acid solution have a low concentration of hydrogen ions. Examples are ethanoic acid (CH3COOH), H2CO3, H3PO4, H2SO3.

 

H2CO3                                   2H+    +   CO32-

H3PO4                                    3H+    +   PO43-

CH3COOH                                  H+    +   CH3COO–

H2SO3                                 2H+   +   SO32-

 

BASICITY OF AN ACID

The basicity of an acid is the number of replaceable hydrogen ions, H+, in one molecule of the acid.

Acid	Basicity
Hydrochloric   acid	Monobasic
Tetraoxosulphate (vi) acid	Dibasic
Tetraoxophosphate(v) acid	Tribasic
Ethanoic acid	Monobasic

 

EVALUATION

1. Define the term acid
2. Differentiate between strong acid and concentrated acid
3. What is the basicity of the following acids: HCl, HNO3, H2SO4

 

PHYSICAL PROPERTIES OF ACID

1. They have a sour taste.
2. They turn blue litmus paper to red.
3. They are corrosive in nature, especially the strong acid.
4. In aqueous solution, they conduct electricity.

 

CHEMICAL PROPERTIES OF ACID

Reaction with metals: They react with metals to liberate hydrogen gas and the salt of the metal, i.e

Acid   +   Metal               Salt           + Hydrogen gas.

E.g.  2HCl(aq)   +   Zn(s)                      ZnCl2(aq)     +    H2(g)

H2SO4(aq)  +   Mg(s)                     MgSO4(aq)    +    H2(g)

 

They react with soluble bases to form salt and water only. This reaction is known as neutralization.

Acid     +     Base                      salt        +      water

E.g  H2SO4(aq)     +   2KOH(aq)                         K2SO4(aq)   +    2H2O(l)

2HCl(aq)     +     CaO(s)                           CaCl2(aq)      +      H2O(l)

 

1. They react with  trioxocarbonates (iv) salts to liberate carbon (iv) oxide, salt and water  i.e. Acid   +  trioxocarbonate (iv)                           Salt  +     Water   +    CO2

E.g  2HCl (aq) +    Na2CO3(aq)                           2NaCl(aq)  +      H2O(l)  +  CO2(g)

 

PREPARATION OF ACIDS

Acid can be prepared by using the following methods:

1. Dissolving an acid anhydride in water: Acid anhydride is oxides of non-metal that dissolve in water to produce the corresponding acids e.g SO2, CO2, CO, NO2, SO3.

SO2(g)      +    H2O(l)                     H2SO3(aq)

CO2(g)    +   H2O(l)                   H2CO3(aq)

SO3(g)    +   H2O(l)                  H2SO4(aq)

 

2. Combination of constituent elements.

(a).Burning hydrogen in chlorine, in the presence of activated charcoal as the catalyst, yields HCl gas which dissolves readily in water to give HCl acid.

H2(g)     +         Cl2(g)            activated charcoal       2HCl(g)

(b)Heating hydrogen gas and bromine vapour, in the presence of platinum as the catalyst, produces hydrogen bromide which dissolves readily in water to form hydrobromic acid.

H2(g)         +          Br2(g)                  Platinum         2HBr(g)

(3)By displacement of a weak or more volatile acid from it salt by a stronger or less

volatile acid. For example

(a)Displacement of the more volatile hydrogen chloride from metallic chloride by the less volatile concentrated tetraoxosulphate (vi) acid.

NaCl(s)   +  H2SO4(aq)             NaHSO4(aq)        +     HCl(aq)

 

(b)Displacement of weaker trioxoborate (iii) acid from ‘borax’ by tetraoxosulphate (vi) acid.

Na2B4O7(s) +  H2SO4(aq)    +   5H2O(l)               Na2SO4(aq)    +    4H3BO3(aq)

Borax                                            Trioxoborate (iii) acid

(4)By precipitating an insoluble sulphide from a metallic salt by hydrogen sulphide

Pb (CH3COO)2(aq)    +  H2S(g)                 PbS(s)    + CH3COOH(aq)

 

Uses of acid

(1) Acids are useful chemicals which are used in many industries to make other consumer chemicals such as fertilizers, detergent and drugs.

(2) They are used in industrial process as drying agents, oxidizing agents and catalysts.

 

USES OF ORGANIC AND INORGANIC ACID

Name	Uses
HCl	Needed by industries to make chemicals used to remove rust. Used to clean the surface of metals before electroplating.
H2SO4	Needed by industries to make chemicals used as a drying and dehydrating agent. Used as an electrolyte in lead-acid accumulators Required in oil refineries.
HNO3	Needed by industries for making fertilizers, explosives etc.
Boric acid	Used as mild antiseptic or germicide.
Tartaric acid	Used in making baking soda, soft drinks and health salts
Acetic acid (ethanoic acid)	Used in preserving food. Used in dyeing silk and other textiles.
Citric acid	Used in making fruits juice.
Fatty acid (palmitic and stearic acid)	Used in the manufacture of soap. This process is known as saponification. Fatty acid + Caustic soda     Soap   +     H2O.

EVALUATION

1. Mention three physical properties of acids
2. Using balanced equations, state the chemical properties of acids
3. State two methods of preparing acids
4. Outline the uses of acids

 

pH SCALE

All acidic solution contains H+ and all alkaline solution contains OH– ions. The PH scale measure the concentration of H+ ions present in a solution and start from 0 to 14

 

DEFINITION OF pH

pH is defined as the negative logarithms of the hydrogen ion [H+] concentration to the base of 10.

i.e. pH = -log [H+].

Thus: If [H+] = 0.00001 or 10-5.

log [H+] = log10-5 = -5

pH= -log [H+] = – (-5) = 5.

If [H+] =10-x

Therefore, pH= -log10-x = – (-x) = x

If [H+] = 10-2, PH = 2

 

DEFINITION OF pOH

POH is defined as the negative logarithms of the hydroxide ion [OH–] concentration to the base of 10.

i.e. pOH= -log [OH–].

PH is the degree of acidity. A solution with PH 7 is neutral. A solution with PH less than 7, i.e. PH 6,5,4, e.t.c, indicate acidity increasing as the numbers decreases. A solution with PH greater than 7, i.e. PH 8,9,10, e.t.c, indicate alkalinity increasing as the numbers increase.

PH 1        2     3     4    5      6      7     8     9     10     11     12     13     14

 Increasing acidity ⇐           ⇑Neutral     —⇒ Increasing alkalinity

 

A solution with PH 1 is very acidic [with high concentration of H+]. A solution with pH 13 is very alkaline [with low concentration of H+, but high concentration of OH–].

Note that: If pH is 1, it has a concentration of H+ 10 times greater than pH 2 and 100 times greater than PH 3 e.t.c.

pH 1 > pH 2 > pH 3.

Concentration of H+    10-1      10-2      10-3.

       0.1       0.01     0.001.

 

Relationship between pH and pOH.

H2O              H+    +     OH–

From conductivity measurement, [H+]=10-7moldm-3, [OH–]=10-7moldm-3.

[H+] [OH–] = Kw=10-7 x 10-7=10-14mol2dm-6.

Taking logarithm of both sides

log ([H+] [OH–]) = logKw

log [H+] + log[OH–] =logKw

Subtracting both sides

-(log[H+]  +  [OH–]) = -logKw

-log [H+] – log[OH–] = -logKw

-log [H+] + (-log [OH–]) = -logKw

pH + pOH = PKw

pKw = -log10-14 = -(-14) = 14

Therefore, pH + pOH = 14.

 

Worked examples

1. Find the hydrogen and hydroxide ion concentrations in

(a) 0.01moldm-3 tetraoxosulphate (vi) acid solution.

(b) 0.001moldm-3 potassium hydroxide solution.

 

Solution

(a).  H2SO4(aq)                      2H+(aq)  +  SO42-(aq)

 

From the equation,  1 moldm-3 H2SO4 ionizes to give 2moldm-3 H+

Therefore, 0.01moldm-3 H2SO4 would ionize to give (2×0.01) moldm-3 H+

[H+] = 2×10-2moldm-3

[H+] [OH–] = 10-14

(2×10-2) [OH–] = 10-14

[OH–] =    10-14

2x 10-2

[OH–] = 0.5x (10-14- -2)

[OH–] =0.5 x10-14+2

[OH–] =0.5×10-12moldm-3.

(b).     KOH(aq)   →   K+(aq)  +     OH–(aq)

From the equation,

1moldm-3 of KOH ionizes to give 1moldm-3 of OH–

10-3moldm-3 of KOH would ionize to give 10-3moldm-3 of OH–

[OH–]=10-3moldm-3.

[H+] [OH–]=10-14

[H+] (10-3) = 10-14.

[H+] = 10-14

10-3

[H+] = 10-14+3

[H+] = 10-11moldm-3

 

2. A glass cup of orange juice is found to have a POH of 11.40. Calculate the concentration of the hydrogen ions in the juice.

Solution

pH + pOH = 14.

pH = 14 – 11.4.

pH = 2.6.

pH = -log [H+]

2.6 =-log [H+].

[H+] = Antilog (-2.6)

[H+] = 0.0025moldm-3

[H+] = 2.5×10-3moldm-3.

 

Measuring pH of a solution.

We use pH meter and a universal indicator to detect PH of a solution.

Universal indicator is a mixture of indicator and can change to several colours corresponding to a particular PH and compared with the standard colour provided by the manufacturer of the universal indicator. Universal indicator measures PH between 3 and 11.

Process:-

Put 10cm3 of test solution in a test tube, add 2 drops of universal indicator, and compare with the colour chart or place 2 drops of test solution on universal indicator paper and compare the colour with the chart.

 

EVALUATION

1. Define the term pH.
2. What is the pH of a solution having a hydrogen ion concentration of 6x 10-9

mol/dm3

 

GENERAL EVALUATION/REVISION

1. Give the chemical formula of the following acids (a) Tetraoxosulphate (vi) acid

(b) Trioxonitrate (v) acid (c) Oxochlorate (i) acid

2. What is the IUPAC nomenclature of the following (a) HNO2 (b) HOBr (c) H3PO4 (d) H2S
3. Determine the oxidation number of Cl and C in each of the following (a) KClO3

(b) HOCl (c) H2CO3 (d) CO2

4. Mention the laboratory apparatus that are used in for an acid-base titration

What can be used to determine the acidity or alkalinity of a solution?

 

 

WEEKEND ASSIGNMENT

1. The following acids are monobasic except (a) HNO2 (b) HBr (c) HOCl (d) H2SO3
2. Which of the following ions is acidic? (a) K+   (b) NO3–     (c) S2-   (d) H3O+.
3. The number of hydroxonium ions produced by one molecule of an acid in aqueous solution is it (a) acidity (b) basicity (c) concentration (d) pH.
4. The basicity of ethanoic acid CH3COOH is: (a) 0 (b) 1 (c) 2 (d) 3
5. A solution with pH 7 is (a) Acidic  (b) dilute    (c) neutral   (d) saturated

 

THEORY

1. 1. What is (i) an acid (ii) the basicity of an acid?
2. What is the basicity of tetraoxophosphate (V) acid?
3. The concentrations of H+ in two solutions are (a) 1 × 10^-4 mol dm^-3 and (b) 5 × 10^-9 mol dm-3. What is the PH of each solution?

 

 

CONTENT

• Bases
• Salts
• Efflorescence, deliquescence and hygroscopic

 

BASES AND ALKALIS

A base is a substance which will neutralize an acid to yield a salt and water only. Most oxide and hydroxide of metals are bases e.g. Na2O, K2O, MgO, NaOH, KOH e.t.c.

 

An alkalis is a basic hydroxide which is soluble in water NaOH, KOH, Ca(OH)2. A basic oxide (or hydroxide) is a metallic oxide (or hydroxide) which contains ions (O2- or OH–) and will react with an acid to form a salt and water only.

 

Note: An exception to this definition is the reaction of lead (IV) oxide with hydrochloric acid to produce lead (ii) chloride (a salt), water and chlorine gas.

PbO2(s) +   4HCl (aq)               PbCl2 (aq) +    2H2O (l)   +   Cl2 (g)

 

From the equation above, PbO2 is not a base.    The nature of the hydroxides of the metals varies according to the position of the metal in the electrochemical series, as illustrated below.

Metal	Solubility	Decomposition by heat.
K Na Ca	The hydroxides of these metals are soluble in water and are alkalis.	Hydroxide of sodium and potassium can not be decomposed by heat.
Mg, Al, Zn Fe, Pb, Cu	These metals form hydroxides which are insoluble in water. They are amphoteric except the three hydroxides of Magnesium, iron, and copper.	Decomposed on heating to form oxide and water.
Hg, Ag, Au.	Hydroxides of these metals do not exist.

 

STRENGTH OF BASE

Like an acid, we have strength of a base. The strength of a base can either be weak or strong.

Weak base: Weak base are base that ionizes slightly in aqueous solution to produce positively charged metallic ion and negatively charged hydroxide ion e.g. CaO, NH3.

 

NH3 (g)    +   H2O (l)                NH4+(aq) OH–(aq)

Strong bases: Strong base are base that ionizes completely in aqueous solution to produce positively charged metallic ion and negatively charged hydroxide ion e.g. Na2O, K2O.

Na2O(s)   +      H2O (l)                   2NaOH(aq)

K2O(s)      +     H2O (l)             2KOH(aq)

NaOH (aq)            Na+(aq)  +    OH–(aq)

KOH (aq)              K+(aq)         +       OH–(aq)

 

PHYSICAL PROPERTIES OF ALKALIS

1. Alkalis have a bitter taste.
2. Alkalis are soapy to the touch.
3. Alkalis turn red litmus blue.
4. Concentrated form of the caustic alkalis of NaOH and KOH are corrosive.

 

CHEMICAL PROPERTIES OF ALKALIS

1. Reaction with acid: All base react with acid to form salt and water only.

NaOH(aq)   +    HCl(aq)                 NaCl(aq)       +    H2O(l)

MgO(s)      +   2HNO3(aq)             Mg(NO3)2(aq)   +   H2O(l)

2. Displacement of volatile ammonia from ammonium salt by a non volatile alkali: if an ammonium salt is warmed with an alkali (in the presence of water) ammonia gas is liberated

NaOH (aq)   +      NH4Cl (aq)                  NaCl (aq)    +     H2O(l)   +     NH3(g)

Ca (OH)2(aq)  +  (NH4)2SO4(aq)               CaSO4(aq)   +    2H2O(l)  +   2NH3(g)

 

USES OF ALKALIS/BASE

Names	Uses
NaOH	Used in the manufacture of soap, Na salts and plastic Used in petrol refining.
KOH	Used in the manufacture of liquid soap Used in dyeing and electroplating.
Ca(OH)2	Used in manufacture of  mortar, cement and plaster Used in dissolving acidic soil.
Mg(OH)2	Used in the manufacture of toothpaste Used as a laxative.
Aqueous NH3	Used for bleaching cloth Used as detergent.

 

EVALUATION

1. Define alkali giving examples
2. State three physical properties of alkalis
3. Using balanced equations, state two chemical properties of bases
4. State the uses of bases

 

NEUTRALIZATION REACTION

Neutralization reaction can be defined in three major ways.

1. In terms of acid and the base present.
2. In terms of H+ ion and OH– present in the acid and base.
3. In terms of oxonium ions (H3O+) and hydroxide ion (OH–).

Neutralization is the process whereby an acid react completely with an alkalis/bases to form salt and water.

HCl(aq)   +   NaOH(aq)                    NaCl(aq)    +    H2O(l)

Neutralization is the combination of hydrogen ion (H+) and hydroxide ions (OH–) to form water molecules. A salt is also form at the same time.

H+(aq)    +     OH–                        H2O(l)

Neutralization can also be defined as the combination of oxonium ions (H3O+) and hydroxide ions (OH–) to form water molecule. A salt is also formed at the same time.

H3OCl(aq)        +      KOH(aq)                         KCl(aq)    +    H2O(l)

(H3O)2SO4(aq)   +      2NaOH(aq)                      Na2SO4(aq)    +     4H2O(l)

H3O+(aq)              +      OH–(aq)                            2H2O(l)

 

OXONIUM ION

In aqueous solution, the hydrogen ion become associated with a water molecule to form oxonium ion (H3O+).This is an example of a coordinate covalent combination.

H+    +   H2O                      H3O+

During neutralization, oxonium ion H3O+ behaves as hydrogen ion and thus reacts with hydroxide ion (OH–) to form water molecules.

 

EVALUATION

1. Define the term neutralization.
2. Write TWO balanced equations to show neutralization reactions.

 

SALTS

A salt is referred to as the compound formed when all or part of the ionisable hydrogen ion in an acid is replaced by a metallic or ammonium ion e.g.

HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)

H2SO4(aq) + KOH(aq) → KHSO4(aq) + H2O(l)

 

TYPES OF SALTS

There are five main types of salts namely:

1. Normal salt.
2. Acid salts
3. Basic salts
4. Double salts.
5. Complex salts.

 

1. Normal salts: are the salts formed when all the replaceable hydrogen ion in the acid has been completely replaced by a metal ion e.g. NaCl, K2SO4, Na3PO4, NaNO3 etc. Normal salts are neutral to litmus

HCl(aq)   +    NaOH(aq)                     NaCl(aq)  +    H2O(l)

H2SO4(aq)  +  KOH(aq)                        K2SO4(aq)  +   H2O(aq)

 

2. Acid salts: Acid salts are formed when the replaceable hydrogen ion in the acids are only partially replaced by a metal e.g. NaHSO4, Na2HPO4, NaH2PO4, NaHCO3. They can be produce from acids which contain more than one replaceable hydrogen ion. Acids with two replaceable hydrogen ions can form only one acid salt while acid with three replaceable hydrogen ions can form two different acid salts

H2SO4(aq)    +    NaOH(aq)                        NaHSO4(aq)  +  H2O(l)

2H3PO4(aq)   +   3NaOH(aq)                       NaH2PO4(aq)  + Na2HPO4(aq)  +3H2O(l)

Acid salts turn blue litmus red. Acid salts can be converted to normal salt if the remaining replaceable hydrogen ions in the acid salt are replaced in with metallic ions.

KHSO4(aq)   +    KOH(aq)                          K2SO4(aq)  +  H2O(l)

 

3. Basic salts: Basic salts are formed when only part of the hydroxide ions of a base are replaced by the negative ions from an acid. It can occur when there is insufficient supply of acid for complete neutralization of the base e.g Zn(OH)Cl , Mg(OH)Cl, Mg(OH)NO3, Bi(OH)2NO3 e. t .c.

Zn(OH)2(aq)   +   HCl(aq)                          Zn(OH)Cl(aq)  +   H2O(l)

 

Because of the presence of hydroxide ion in the salt, it has basic properties. Basic salts turn red litmus blue. Basic salts react with excess acid to form a normal salt and water only.

Mg(OH)NO3(aq)   +  HNO3(aq)                     Mg(NO3)2(aq)   +  H2O(l)

 

4. Double salts: Double salts are salt which ionize to produce three different types of ions in solution. Usually two of these are positively charged (metallic or NH4+ ion) while the other is negatively charged e.g. (NH4)2Fe(SO4)2.6H2O, KAl(SO4)2.12H2O, KCr(SO4)2.12H2O.

(NH4)2Fe(SO4)2.6H2O: Ammonium iron (II) tetraoxosulphate (VI) hexahydrate.

KAl (SO4)2.12H2O: Aluminium Potassium tetraoxosulphate (V) dodecahydrate (Potash alum).

KCr (SO4)2.12H2O: Chromium (III) Potassium tetraoxosulphate (VI) dodecahydrate (Chrome alum).

 

5. Complex salts: Complex salts contains complex ion i.e ion consisting of a charged group of atom e.g. Na2Zn(OH)4, K4Fe(CN)6, NaAl(OH)4.

Na2Zn(OH)4: Sodium tetrahydroxozincate (ii)

K4Fe(CN)6 : Potassium hexacyanoferrate (iii)

NaAl(OH)4: Sodium tetrahydroxoaluminate (iii)

Na2Zn(OH)4                  2Na+   +   [Zn(OH)4]2-

K4Fe(CN)6                       4K+    +     [Fe(CN)6]4-.