Acids & Bases (Part II): Preparation, Properties & Uses of Bases

In a layman's term, a base is the opposite of an acid. In other words, a base is everything an acid is not. Recall that in the Part I of this series, we looked at the definition of an acid using three different concepts. Similarly, we are going to define a base using the same concepts.

Definition

Lewis Bases
According to G. N. Lewis, a base is any species that can readily donate a pair of electrons. The availability of lone pair(s) of electrons increases a substance's ability to behave as a base. Examples include H2Ö, ÑH3, Cl-, F- etc. They are also considered to be nucleophiles. Any species with an electron-rich centre is said to be a nucleophile.

Brønsted-Lowry Bases
According to Brønsted and Lowry, an acid is a proton donor, while a base is a proton acceptor. In other words, any substance that has the ability to accept a proton (hydrogen ion, H+) by donating a pair of electrons to it, is said to be a base. Examples are H2Ö, ÑH3, Br- etc.

H2Ö(l) + H+(aq) <----> H3Ö+(aq)

ÑH3(l) + H+(aq) ----> NH4+(aq)

Cl-(aq) + H+(aq) ----> HCl(aq)

Arrhenius Bases
Recall that an Arrhenius acid produces hydrogen ion as the only positive ion, when dissolved in water. Similarly, a base is any substance, which produces hydroxide ion, OH- when dissolved in water. For example, sodium hydroxide dissolves in water to produce sodium ion, Na+ and OH- as follows:

NaOH(aq) ----> Na+(aq) + OH-(aq)

Also, potassium hydroxide dissolves in water to produce potassium ion, K+ and OH-.

KOH(aq) ----> K+(aq) + OH-(aq)

A base can also be defined as any substance that neutralizes an acid, or a substance that reacts with an acid to form salt and water. This is because not all bases produce hydroxide ions in water, as some of them are oxides and hydroxides of certain metals, which are insoluble in water.

Classification of Bases

From the aforementioned, bases can be classified into two major categories based on their solubility in water. These are the soluble and insoluble bases.

Soluble Bases: These comprise of oxides of Group I metals and calcium; and hydroxides of sodium, potassium, calcium and ammonium. These soluble hydroxides are known as alkalis.

Insoluble Bases: These are the oxides and hydroxides of other metals not mentioned above.

Examples of alkalis and insoluble bases are:

Alkalis      |        Names
NaOH       |    Sodium hydroxide (Caustic Soda)
KOH          |    Potassium hydroxide (Caustic Potash)
Ca(OH)2   |   Calcium hydroxide (Slaked Lime/Lime Water)
NH4OH     |   Ammonium hydroxide (Liquid Ammonia)

Insoluble Bases    |          Names
          MgO             |    Magnesium oxide
          CuO              |    Copper (II) oxide
          FeO              |     Iron (II) oxide
          Mg(OH)2     |     Magnesium hydroxide
          Fe(OH)2      |     Iron (II) hydroxide
          Cu(OH)2      |    Copper (II) hydroxide
         Zn(OH)2        |     Zinc hydroxide
         Al(OH)3         |     Aluminium hydroxide

Classification of Alkalis

Alkalis are classified based on their strength, which is their degree of ionization in water. These are strong and weak alkalis.

Strong Alkalis: These are those, which ionize completely when dissolved in water. Examples are sodium hydroxide and potassium hydroxide.

NaOH(aq) ----> Na(s) + OH-(aq)

KOH(aq) ----> K+(aq) + OH-(aq)

Weak Alkalis: These are those hydroxides which undergo incomplete dissociation in water. Examples are calcium hydroxide and liquid ammonia.

Ca(OH)2(aq) <----> Ca2+(aq) + 2OH-(aq)

NH4OH(aq) <----> NH4+(aq) + OH-(aq)

NH4OH was was initially known  as ammonium hydroxide, but was later changed to liquid ammonia (NH3.H2O), because unlike other hydroxides, it decomposes in the presence of heat to liberate ammonia gas.

NH4OH(aq) + heat ----> NH3(g) + H2O(l)

Physical Properties of Alkalis

1. They possess a bitter taste.
2. They are slippery to touch because of their soapy feel.
3. They turn red litmus paper blue.
4. They change the colour of phenolphthalein paper from colourless to pink.
5. Their concentrated forms are corrosive.
6. They are electrolytes.

Chemical Properties of Bases

1. Neutralization Reactions: All bases react with acids to form salts and water. This is known as neutralization reaction.

base + acid ----> salt + water

Examples
(a) Potassium hydroxide reacts with hydrochloric acid to form potassium chloride and water.

KOH(aq) + HCl(aq) ----> KCl(aq) + H2O(l)

Ionically,

K+(aq).OH-(aq) + H+(aq).Cl-(aq) ----> K+(aq).Cl-(aq) + H2O(l)

H+(aq) + OH-(aq) ----> H2O(l)

(b) Copper (II) oxide reacts with tetraoxosulphate (VI) acid to produce copper (II) tetraoxosulphate (VI) and water.

CuO(s) + H2SO4(aq) ----> CuSO4(aq) + H2O(l)

(c) Insoluble magnesium hydroxide reacts with dilute trioxonitrate (V) acid to form magnesium trioxonitrate (V) and water.

Mg(OH)2(s) + 2HNO3(aq) ----> Mg(NO3)2(aq) + 2H2O(l)

(d) Sodium hydroxide combines with ethanoic acid to produce sodium ethanoate and water.

NaOH(aq) + CH3COOH(aq) ----> CH3COONa(aq) + H2O(l)

Ionically,

Na+(aq).OH-(aq) + CH3COO-(aq).H+(aq) ----> CH3COO-(aq).Na+(aq) + H2O(l)

H+(aq) + OH-(aq) ----> H2O(l)

From (a) and (d) above, we can see that the actual species that 'partake' in a neutralization reaction are hydrogen/hydroxonium ions, and hydroxide ions.

Hence a neutralization reaction can be redefined as the reaction between hydrogen/hydroxonium ions and hydroxide ions to produce water molecules.

H+(aq) + OH-(aq) ----> H2O(l)

H3O+(aq) + OH-(aq) ----> 2H2O(l)

2. Reaction with Ammonium Salts: All alkalis, except liquid ammonia, react with ammonium salts to liberate ammonia gas when heated. This reaction is used in qualitative analysis for the identification of salts that contain ammonium ions.

NH4+(aq) + NaOH(aq) ----> NH4OH(aq) + Na+(aq)

NH4OH(aq) + heat ----> NH3(g) + H2O(l)

Methods of Preparation of Bases

The following are some of the ways bases can be prepared in the laboratory:

1. Burning metals in air. Metals burn in air to form basic oxides.

4Na(s) + O2(g) ----> 2Na2O(s)

2Ca(s) + O2(g) ----> 2CaO(s)

2Fe(s) + O2(g) ----> 2FeO(s)

2. Dissolving soluble basic oxides in water. Some metallic oxides dissolve in water to form alkalis. For instance, sodium oxide and potassium oxide dissolve in water to form sodium hydroxide and potassium hydroxide respectively.

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

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

Calcium oxide dissolves sparingly in water to form calcium hydroxide (lime water).

CaO(s) + H2O(l) ----> Ca(OH)2(aq)

3. Dissolving some metals in water. The very reactive Group 1 metals like potassium and sodium, and a few Group 2 metals such as calcium and barium react with water to form alkalis with the liberation of hydrogen.

2K(s) + 2H2O(l) ----> 2KOH(aq) + H2(g)

2Na(s) + 2H2O(l) ----> 2NaOH(aq) + H2(g)

Ca(s) + 2H2O(l) ----> Ca(OH)2(aq) + H2(g)

4. Dissolving some metallic hydrides in water. Sodium, potassium and calcium hydrides dissolve in water to form alkalis, with the evolution of hydrogen gas.

NaH(s) + H2O(l) ----> NaOH(aq) + H2(g)

KH(s) + H2O(l) ----> KOH(aq) + H2(g)

CaH2(s) + 2H2O(l) ----> Ca(OH)2(aq) + 2H2(g)

5. Thermal decomposition of some trioxocarbonate (IV) and trioxonitrate (V) salts. Heating some trioxocarbonate (IV) and trioxonitrate (V) salts will produce the oxides of the metallic cations in the salt.

Examples
(a) Copper (II) trioxocarbonate (IV) salt undergoes decomposition when heated to form copper (II) oxide, with the liberation of carbon (IV) oxide.

CuCO3(s) + heat ----> CuO(s) + CO2(g)

Other examples are:

ZnCO3(s) + heat ----> ZnO(s) + CO2(g)

CaCO3(s) + heat ----> CaO(s) + CO2(g)

(b) Calcium trioxonitrate (V) salt decomposes on heating to form calcium oxide. Oxygen and brown fumes of nitrogen (IV) oxide are also given off.

2Ca(NO3)2(s) + heat ----> 2CaO(s) + 4NO2(g) + O2(g)

Other examples include:

2Pb(NO3)2(s) + heat ----> 2PbO(s) + 4NO2(g) + O2(g)

2Zn(NO3)2(s) + heat ----> 2ZnO(s) + 4NO2(g) + O2(g)

6. Double decomposition. Some insoluble hydroxides can be prepared through the process of double decomposition by reacting a soluble salt containing the metallic cation of the intended hydroxide with an alkali.

For instance, copper (II) hydroxide can be prepared by reacting copper (II) trioxonitrate (V) solution with sodium hydroxide.

Cu(NO3)2(aq) + 2NaOH(aq) ----> Cu(OH)2(s) + 2NaNO3(aq)

Uses of Bases

The following are some of the important uses of bases:

1. They are used for making soaps. E.g. NaOH, KOH

2. They are used in the manufacture of fertilizer. E.g. NaOH, liquid ammonia.

3. Calcium hydroxide is used for reducing the level of acidity in acidic soils, and for the production of cement, mortar and Plaster of Paris (POP) etc.

4. Magnesium hydroxide is used for making antacid (milk of magnesia), laxatives and toothpaste.

5. Liquid ammonia is used in laundries as a solvent for removing grease and oil stains.

6.  They are used in the petroleum industry for the refining of some petroleum products. E.g. NaOH

7. They are used as dyes in tanning industry. E.g. KOH

8. Alkalis like KOH, NaOH and liquid ammonia are used in water treatment, and in qualitative analysis for the identification of some cations.

Do These:

Question 1
(a) Elaborate on the statement: "All alkalis are bases, but all bases are not alkalis".
(b) Using appropriate examples, state five methods of preparing bases in the laboratory.
(c) List four uses of bases

Question 2
(a) What is a base?
(b) What is a neutralization reaction?
(c) Aqueous solution of trioxonitrate (V) acid neutralizes a solution of sodium hydroxide. Write the ionic equation for this reaction.
(d) In what other way, can it be shown that potassium hydroxide is a base without using litmus or phenolphthalein paper?

Question 3
Answer True (T) or False (F)
A base
(a) is a proton donor according to Brønsted and Lowry.
(b) produces hydroxide ions in water.
(c) accepts a pair of electrons based on Lewis Theory.
(d) is an electrolyte.
(e) does not have the ability to change the colour of an indicator.
(f) forms an aqueous solution with pH6.
(g) is neutralized by a salt.
(h) reacts with ammonium tetraoxosulphate (VI) to liberate ammonia when heated.

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