Introduction. All known substances are divided into simple and complex. Simple substances are divided into metals and non-metals. Today, the known 105 chemical elements make up more than 30,000,000 compounds, of which about 300,000 form inorganic compounds. Although each compound has its own unique properties, many chemical compounds share similar properties with each other. Based on these general properties, they divide inorganic compounds into oxides, acids, bases and salts [4, p. 98].

 

Oxides. Compounds of oxygen with other elements are called oxides. There are only two kinds of atoms in oxides.

1.      NO₂, SO₂, H₂O, CO₂, N₂O₅, NO, N₂O are common non metal oxides, they have covalent bond structure.

2.      Na₂O, FeO, Al₂O₃, CaO, SiO₂, MgO, CuO, PbO are some common metal oxides they have ionic structure.

 

Some oxides can react directly with water to form an acidic, basic, or amphoteric solution. An amphoteric solution is a substance that can chemically react as either acid or base. However, it is also possible for an oxide to be neither acidic nor basic. There are different properties which help distinguish between the three types of oxides. The term anhydride ("without water") refers to compounds that assimilate H₂O to form either an acid or a base upon the addition of water.

Oxides are binary compounds of oxygen with another element, e.g., CO₂, SO₂, CaO, CO, ZnO, BaO, H₂O, etc. These are termed as oxides because here, oxygen is in combination with only one element. Based on their acid-base characteristics oxides are classified as acidic or basic. An oxide that combines with water to give an acid is termed as an acidic oxide. The oxide that gives a base in water is known as a basic oxide [5, p. 1].

 

Naming of Oxides. They are named like binary compounds.

MgO : Magnesium oxide

FeO : Iron (II) oxide

Na₂O : Sodium oxide

SO₂ : Sulfur dioxide

CO₂ : Carbon dioxide

P₂O₅ : Diphosphorus pentoxide

SnO₂ : Tin(IV) oxide

NO : Nitrogen monoxide

 

Classification of Oxides

A.     Neutral Oxides.

·         They are oxygen poor compounds of non metals.

·         They have neither acidic nor basic properties.

·         They do not react with acids, bases and water. They are slightly soluble in water.

·         CO, NO and N₂O are neutral oxides

 

B.     Acidic Oxides. Oxygen rich compounds of non metals are called acidic oxides. SO₂, NO₂, P₂O₅, Cl₂O₇, CO₂ are examples. Their solutions are acidic.

They are known as acidic anhydrides. Acidic oxide + water → Acid

 

CO₂ + H₂O → H₂CO₃

P₂O₅ + 3H₂O → 2H₃PO₄

N₂O₅ + H₂O → 2HNO₃

 

Acidic oxides are, therefore, known as acid anhydrides, e.g., sulfur dioxide is sulfurous anhydride; sulfur trioxide is sulfuric anhydride. When these oxides combine with bases, they produce salts, e.g.

 

SO₂+ 2NaOH → H₂SO₃ + H₂O

 

C.     Basic Oxides. Generally Group 1 and Group 2 metal oxides are called basic oxides. Na₂O, CaO, Li₂O, MgO, K₂O are examples. Their solutions are basic. They are known as basic anhydrides. Basic oxide + water → Base

 

Na₂O + H₂O → 2NaOH

MgO + H₂O → Mg(OH)₂

K₂O + H₂O → 2KOH

 

These metallic oxides are therefore, known as basic anhydrides. They react with acids to produce salts, e.g.

 

MgO + 2HCl→MgCl₂ + H₂O

Na₂O + H₂SO₄ → Na₂SO₄ + H₂O

 

D.    Mixed Oxides. Compounds that contain two oxides of the same metal are called mixed oxides. Fe₃O₄, Mn₃O₄, Pb₃O₄ are examples.

They behave as if they are two separate oxides in chemical reactions.

 Fe₃O₄ : FeO.Fe₂O₃ : Iron (II, III) oxide

Mn₃O₄ : MnO.Mn₂O₃ : Manganese(II, III) oxide

Pb₃O₄ : 2PbO.PbO₂ : Lead (II, IV) oxide

 

E.     Amphoteric Oxides. Amphoteric oxides are metallic oxides, which show both basic as well as acidic properties. When they react with an acid, they produce salt and water, showing basic properties. While reacting with alkalies they form salt and water showing acidic properties, e.g.

 

ZnO + 2HCl → ZnCl₂ + H₂O (basic nature)

ZnO + 2NaOH → Na₂ZnO₂ + H₂O (acidic nature)

Al₂O₃ + 3H₂SO₄ → Al₂(SO₄)₃ + 3H₂O (basic nature)

Al₂O₃ + 2NaOH → 2NaAlO₂ + H₂O (acidic nature)

 

Amphoteric oxides have both acidic and basic properties. A common example of an amphoteric oxide is aluminum oxide. In general, amphoteric oxides form with metalloids. (see chart below for more detail). Example with acidic properties:

 

Al₂O₃ + H₂O → 2Al(OH)₃ + 2H⁺

 

Example with basic properties:

 

Al₂O₃ + H₂O → 2Al³⁺ + 3OH⁻

 

F.     Peroxides and dioxides. Oxides: Group 1 metals react rapidly with oxygen to produce several different ionic oxides, usually in the form of M₂O. With the oyxgen exhibiting an oxidation number of -2.

 

4Li + O₂ → 2Li₂O

 

Peroxides: Often Lithium and Sodium reacts with excess oxygen to produce the peroxide, M₂O₂. with the oxidation number of the oxygen equal to -1.

Compounds that contain O₂^-2 ion are called peroxides: H₂O₂, K₂O₂, Li₂O₂, CaO₂, MgO₂ are examples.

H₂ + O₂ → H₂O₂

 

Superoxides: Often Potassium, Rubidium, and Cesium react with excess oxygen to produce the superoxide, MO₂ with the oxidation number of the oxygen equal to -1/2.

 

Cs + O₂ → CsO₂

 

A peroxide is a metallic oxide which gives hydrogen peroxide by the action of dilute acids. They contain more oxygen than the corresponding basic oxide, e.g., sodium, calcium and barium peroxides.

 

BaO₂ + H₂SO₄ → BaSO₄ + H₂O₂

Na₂O₂ + H₂SO₄ → Na₂SO₄ + H₂O₂

 

Dioxides like PbO₂ and MnO₂ also contain higher percentage of oxygen like peroxides and have similar molecular formulae. These oxides, however, do not give hydrogen peroxide by action with dilute acids. Dioxides on reaction with concentrated HCl yield Cl₂ and on reacting with concentrated H₂SO₄ yield O₂.

 

PbO₂ + 4HCl → PbCl₂ + Cl₂ + 2H₂O

2PbO₂ + 2H₂SO₄ → 2PbSO₄ + 2H₂O + O₂

 

G.    Compound oxides. Compound oxides are metallic oxides that behave as if they are made up of two oxides, one that has a lower oxidation and one with a higher oxidation of the same metal, e.g.,

                 

Red lead: Pb₃O₄ = PbO₂ + 2PbO

Ferro-ferric oxide: Fe₃O₄= Fe₂O₃ + FeO

 

On treatment with an acid, compound oxides give a mixture of salts.

                    

Fe₃O₄ + 8HCl → 2FeCl₃ + FeCl₂ + 4H₂O [5, p. 1]

 

Preparation of Oxides. Oxides can be generated via multiple reactions. Below are a few.

1.      By direct heating of an element with oxygen Many metals and non-metals burn rapidly when heated in oxygen or air, producing their oxides, e.g.,

 

2Ca + O₂ → 2CaO

P₄ + 5O₂ → P₂O₅

 

2.      By reaction of oxygen with compounds at higher temperatures. At higher temperatures, oxygen also reacts with many compounds forming oxides, e.g. Sulphides are usually oxidized when heated with oxygen.

 

2PbS + 3O₂ → 2PbO + 2SO₂

2ZnS + 3O₂ → 2ZnO + 2SO₂

 

When heated with oxygen, compounds containing carbon and hydrogen are oxidized.

 

C₂H₅OH + 3O₂ → 2CO₂ + 3H₂O

 

By thermal decomposition of certain compounds like hydroxides, carbonates, and nitrates

 

CaCO₃ → CaO + CO₂

Cu(OH)₂ → CuO + H₂O

 

3.      By oxidation of some metals with nitric acid

 

2Cu + 8HNO₃ → 2CuO + 8NO₂ + 4H₂O + O₂

Sn + 4HNO₃  → SnO₂ + 4NO₂ + 2H₂O

 

4.      By oxidation of some non-metals with nitric acid

 

C + 4HNO₃ → CO₂ + 4NO₂ + 2H₂O [5, p. 4]

 

Chemical properties

1.      Interaction with water

Na₂O + H₂O → 2NaOH

SO₃ + H₂O → 2H₂SO₄

 

2.      Interaction with acid or base

 

MgO + H₂SO₄ → MgSO₄ + H₂O

CuO + 2HCl → CuCl₂ + H₂O

CO₂ + Ba(OH)₂ → BaCO₃ + H₂O

SO₂ + 2NaOH → Na₂SO₃ + H₂O

ZnO + 2NaOH → Na₂ZnO₂ + H₂O

ZnO + H₂SO₄ → ZnSO₄ + H₂O

ZnO + 2NaOH + H₂O → Na₂[Zn(OH)₄]

 

3.      Interaction of basic and acid oxide with each other leads to salt formation

 

Na₂O + CO₂ → Na₂CO₃

 

4.      Reduction up to simple substances:

 

3CuO + 2NH₃ → 3Cu + N₂ + 3H₂O

P₂O₅ + 5C → 2P + 5CO

 

 Salts are ionic compounds of anions and cations (figure2.1). NaCl, CaCO₃, ZnBr₂, FeSO₄, NH₄Cl, Mg(NO₃)₂, LiF, AlPO₄ …etc are examples.

 

Figure 2.1 Structure of salt

 

Melting and boiling points: Salts are mostly solids which melt as well as boil at high temperatures.

Solubility in water: Salts are generally soluble in water. For example, sodium chloride, potassium sulphate, aluminium nitrate, ammonium carbonate, etc., are soluble salts while silver chloride, lead chloride, copper carbonate, etc., are insoluble in water.

Water of crystallization: Generally, salts are found as crystals with water molecules present in them. This water is called water of crystallization and such salts are called hydrated salts.

For example, copper sulphate crystal has five molecules of water for each copper sulphate molecule. This is written as CuSO₄∙5H₂O. This water of crystallization gives the crystal its shape. It also gives colour to some crystals. On heating, hydrate salts lose their water of crystallization and, as a result, the crystals lose their shape and colour and change to a powdery substance.

The hydrated salts that have lost their water of crystallization are called anhydrous salts.

When hydrated copper sulphate is heated, it gives out water molecules to form white powdery anhydrous copper sulphate. On addition of water, this substance can convert back to a hydrated copper sulphate solution again [6, p, 1].

 

Naming of Salts

• In the naming of salts first metal ion (positive ion) then name of negative ion is read.

NaCl: Sodium chloride

BaC₂O₄: Barium oxalate

FeCl₂: Iron (II) chloride

KMnO₄: Potassium permanganate

NH₄Br: Ammonium bromide

PbI₂: Lead (II) iodide [6, p. 1]

 

Classification of Salts

A.     Neutral Salts

• They are formed from the reactions of strong acids with strong bases.

• They have neither acidic nor basic character.

• Their solutions are neutral.

• NaCl, LiNO₃, KNO₃, NaBr, Li₂SO₄ …etc are examples.

 

B.     Acidic Salts

• They are formed from the reactions of strong acids with weak bases.

• They have acidic character.

• The salt still has hydrogen atom(s) from an acid which can further be replaced by metallic ions. Examples include: NaHSO₄, NaHCO₃ and NaHS

 

C.     Basic Salt

• They are formed from the reactions of weak acids with strong bases.

• They have basic character.

• The salt contains hydroxides together with metallic ions and negative ions from an acid. Examples are basic zinc chloride, ZnOHCl, basic magnesium chloride.

 

D.    Double salt

Salt that ionizes to produce three different types of ions in solution, two of these are usually positively charged and the other negatively charged. Examples are ammonium iron(II) tetraoxosulphate(VI) hexahydrate, (NH₄)₂ Fe(SO₄)₂ × 6H₂O; potash alum or aluminium potassium tetraoxosulphate(VI) dodecahydrate, KAl(SO₄)₂ × 12H₂O; and chrome alum or chromium(III) potassium tetraoxosulphate(VI) dodecahydrate, KCr(SO₄)₂ × 12H₂O.

 

E.     Complex salt

The salt contains complex ions, i.e. ions consisting of a charged group of atoms. Examples are sodium tetrahydroxozincate(II)

 

Na₂Zn(OH)₄ (aq)   ↔ 2Na⁺ (aq) + Zn(OH)₄^2- (aq) [7, p. 1]

 

Preparation of Salts. Salts are formed by a chemical reaction between:

·         By the reaction between a base and an acid:

 

NH₃ + HCl → NH₄Cl

Cu(OH)₂ + 2HCl → CuCl₂ + 2H₂O

 

·         By the reaction between metal and an acid:

 

Mg + H₂SO₄ → MgSO₄ + H₂

Fe + 2HCl → FeCl₂ + H₂

 

·         By direct union of a metal and a nonmetal: 

 

Ca + Cl₂ → CaCl₂

Fe + S → FeS

 

·         A base and an acid anhydride, e.g.,

 

2 NaOH + Cl₂O → 2 NaClO + H₂O

2NaOH + CrO₃ → Na₂CrO₄ + H₂O

 

·         By the union an acid(or an acidic oxide) and a basic anhydride:

 

2 HNO₃ + Na₂O → 2NaNO₃ + H₂O

3CaO + P₂O₅ → Ca₃(PO₄)₂

 

·         By the reaction between a metal and a base:

 

Zn + 2NaOH → Na₂ZnO₂ + H₂

3S + 6NaOH → 2Na₂S + Na₂SO₃ + 3H₂O

 

·         Salts can also form if solutions of different salts are mixed, their ions recombine, and the new salt is insoluble and precipitates (see: solubility equilibrium), for example:

       

Pb(NO₃)₂ + Na2SO₄ → PbSO₄↓ +2NaNO₃ [8, p. 1]

 

Chemical Properties of Salts

1.      Reaction with a metal: Salts can react with metals according to activity strength.

 

Zn(s) + 2AgNO₃ (aq) → 2Ag(s) + Zn(NO₃)₂ (aq)

 

2.      Reaction with a base: A salt reacts with a base to produce another salt and base.

 

(NH₄)₂SO₄ + 2NaOH → Na₂SO₄ + 2NH₄OH

 

3.      Reaction with an acid: When a salt reacts with an acid, another salt and are formed.

 

NaHCO₃ + H₂SO₄ → NaHSO₄ + H₂CO₃

 

4.      Behaviour of salts towards water: When a salt is dissolved in water, the solution may be neutral, acidic or alkaline. This depends upon the nature of the salt used.

a)      A normal salt derived from a strong base gives an alkaline solution.

 

NaCl + 2H₂O → NaHSO₄ + HCl (at law temperature)

2NaCl + H₂O → Na₂SO₄ + 2HCl (at high temperature)

 

b)      A normal salt derived from a weak acid and strong base gives an alkaline solution.

 

Na₂CO₃ + 2H₂O → 2NaOH + CO₂ + H₂O

CH₃COONa + H₂O → CH₃COOH + NaOH

 

c)      A salt derived from a strong acid and weak base gives an acidic solution.

 

AlCl₃ + 3H₂O → Al(OH)₃ + 3HCl

NH₄Cl + H₂O → NH₄OH + HCl

 

d)     Solution of acidic salts are acidic to litmus, i.e., these solutions turn blue litmus paper red [6, p. 1].

 

Bases. In chemistry, bases are substances that, in aqueous solution, are slippery to the touch, taste astringent, change the color of indicators (e.g., turn red litmus paper blue), react with acids to form salts, promote certain chemical reactions (base catalysis), accept protons from any proton donor, and/or contain completely or partially displaceable OH⁻ ions. Examples of bases are the hydroxides of the alkali metals and the alkaline earth metals (NaOH, Ca(OH)₂, etc.).

Properties. Some general properties of bases include

• Slimy or soapy feel on fingers, due to saponification of the lipids in human skin.
• Concentrated or strong bases are caustic on organic matter and react violently with acidic substances.
• Aqueous solutions or molten bases dissociate in ions and conduct electricity.
• Reactions with indicators: bases turn red litmus paper blue, phenolphthalein pink, keep bromothymol blue in itsnatural colour of blue, and turns methyl orange yellow.
• The pH level of a basic solution is higher than 7.
• Bases are bitter in taste [9, p. 1].

Naming of Bases. The word “hydroxide” is added after the name of metal ion in the naming of bases.

 Mg(OH)₂ : Magnesium hydroxide

KOH : Potassium hydroxide

NaOH : Sodium hydroxide

Ba(OH)₂ : Barium hydroxide

Cu(OH)₂ : Copper (II) hydroxide

NH₃ : Ammonia

 

Classification of Bases

Strong bases. A strong base is a basic chemical compound that is able to deprotonate very weak acids in an acid-base reaction. Compounds with a pKa of more than about 13 are called strong bases. Common examples of strong bases are the hydroxides of alkali metals and alkaline earth metals like NaOH and Ca(OH)₂. Very strong bases are even able to deprotonate very weakly acidic C-H groups in the absence of water. Hydroxide compounds in order of strongest to weakest:

·         Potassium hydroxide (KOH)

·         Barium hydroxide (Ba(OH)₂)

·         Cesium hydroxide (CsOH)

·         Sodium hydroxide (NaOH)

·         Strontium hydroxide (Sr(OH)₂)

·         Calcium hydroxide (Ca(OH)₂)

·         Lithium hydroxide (LiOH)

·         Rubidium hydroxide (RbOH)

 

The cations of these strong bases appear in the 1st and 2nd groups of the periodic table (alkali and earth alkali metals).

Superbases. Group 1 salts of carbanions, amides, and hydrides tend to be even stronger bases due the conjugate acids, which are stable hydrocarbons, amines, and water. Usually these bases are created by adding pure alkali metals such as sodium into the conjugate acid. They are called superbases and it is not possible to keep them in water solution, due to the fact they are stronger bases than the hydroxide ion and as such it will deprotonate the conjugate acid water. For example the ethoxide ion (conjugate base of ethanol) in the presence of water will undergo this reaction.

 

CH₃CH₂O^- + H₂O  → CH₃CH₂OH + OH^-

 

·         Butyl lithium (n-BuLi)

·         Lithium diisopropylamide (C₆H₁₄LiN)

·         Sodium amide (NaNH₂)

·         Sodium hydride (NaH)

 

Weak Base. A "Weak Base" is one that does not fully ionize in solution. When a base ionizes, it takes up a hydrogen ion from the water around it, leaving an OH- ion behind. Weak bases have a higher H⁺ concentration than strong bases. Weak bases exist in chemical equilibrium in the same way weak acids do. The Base Ionization Constant K_b indicates the strength of the base. Large K_bs belong to stronger bases. The pH of a base is greater than 7 (where 7 is the neutral number; below 7 is an acid), normally up to 14. Common example of a weak base is ammonia, which is used for cleaning.

Examples of Weak Bases:

·         Alanine (C₃H₅O₂NH₂)

·         Ammonia (water) (NH₃ (NH₄OH))

·         Dimethylamine ((CH₃)₂NH)

·         Ethylamine (C₂H₅NH₂)

·         Glycine (C₂H₃O₂NH₂)

·         Hydrazine (N₂H₄)

·         Methylamine (CH₃NH₂)

·         Trimethylamine ((CH₃)₃N) [9, p. 1]

 

Preparation of Bases

1.      Reactions of active metals (alkaline and alkaline earth metals) with water

 

2Na + 2H₂O → 2NaOH + H₂

Ca + 2H₂O → Ca(OH)₂ + H₂

 

2.      Interaction oxides of active metals with water

 

BaO + H₂O → Ba(OH)₂

 

3.      Electrolysis water solutions of salts

 

2NaCl + 2H₂O → 2NaOH + H₂ + Cl₂

 

Chemical Properties of Bases. Alkalis Insoluble bases

1.      Action to indicators  litmus – blue methylorange – yellow phenolphthalein – crimson

 

2.      Interaction with acid oxides

 

2KOH + CO₂ → K₂CO₃ + H₂O

KOH + CO₂ → KHCO₃

 

3.      Interaction with acids (reaction of neutralization)

 

NaOH + HNO₃ → NaNO₃ + H₂O

Cu(OH)₂ + 2HCl → CuCl₂ + 2H₂O

 

4.      Reaction of exchange with salts

 

Ba(OH)₂ + K₂SO₄ → 2KOH + BaSO₄↓

3KOH + Fe(NO₃)₃ → Fe(OH)₃↓ + 3KNO₃

 

5.      Thermal decomposition t^оС

 

Cu(OH)₂ → CuO + H₂O

Mg(OH)₂ → MgO + H₂O [10, p. 1]

 

Acids. The term acid was first used in the seventeenth century; it comes from the Latin root ac-, meaning “sharp”, as in acetum, vinegar. Acids have long been recognized as a distinctive class of compounds whose aqueous solutions exhibit the following properties:

·         A characteristic sour taste;

·         ability to change the color of litmus1 from blue to red;

·         react with certain metals to produce gaseous H₂;

·         react with bases to form a salt and water.

The first chemical definition of an acid turned out to be wrong: in 1787, Antoine Lavoisier, as part of his masterful classification of substances, identified the known acids as a separate group of the “complex substances” (compounds). Their special nature, he postulated, derived from the presence of some common element that embodies the “acidity” principle, which he named oxygen, derived from the Greek for “acid former”. Lavoisier had assigned this name to the new gaseous element that Joseph Priestly had discovered a few years earlier as the essential substance that supports combustion. Many combustion products (oxides) do give acidic solutions, and oxygen is in fact present in most acids, so Lavoisier’s mistake is understandable.

In 1811 Humphrey Davy showed that muriatic (hydrochloric) acid (which Lavoisier had regarded as an element) does not contain oxygen, but this merely convinced some that chlorine was not an element but an oxygen-containing compound. Although a dozen oxygen-free acids had been discovered by 1830, it was not until about 1840 that the hydrogen theory of acids became generally accepted. By this time, the misnomer oxygen was too well established a name to be changed.

Compounds dissolving in water by producing H⁺ ion are called acids. Many of the fruits and vegetable contain acids; in lemon, apple, tomatoes, orange, as well as in car batteries, and in cleaning materials.

 

HCl(g) ®  H⁺(aq) + Cl^-(aq)

H₂SO₄ ® 2H⁺(aq) + SO₄^-2(aq)

HNO₃ ®  H⁺(aq) + NO^3-(aq)

CH₃COOH ®  H⁺(aq) + CH₃COO^-(aq)

 

·         They have sour taste.

·         They change the color of litmus paper to red.

·         Their aqueous solutions conduct electricity.

·         They are corrosive substances.

·         Most of them are soluble in water.

 

 Naming of Acids

·         Acids containing two types of atoms are called binary acids.

·         Their names follow the form hydro + nonmetal name + acid.

HCl : Hydrochloric acid

HI : Hydroiodic acid

H₂S : Hydrosulfuric acid

HF : Hydrofluoric acid

 

Acids containing oxygen atoms are called oxy acids. Their names follow the form –ic + acid, or –ous + acid.

H₃BO₃ : Boric acid

H₃PO₄ : Phosphoric acid

H₂SO₄ : Sulfuric acid

H₂SO₃ : Sulfurous acid

HNO₃ : Nitric acid

HNO₂ : Nitrous acid

 

Classification of Acids

According to Strength. If an acid ionizes completely, it is an strong acid, and if it ionizes partially it is a weak acid.

Strong acids: HCl, H₂SO₄, HNO₃, HI, HBr, HClO₄

Weak acid: HF, H₂SO₃, HNO₂, H₂S, H₃PO₄, CH₃COOH, HCN, H₂CO₃

 

According to Number of Hydrogen Atoms. According to number of H+ ion produced acids are classified as monoprotic, diprotic or triprotic.

 

Monoprotic acids: HCl, HNO₃, HI, HBr, HClO₄

 

HCl(g) + H₂O → H⁺(aq) + Cl^-(aq)

CH₃COOH (l) + H₂O → H⁺(aq) + CH₃COO^-(aq)

 

Diprotic acids: H₂SO₃, H₂S, H₂CO₃, H₂SO₄

 

H₂CO₃(l) + H₂O → 2H⁺(aq) + CO₃^-2(aq)

 

Triprotic acids: H₃PO₄, H₃AsO₄

 

H₃AsO₄(s) + H₂O → 3H⁺ (aq) + AsO₄^-3(aq) [11, p. 1]

 

Preparation of Acids

1.      Some acids can be prepared by the direct combination of a non-metal with hydrogen. For e.g.,

 

H₂(g) + Cl₂(g) → 2HCl(g)

H₂(g) + I₂(s) → 2HI(g)

H₂(g) + S(l) → 2H₂S(g)

 

In the above reactions, the obtained gases like hydrogen chloride, hydrogen iodide and hydrogen sulphide are also called the acid anhydrides. They exhibit acidic properties, only when dissolved in water.Therefore an acid anhydride is a substance that dissolves in water to form an acid. In other words, it is the acid minus water (acid - water = acid anhydride).

Examples: CO₂ gas is the acid anhydride of carbonic acid, SO₃ gas is the acid anhydride of carbonic acid, HCl gas is the acid anhydride of hydrochloric acid.

 

H₂O(l) + CO₂(g) → H₂CO₃(aq)

H₂O(l) + SO₃(s) → H₂SO₄(aq)

H₂O(l) + HCl(g) → H₂CO₃(aq)

 

2.      Preparation by dissolving acidic oxide in water. Oxides that can add on hydrogen ions to their molecules are called acidic oxides. Thus when they dissolve in water they associate with hydrogen ions to form acids.

 

6H₂O(l) + 2P₂O₅(s) → 4H₃PO₄(aq)

 

3.      Preparation by the displacement of metal ions of salts of volatile acids by less volatile acid. The hydrogen ions of less volatile acids, like concentrated sulphuric acid, can help displace metal ions of salts of volatile acids, to produce the acid.

 

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

KNO₃(s) + H₂SO₄ → KHSO₄ (aq) + HNO₃ (aq)

 

4.      Preparation by the action of dilute acids on salts

 

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

FeS(s) + H₂SO₄ → FeSO₄ (aq) + H₂O (l)

 

             Hydrogen sulphide is the acid anhydride.

 

5.      Preparation by the oxidation of non-metals by concentrated nitric acid. Non metals like sulphur, phosphorous get oxidized by concentrated nitric acid to form their respective acids.

 

S(s) + 6HNO₃ (aq) → H₂SO₄ (aq) + 6NO₂ (g) + 2H₂O (l)

P(s) + 5HNO₃ (aq) → H₃PO₄ (aq) + 5NO₂ (g) + H₂O (l) [12, p. 1]

 

Chemical Properties of Acids

1.      Action to indicators

Litmus – red

Methylorange – pink

 

2.      Interaction with bases (reaction of neutralization)

 

H₂SO₄ + 2KOH → K₂SO₄ + 2H₂O

2HNO₃ + Ca(OH)₂ → Ca(NO₃)₂ + 2H₂O

 

3.      Interaction with basic oxides

 

CuO + 2HNO₃ → K₂SO₄ + 2H₂O

 

4.      Interaction with metals

 

Zn + 2HCl → ZnCl₂ + H₂

2Al + 6HCl → 2AlCl₃ + 3H₂

 

5.      Interaction with salts (reactions of exchange) at which stands out gas or formed residual

 

H₂SO₄ + BaCl₂ → BaSO₄ + 2HCl

2HCl + K₂CO₃ → 2KCl + H₂O + CO₂ [10, p. 1]

 

Amphoteric Compounds

·         Most of the compounds of Zn, Al, Cr, Sn, Pb, and Be are amphoteric compounds.

·         Oxides and hydroxides of these metals have both acidic and basic characters.

·         They are in soluble in water and do not react with it.

·         ZnO, Al₂O₃, SnO, BeO, Cr₂O₃, PbO are oxides, and Zn(OH)₂, Al(OH)₃, Be(OH)₂, Sn(OH)₂, Pb(OH)₂, Cr(OH)₃ are hydroxides.

 

ZnO + 2HCl → ZnCl₂ + H₂O

ZnO + 2NaOH → Na₂ZnO₂ + 2H₂O