Resonance Pre-foundation Programmes

(PCCP) Division

Career Care

COURSE : OLYMPIAD

CHEMISTRY

WORKSHOP TAPASYA

SHEET

CLASS-VIII

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S. No. Topics Page No.

1.

3 Acids, Bases and Salts 18 27

General Chemistry 1 - 6

2. Metals and Non-Metals 7-17

. -

13RPCCP

SUBJECT : CHEMISTRY CLASS-VIII(OLYMPIAD)

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PAGE # 1

Atom

An atom is the smallest particle of an element thatcan take part in a chemical reaction. The size of anatom is indicated by its radius which is called "atomicradius" (radius of an atom). Atomic radius ismeasured in "nanometres"(nm).

1 metre = 109 nanometres or 1nm = 10-9 m.

Hydrogen atom is the smallest atom of all having an

atomic radius 0.037nm.Maharishi Kanad told that if we keep dividing matteron and on, we will get the smallest particle called asparmanu (which was later termed as atom by JohnDalton).

Modern atomic theory

(a) Structure of an atom :

An atom consists of two parts -

(i) Nucleus : Nucleus is situated in the centre of anatom.All the protons & neutrons are situated in the nucleus,therefore, the entire mass of an atom is almostconcentrated in the nucleus.

GENERAL CHEMISTRY

The overall charge of nucleus is positive due to thepresence of positively charged protons.The protons & neutrons are collectively callednucleons.The radius of the nucleus of an atom is of the order of10�13 cm and its density is of the order of 1015 g/cm3.(ii) Extra nuclear region : In extra nuclear partelectrons are present which revolve around thenucleus in orbits of fixed energies.

The maximum number of electrons that can beaccommodated in a shell is given by the formula 2n2.(n= number of shells i.e. 1,2,3 -------)

Shell n 2n2 max. no.of electrons

K 1 2(1)2 2

L 2 2(2)2 8

M 3 2(3)2 18

N 4 2(4)2 32

(b) Composition of an atom :It consists of three elementary particles electron,proton and neutron. These are known as sub-atomicparticles.

Property Electron Proton Neutron

1. Discovery J.J. Thomson E. Goldstein James Chadwick

2. Symbol e p n

3. Nature Negatively charged Positively charged Neutral

4. Relative charge -1 +1 0

5. Absolute charge 1.602 × 10-19

C 1.602 × 10-19

C 0

6. Relative mass 1 1

7. Absolute mass 9.109 × 10-28

g 1.6725 × 10-24

g 1.6748 × 10-24

g

18371

Atomic number (Z)

Z = no. of protons = no. of electrons (in electricallyneutral atoms).

e.g. 13

Al Z = 13p = 13e = 13

Each element has a unique atomic number.

Mass number (A)

A = no. of protons + no. of neutrons (total no. ofnucleons)A = p + nA = Z + n

e.g. Al2713Z = 13A = 27p = 13e = 13n = A � Z

= 27 � 13 = 14

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PAGE # 2

Symbol

"A symbol is a brief representation of the name of theelement".Berzelius, a Swedish chemist, was the first tointroduce the system of using letters as symbols forthe elements.

(a) Symbols may be derived from the first letter of theEnglish name of the element.e.g. Hydrogen (H), Oxygen (O), Nitrogen (N) etc.

(b) Symbols may be derived from the first letter andanother significant letter of the name of element.e.g. Aluminium (Al), Bromine (Br), Calcium (Ca) etc.

(c) Symbols may be derived from their Latin namesof the elements.e.g. Copper (Cuprum-Cu), Iron (Ferrum-Fe), Silver(Argentum-Ag), Sodium (Natrium- Na), Potassium(Kalium -K),Gold ( Aurum -Au) etc.

Isotopes

The atoms of the same element having same atomicnumber, but different mass numbers are calledisotopes.

orThe atoms of the same element having equalnumber of protons but different number of neutrons.

e.g. Isotopes of HydrogenZ A n

Protium H11 1 1 0

Deuterium H21 1 2 1

Tritium H31 1 3 2

The term isotope was given by Margaret Todd.

(a) Properties of Isotopes :

(i) Chemical properties of all the isotopes of the sameelement are same.

(ii) The electronic configuration of all the isotopes ofsame element is same.

(iii) The physical properties like mass, boiling point,melting point of isotopes of same element aredifferent.

Isobars

The atoms of different elements with different atomicnumbers, but same mass number are called isobars.

e.g. C146

and N14

7are isobars.

Ca40

20 and rA

40

18 are isobars.

Isotones

The isotones may be defined as the atoms of differentelements containing same number of neutrons.

e.g. C13

6 and N147 [Number of neutrons (N) = A - Z]

For C13

6 N = 13 - 6 = 7

For N147 N = 14 - 7 = 7

Other example Si30

14, P3115 and S3216

Isoelectronic

Ion or atom or molecule which have the same

number of electrons are called as isoelectronic

species.

e.g. Cl Ar K

2Ca

No. of electrons 18 18 18 18

(i) The idea of tiniest unit of matter (Anu and Parmanu)

was propounded by Maharishi Kanad in Vedic period

in our country.

(ii) Democritus, a Greek philosopher also proposedthat matter is made up of extremely small particles,the �atom�. The name atom comes from Greeklanguage.

(iii) In 18th century two laws of chemical combinationwere proposed.

(iv) John Dalton in 1808 published theory of atomassuming that atoms are the ultimate indivisibleparticles of matter.

Laws of Chemical Combination

There are two important laws of chemicalcombination.

(a) Law of Conservation of Mass or Matter :

Given by Lavoisier in 1774 . According to the law ofconservation of mass, matter can neither be creatednor be destroyed in a chemical reaction.

OrThe law of conservation of mass means that in achemical reaction, the total mass of products is equalto the total mass of the reactants. There is no changein mass during a chemical reaction.e.g.,

CaCO + Heat3 CaO + CO2

56g 44g

56g + 44g =100g

(100 g)

(b) Law of Constant Proportions / Law of DefiniteProportions :

Given by Proust, in 1799. According to the law of

constant proportions a chemical compound alwaysconsists of the same elements combined togetherin the same proportion by mass.e.g.,Water is a compound of hydrogen and oxygen. 9 partsby weight of water is always found to contain 1 part byweight of hydrogen and 8 parts by weight of oxygen.

PAGE # 3

Dalton's Atomic Theory

(i) Matter consists of small indivisible particles calledatoms.

(ii) All atoms of an element are identical.

(iii) The atoms of an element are different from theatoms of any other element.

(iv) A compound is formed by combination of atomsof two or more elements in simple ratio.e.g. Ratio between H and O in water is 1 : 8 by mass.

(v) Atoms take part in chemical reactions.

(vi) Atoms can neither be created nor be destroyed.

Merits

(i) Dalton�s theory explains the law of conservation ofmass and law of constant proportion.

(ii) Atoms of elements take part in chemical reactionthis is true till today.

Demerits

(i) The atom is no longer supposed to be indivisible.

(ii) He could not explain that why do atoms of sameelement combined with each other.

(iii) Atoms of the same element may not necessarilybe identical in all aspects. e.g,. isotopes.

(iv) Atoms of different elements may not necessarily

be different in all aspects. e.g., isobars.

Atomic Models

(a) Thomson's Model of an Atom : Plum - PuddingModel :

J.J. Thomson proposed this model of the atom in theyear 1903 (then only electrons and protons were knownto be present in the atom).

(i) An atom consists of a sphere (or ball) of positivecharge in which negatively charged electrons areembedded like plums in pudding.(ii) The positive and negative charges in an atom areequal in magnitude, due to which an atom iselectrically neutral.

(b) Rutherford's Nuclear Model of an Atom :

(i) Rutherford's alpha-particle scattering experiment

(1911) :

Rutherford chose an extremely thin gold foil and

bombarded it with alpha particles. The alpha particles

are emitted by radioactive substances. Each alpha

particle has a charge of two protons and their mass

is 4 atomic mass units.

(ii) Observations and conclusions :

(A) Most of the particles passed through the gold foil

along a straight line path hence, there must be large

empty space within the atom.

(B) Some particles were deflected through small

angles and large angles , therefore, there must be

something positively charged present in the atom.

(C) A very small fraction of the particles returned on

their path or bounced back , therefore, there must be

something massive concentrated in a very small

space.

(iii) Rutherford�s nuclear model of an atom :

(A) Most of the space in an atom is empty.

(B) An atom consists of a positively charged, dense

and very small nucleus containing all the protons.

(C) Almost the entire mass of an atom is concentrated

inside the nucleus.

(D) The nucleus is surrounded by negatively charged

electrons. The electrons are revolving around the

nucleus in circular paths at very high speed.

(E) An atom is electrically neutral. This is because

the number of protons and electrons in an atom is

equal.

PAGE # 4

(c) Bohr's Atomic Model (1913) :

(i) The nucleus is situated at the centre of the atom.

(ii) Electrons revolve around the nucleus in certaindefinite circular paths, called orbits or shells.

(iii) In each orbit, electrons possess a certain definiteamount of energy. This energy is different for differentorbits, but it is fixed for any one orbit, hence, theseorbits are also known as energy levels.

(iv) The energy levels are designated as K, L, M, N, O,......................, with K being the nearest to the nucleus.

(v) As long as an electron moves in a particular energylevel, its energy remains constant.

The arrangement of the electrons in different shellsis known as the electronic configuration of theelement.

If the outermost shell has its full quota of 8 electronsit is said to be an octet. If the first shell has its fullquota of 2 electrons, it is said to be duplet.

(a) Bohr -Bury scheme for distribution of electrons in

various shells :

Name of shells K L M NShell number 1 2 3 4

Maximum Number of

electrons (2n2 )2 8 18 32

(b) Electronic Configuration of some Elements -

Atomic number

Symbols of the element

Name of the element

Electronic configuration

1 H Hydrogen 12 He Helium 23 Li Lithium 2,14 Be Beryllium 2,25 B Boron 2,36 C Carbon 2,47 N Nitrogen 2,58 O Oxygen 2,69 F Fluorine 2,710 Ne Neon 2,811 Na Sodium 2,8,112 Mg Magnesium 2,8,213 Al Aluminium 2,8,314 Si Silicon 2,8,415 P Phosphorus 2,8,516 S Sulphur 2,8,617 Cl Chlorine 2,8,718 Ar Argon 2,8,819 K Potassium 2,8,8,120 Ca Calcium 2,8,8,2

(c) Significance of Electronic Configuration :

(i) Electronic configuration of an atom helps us tounderstand the chemical reactivity of the element.

(ii) When the outermost shell of an atom is completelyfilled as per Bohr-Bury scheme then the element isunreactive.

(iii) When the outermost shell of an atom is notcompletely filled according to Bohr-Bury rule, theelement is reactive.

An atom can get the noble gas electronic configurationin three ways -� By losing one or more electrons.� By gaining one or more electrons.� By sharing one or more electrons with other atom oratoms.

Valance shell and valence electrons

The outermost shell of an atom is known as thevalence shell. The electrons present in the valenceshell of an atom are known as valence electrons.The remainder of the atom i.e. the nucleus and otherelectrons is called the core of the atom. Electronspresent in the core of an atom are known as coreelectrons.e.g.The electronic configuration of the sodium (Na) atomis :-Na (11) K L M

2 8 1

Thus, valence electrons in Na atom = 1 and coreelectrons in Na atom = 2 + 8 = 10

(a) Valency :

Valency of an element is the combining capacity ofthe atoms of the element with atoms of the same ordifferent elements.

The valency of an element = number of valenceelectrons(when number of valence electrons are from 1 to 4)The valency of an element = 8� number of valenceelectrons. (when number of valence electrons aremore than 4)

(b) Variable Valency : Certain elements (metals andnon - metals ) exhibit more than one valency.(i) Among the metals iron, copper, silver etc. showvariable valency. For lower valency a suffix -ous andfor higher valency a suffix -ic is attached at the end ofthe name of the metal.e.g.Ferrous = Fe+2

Ferric = Fe+3

(ii) Among the non - metals nitrogen, phosphorus,

sulphur etc. show variable valency.

PAGE # 5

Ions

The charged particles formed by an atom on the gainor loss of one or more electron(s) are called ions.

Cation Anion

Ion

Cation : The loss of an electron by an atom leads tothe formation of a cation.e.g. Na � e� Na+

11e� 10e�

Anion : The gain of an electron by an atom leads toformation of an anion.e.g. Cl + e� Cl�

17e� 18e�

Radicals

A molecule of an inorganic compound is made up oftwo electrically charged species which are known asradicals.The positively charged radical is known as the basicradical while negatively charged radical is calledacidic radical.

(a) Types of radicals :

(i) Simple radicals : When a radical consists of onlyone element , it is called simple radical.e.g. Ag+, Ba2+, Cl�, Br�, S2� etc.

(ii) Compound radicals: When a radical consists ofmore than one type of elements , it is called acompound radical.e.g. NO

3�, SO

42�, NH

4+, BO

33�, PO

43� etc.

LIST OF COMMON ELECTROVALENT POSITIVE RADICALS

LIST OF COMMON ELECTROVALENT NEGATIVE RADICALS

Monovalent Electronegative Bivalent Electronegative

Trivalent Electronegative

Tetravalent Electronegative

1. Fluoride F�

1. Sulphate SO 42- 1. Nitride N3- 1. Carbide C4-

2. Chloride Cl� 2. Sulphite SO 32- 2. Phosphide P3-

3. Bromide Br� 3. Sulphide S2-3. Phosphite PO3

3-

4. Iodide I� 4. Thiosulphate S2O32- 4. Phosphate PO4

3-

5. Hydride H� 5. Zincate ZnO22-

6. Hydroxide OH� 6. Oxide O2-

7. Nitrite NO2� 7. Peroxide O2

2-

8.Nitrate NO3� 8. Dichromate Cr2O7

2-

9. Bicarbonate or Hydrogen carbonate HCO3� 9. Carbonate CO3

2-

10. Bisulphite or Hydrogen sulphite HSO3� 10. Silicate SiO 3

2-

11. Bisulphide or Hydrogen sulphide HS�

12. Bisulphate or Hydrogen sulphate HSO4�

13. Acetate CH COO3�

PAGE # 6

Chemical formula

Molecule of an element or a compound may berepresented by symbols of the elements present inone molecule of the compound. It is known as achemical formula.

e.g. HCl is the formula of hydrogen chloride and NaCl isthat of sodium chloride.

(a) Significance of a Chemical Formula :

(i) Names of the elements present in the compound.

(ii) Number of atoms of each element.

(iii) Molecular weight of the compound.

(iv) The relative proportion of weights of the elements.

(b) Writing a Chemical Formula :

Step-1Write the symbol of the positive ion or the radical tothe left and that of the negative ion or radical to theright.

Step-2Put the valency of each radical or the ion on its topright. Divide the valency by the highest common factor,if any, to get a simple ratio. Now ignore the (+) & (�)signs. Interchange the valency of radicals or ions.

Step-3Shift the valency to the lower right side of the radicalor ion. If the compound radical receives number morethan 1, enclose it within brackets. Do not enclosesimple radicals within brackets.

Names of compounds

y y

(c) Formula of some useful compounds

S.No. Compounds Common Names Chemical Names

1 CaO Lime Calcium oxide

2 NaHCO3 Baking soda Sodium hydrogen carbonate

3 Na2CO3.10H2O Washing soda Sodium carbonate decahydrate

4 CaCO3 Limestone Calcium carbonate

5 Ca(OH)2 Slaked lime Calcium hydroxide

6 CuSO4. 5H2O Blue vitriol Copper sulphate pentahydrate

7 NaCl Common salt Sodium chloride

8 Na2CO3 Soda ash Sodium carbonate

9 NaOH Caustic soda Sodium hydroxide

10 KOH Caustic potash Potassium hydroxide

11 CaOCl2 Bleaching powder Calcium oxychloride

12 CaSO4.1/2 H2O Plaster of paris Calcium sulphate hemihydrate

13 CaSO4. 2H2O Gypsum Calcium sulphate dihydrate

14 FeSO4. 7H2O Green vitriol Ferrous sulphate heptahydrate

15 H2SO4 Oil of vitriol Sulphuric acid

Formula of some simple useful compounds

7PAGE # 7

METALS AND NON-METALS

INTRODUCTION

There are 118 chemical elements known at present.On the basis of their properties, all these elementscan be broadly divided into two main groups: Metalsand Non-Metals. A majority of the known elementsare metals. All the metals are solids, except mercury,which is a liquid metal at room temperature. Thereare 22 non-metals, out of which, 10 non-metals aresolids, one non-metal (bromine) is a liquid and theremaining 11 non-metals are gases at roomtemperature.

POSITION OF METALS AND NON-METALS

IN THE PERIODIC TABLE

The metals are placed on the left hand side and inthe centre of the periodic table. On the other hand, thenon-metals are placed on the right hand side of theperiodic table. This has been shown in the figure. Itmay be noted that hydrogen (H) is an exceptionbecause it is non-metal but is placed on the left handside of the periodic table.

Metals and non-metals are separated from each otherin the periodic table by a zig-zag line. The elementsclose to zig-zag line show properties of both themetals and the non-metals. They show someproperties of metals and some properties of non-metals. These are called metalloids. The commonexamples of metalloids are boron (B), silicon (Si),germanium (Ge), arsenic (As), antimony (Sb), tellurium(Te) and polonium (Po).

In general, the metallic character decreases on goingfrom left to right side in the periodic table. However,on going down the group, the metallic characterincreases.

Note :The elements at the extreme left of the periodic tableare most metallic and those on the right are leastmetallic or non-metallic.

GENERAL PROPERTIES OF METALS

AND NON-METALS

(a) Electronic Configuration of Metals :

The atoms of metals have 1 to 3 electrons in theiroutermost shells. For example, all the alkali metalshave one electron in their outermost shells (lithium -2, 1 ; sodium-2, 8, 1; potassium-2, 8, 8, 1 etc.).

Sodium, magnesium and aluminium are metalshaving 1, 2 and 3 electrons respectively in theirvalence shells. Similarly, other metals have 1 to 3electrons in their outermost shells.

It may be noted that hydrogen and helium areexception because hydrogen is a non-metal havingonly 1 electron in the outermost shell (K shell) of itsatom and helium is also a non-metal having 2electrons in the outermost shell (K shell).

(b) Physical Properties of Metals :

The important physical properties of metals arediscussed below :

(i) Metals are solids at room temperature: All metals(except mercury) are solids at room temperature.

Note :Mercury is a liquid at room temperature.

(ii) Metals are malleable :Metals are generally malleable. Malleability meansthat the metals can be beaten with a hammer intovery thin sheets without breaking. Gold and silver areamong the best malleable metals. Aluminium andcopper are also highly malleable metals.

(iii) Metals are ductile :It means that metals can be drawn (stretched) intothin wires. Gold and silver are the most ductile metals.Copper and aluminium are also very ductile andtherefore, these can be drawn into thin wires whichare used in electrical wiring.

(iv) Metals are good conductors of heat andelectricity :All metals are good conductors of heat. The conductionof heat is called thermal conductivity. Silver is the bestconductor of heat. Copper and aluminium are alsogood conductors of heat and therefore, they are usedfor making household utensils. Lead is the poorestconductor of heat. Mercury metal is also a goodconductor of heat.

Metals are also good conductors of electricity. Theelectrical and thermal conductivities of metals aredue to the presence of free electrons in them. Amongall the metals, silver is the best conductor of electricity.Copper and aluminium are the next best conductorsof electricity. Since silver is expensive, therefore,copper and aluminium are commonly used formaking electric wires.

Note :Silver is best conductor of heat and electricity.

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8PAGE # 8

(v) Metals are lustrous and can be polished :

Most of the metals have shine and they can be polished.The shining appearance of metals is also known asmetallic lustre. For example, gold, silver and copper metalshave metallic lustre.

Activity :To demonstrate action of air, moisture and othergases on the metal surface.Take a piece of Magnesium ribbon, Aluminium wireand a sheet of Copper and note their appearance. Nowrub the surface of each metal with a sand paper andobserve their appearance again. You will observe thatthe metal articles, which had a dull appearance,become bright on rubbing. This can be explained asfollows :

When a metal has been kept exposed to air for a longtime, its appearance becomes dull as it loses most ofits shine or brightness. This is due to formation of athin layer of oxide, hydroxide, carbonate or sulphide ofthe metal by slow action of the moisture and gasespresent in the air with the metal. This process ofdeposition of a layer of metal oxides or other such metalcompounds is termed as corrosion of metal. Forexample, the surface of aluminium on exposure to airis covered with a thin layer of aluminium oxide whichprevents further reaction between aluminiumunderneath and air. Similarly, copper is coated with agreen layer when kept in moist air due to formation ofBasic copper carbonate while silver articles acquireblackish colour due to formation of silver sulphide onits surface.

On rubbing the dull metal surface with a sand paper,the outer corroded layer can be removed and the metalsurface again becomes lustrous and bright.

(vi) Metals have high densities :

Most of the metals are heavy and have high densities.For example, the density of mercury metal is very high(13.6 g cm�3). However, there are some exceptions.Sodium, potassium, magnesium and aluminium havelow densities. Densities of metals are generallyproportional to their atomic masses.

(vii) Metals are hard :Most of the metals are hard. But all metals are notequally hard. Metals like iron, copper, aluminium etc.are quite hard. They cannot be cut with a knife. Sodiumand potassium are common exceptions which are softand can be easily cut with a knife.

(viii) Metals have high melting and boiling points :Most of the metals (except Na, K, Rb, Cs, Ga) havehigh melting and boiling points.

Note :Tungsten has highest melting point (34100C) amongall the metals.

(ix) Metals are rigid : Most of the metals are rigid andthey have high tensile strength.

(x) Metals are sonorous : Most of the metals aresonorous i.e, they make sound when hit with an object.

(c) Electronic Configuration of Non-Metals :

The atoms of non-metals have usually 4 to 8 electronsin their outermost shells. For example, Carbon (At. No.6), Nitrogen (At. No. 7), Oxygen (At. No. 8), Fluorine (At.No. 9) and Neon (At. No. 10) have respectively 4,5,6,7,8electrons in their outermost shells.

(d) Physical Properties of Non-Metals :

The important physical properties of non-metals arelisted below:(i) Non-metals are brittle.

(ii) Non-metals are not ductile.

(iii) Non-metals are bad conductor of heat andelectricity. (Exception: Graphite is a good conductorbecause of the presence of free electrons.)

(iv) Non-metals are not lustrous and cannot bepolished. (Exception: Graphite and Iodine are lustrousnon-metals.)

(v) Non-metals may be solid, liquid, or gases at roomtemperature.

(vi) Non-metals are generally soft. (Exception:Diamond, an allotropic form of non-metal Carbon, isthe hardest natural substance known).

(vii) Non-metals have generally low melting and boilingpoints. (Exception: Graphite, another allotropic form ofCarbon, has a melting point of about 3730°C).

(viii) Non-metals have low densities. (Exception: Iodinehas high density).

Note :

Graphite is a good conductor of electricity, lustrous andhas very high melting point.

(e) Chemical Properties of Metals :

The atoms of the metals have usually 1, 2 or 3electrons in their outermost shells. These outermostelectrons are loosely held by their nuclei. Therefore,the metal atoms can easily lose their outermostelectrons to form positively charged ions. For example,sodium metal can lose outermost one electron to formpositively charged ion, Na+. After losing the outermostelectron, it gets stable electronic configuration of thenoble gas (Ne : 2, 8). Similarly, magnesium can losetwo outermost electrons to form Mg2+ ion andaluminium can lose its three outermost electrons toform Al3+ ion.Na Na+ + e�

(2, 8, 1) (2, 8)Mg Mg2+ + 2e�

(2, 8, 2) (2, 8)Al Al3+ + 3e�

(2, 8, 3) (2, 8)

Note :The metal atoms lose electrons and form positivelycharged ions, therefore, the metals are calledelectropositive elements.

9PAGE # 9

Difference between the physical properties of metals & nonmetals

10PAGE # 10

Some of the important chemical properties of metals arediscussed below :

(i) Reaction with oxygen :

Metals react with oxygen to form oxides. These oxidesare basic in nature. For example, sodium metal reactswith oxygen of the air and form sodium oxide.

4 Na(s) + O2(g) 2 Na

2O(s)

Sodium oxide

Sodium oxide reacts with water to form an alkali calledsodium hydroxide. Therefore, sodium oxide is a basicoxide.Na

2O(s) + H

2O() 2NaOH(aq)

Sodium hydroxideDue to the formation of sodium hydroxide (which is analkali), the solution of sodium oxide in water turns redlitmus blue (common property of all alkaline solutions).

Note :

When metal oxides are dissolved in water, they givealkaline solutions.Similarly, magnesium is a metal and it reacts with oxygento form magnesium oxide. However, magnesium is lessreactive than sodium and therefore, heat is requiredfor the reaction.

2Mg(s) + O2(g) 2 MgO(s)

Thus, when a metal combines with oxygen, it loses itsvalence electrons and forms positively charged metalion. We can say that oxidation of metal takes place.

Reactivity of metals towards oxygen :

All metals do not react with oxygen with equal ease.The reactivity of oxygen depends upon the nature ofthe metal. Some metals react with oxygen even at roomtemperature, some react on heating while still othersreact only on strong heating. For example :

(A) Metals like sodium, potassium and calcium reactwith oxygen even at room temperature to form theiroxides.

4Na(s) + O (g) 2Na O(s)2 2

Sodium Oxygen Sodium oxide

4K(s) + O2(g) 2K

2O(s)

Potassium Oxygen Potassium oxide

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

Calcium Oxygen Calcium oxide

(B) Metals like magnesium do not react with oxygen atroom temperature. They burn in air on heating to formcorresponding oxides.

2Mg(s) + O2(g) 2MgO(s)

Magnesium Oxygen Magnesium oxide

(C) Metals like zinc do not react with oxygen at roomtemperature. They burn in air only on strong heatingto form corresponding oxides.

2 Zn(s) + O2 (g) 2 ZnO(s)

Zinc Oxygen Zinc oxide

(D) Metals like iron and copper do not burn in air evenon strong heating. However, they react with oxygen onlyon prolonged heating.

3Fe(s) + 2O2(g) Fe

3O

4(s)

Iron Oxygen Iron (II, III) oxide

2Cu(s) + O2(g) 2CuO(s)

Copper Oxygen Copper (II) oxide

(ii) Reaction with water :

Metals react with water to form metal oxide or metalhydroxide and hydrogen. The reactivity of metalstowards water depends upon the nature of the metals.Some metals react even with cold water, some reactwith water only on heating while there are some metalswhich do not react even with steam. For example,

(A) Sodium and potassium metals react vigorouslywith cold water to form their respective hydroxides andhydrogen gas is liberated.

2 Na(s) + 2H2O() 2NaOH (aq) + H

2(g)

Sodium Cold water Sodium Hydrogen hydroxide

2K (s) + 2H2O() 2KOH (aq) + H

2 (g)

Potassium Cold water Potassium Hydrogenhydroxide

Note :

The reaction between sodium and water is so violent

that the hydrogen evolved catches fire.

(B) Calcium reacts with cold water to form calcium

hydroxide and hydrogen gas. The reaction is less

violent.

Ca(s) + 2H2O() Ca (OH)

2 (aq) + H

2(g)

Calcium Cold water Calcium hydroxide

(C) Magnesium reacts very slowly with cold water but

reacts rapidly with hot boiling water forming

magnesium oxide and hydrogen.

Mg (s) + H2O () MgO(s) + H

2(g)

Magnesium Boiling Magnesium

water oxide

(D) Metals like zinc and aluminium react only with steam

to form their respective oxides and hydrogen.

Zn (s) + H2O(g) ZnO (s) + H

2(g)

Zinc Steam Zinc oxide

2Al (s) + 3H2O (g) Al

2O

3(s) + 3H

2(g)

Aluminium Steam Aluminium

oxide

(E) Iron metal does not react with water under ordinary

conditions. The reaction occurs only when steam is

passed over red hot iron and the products are iron

(II,III) oxide and hydrogen.

3Fe(s) + 4H2O(g) Fe

3O

4(s) + 4H

2(g)

Iron Steam Iron (II,III) Hydrogen

(Red hot) oxide

11PAGE # 11

(F) Metals like copper, silver and gold do not react withwater even under strong conditions. The order ofreactivities of different metals with water is :Na > Mg > Zn > Fe > CuReactivity with water decreases

(iii) Reaction with dilute acids :Many metals react with dilute acids and liberatehydrogen gas. Only less reactive metals such ascopper, silver, gold etc. do not liberate hydrogen fromdilute acids. The reactions of metals with dilutehydrochloric acid (HCl) and dilute sulphuric acid(H

2SO

4) are similar. With dil. HCl, they give metal

chlorides and hydrogen whereas with dil. H2SO

4, they

give metal sulphates and hydrogen.

Note :

Dilute nitric acid (HNO3) is an oxidising agent which

oxidises metals,but does not produce hydrogen.

But Mg & Mn produce hydrogen on reacting with diluteHNO

3.

The reactivity of different metals is different with thesame acid. For example :

(A) Sodium, magnesium and calcium react violentlywith dilute hydrochloric acid (HCl) or dilute sulphuricacid (H

2SO

4) liberating hydrogen gas and

corresponding metal salt.

2Na(s) + 2HCl (aq) 2NaCl(aq) + H2(g)

Sodium Hydrochloric Sodium Hydrogen acid chloride

2Na(s) + H2SO

4(aq) Na

2SO

4 (aq) + H

2(g)

Sodium Sulphuric Sodium Hydrogen acid sulphate

Similarly,Mg (s) + 2HCl (aq) MgCl

2(aq) + H

2(g)

Magnesium Hydrochloric Magnesium Hydrogen acid chloride

Mg(s) + H

2SO

4 (aq) MgSO

4 (aq) + H

2(g)

Magnesium Sulphuric Magnesium Hydrogenacid sulphate

(B) Aluminium and zinc react with dilute hydrochloricacid (HCl) or dilute sulphuric acid (H

2SO

4) liberating

hydrogen gas and corresponding metal salt.2Al(s) + 6HCl (aq) 2AlCl

3(aq) + 3H

2(g)

Aluminium Hydrochloric Aluminium Hydrogen acid chloride

2Al(s) + 3H2SO

4(aq) Al

2(SO

4)

3(aq)+ 3H

2(g)

Aluminium Sulphuric Aluminium Hydrogen acid sulphate

Zn(s) + 2HCl (aq) ZnCl2(aq) + H

2(g)

Zinc Hydrochloric Zinc Hydrogenacid chloride

Zn(s) + H2SO

4(aq) ZnSO

4 (aq) + H

2(g)

Zinc Sulphuric Zinc Hydrogenacid sulphate

(C) Iron reacts slowly with dilute HCl or dil. H2SO

4 and

therefore, it is less reactive than zinc and aluminium.Fe(s) + 2HCl(aq) FeCl

2(aq) + H

2(g)

Iron Hydrochloric Ferrous Hydrogen acid chloride

Fe(s) + H2SO

4(aq) FeSO

4(aq) + H

2(g)

Iron Sulphuric Ferrous Hydrogen acid sulphate

(C) Copper does not react with dil. HCl or dil H2SO

4 .

Cu(s) + HCl (aq) No reactionCu(s) + H

2SO

4(aq) No reaction

Therefore copper is even less reactive than iron.

The order of reactivity of different metals with diluteacid:

Na > Mg > Al > Zn > Fe > CuReactivity with dilute acids decreases from sodium tocopper.(iv) Reactions of metals with salt solutions :When a more reactive metal is placed in a salt solutionof less reactive metal, then the more reactive metaldisplaces the less reactive metal from its salt solution.For example, we will take a solution of copper sulphate(blue coloured solution) and put a strip of zinc metal inthe solution. It is observed that the blue colour of coppersulphate fades gradually and copper metal is depositedon the zinc strip. This means that the following reactionoccurs :

Zn(s) + CuSO4(aq) ZnSO

4(aq) + Cu(s)

Zinc Copper Zinc sulphate Copper sulphate (Colourless solution)(Blue solution)

Here, zinc displaces copper from its salt solution.However, if we take zinc sulphate solution and put astrip of copper metal in this solution, no reaction occurs.ZnSO

4 (aq) + Cu(s) No reaction

Zinc Coppersulphate

This means that copper cannot displace zinc metalfrom its solution. Thus, we can conclude that zinc ismore reactive than copper. However, if we put gold orplatinum strip in the copper sulphate solution, thencopper is not displaced by gold or platinum. Thus, goldand platinum are less reactive than copper.

REACTIVITY SERIES OF METALS

(a) Introduction :

We have learnt that some metals are chemically veryreactive while others are less reactive or do not react at all.

On the basis of reactivity of different metals with oxygen,water and acids as well as displacement reactions,the metals have been arranged in the decreasing orderof their reactivities.The arrangement of metals in order of decreasingreactivities is called reactivity series or activity seriesof metals.The activity series of some common metals is given in

Table. In this table, the most reactive metal is placed at

the top whereas the least reactive metal is placed at

the bottom. As we go down the series the chemical

reactivity of metals decreases.

12PAGE # 12

REACTIVITY SERIES OF METALS

Potassium K

Barium Ba

Sodium Na

Calcium Ca

Magnesium Mg

Aluminium Al

Zinc Zn

Iron Fe

Nickel Ni

Tin Sn

Lead Pb

Hydrogen H

Copper Cu

Mercury Hg

Silver Ag

Gold Au

Platinum Pt

Lithium Li Most reactive metal

Least reactive metal

Metalsmore reactivethan hydrogen

Metalsless reactive than hydrogen

Reactivity decreases

Rea

ctiv

ity in

crea

ses

(b) Reasons for Different Reactivities :

In the activity series of metals, the basis of reactivity isthe tendency of metals to lose electrons. If a metal canlose electrons easily to form positive ions, it will reactreadily with other substances. Therefore, it will be areactive metal. On the other hand, if a metal loseselectrons less readily to form a positive ion, it will reactslowly with the other substances. Therefore, such ametal will be less reactive. For example, alkali metalssuch as sodium and potassium lose electrons veryreadily to form alkali metal ions, therefore, they arevery reactive.

(c) Displacement of Hydrogen from Acids

by Metals :

All metals above hydrogen in the reactivity series (i.e.more active than hydrogen) like zinc, magnesium, nickeletc can liberate hydrogen from acids like HCl and H

2SO

4.

These metals have greater tendency to lose electronsthan hydrogen. Therefore, the H+ ions in the acids willaccept electrons and give hydrogen gas as :M M+ (aq) + e�

Metal

H+ (aq) + e� H(From acid)

H + H 2H

The metals which are below hydrogen in the reactivityseries (i.e. less reactive than hydrogen) like copper,silver, gold cannot liberate hydrogen form acids likeHCl, H

2SO

4 etc. These metals have lesser tendency to

lose electrons than hydrogen. Therefore, they cannotgive electrons to H+ ions.

(d) Reactivity Series and Displacement

Reactions :

The reactivity series can also explain displacementreactions. In general, a more reactive metal (placedhigher in the activity series) can displace the lessreactive metal from its salt solution. For example, zinc,displaces copper from its salt solution.Zn (s) + CuSO

4 (aq) ZnSO

4(aq) + Cu(s)

(e) Usefulness of Activity Series :

The activity series is very useful and it gives thefollowing informations :

(i) The metal which is higher in the activity series ismore reactive than the others. Lithium is the mostreactive and platinum is the least reactive metal.

(ii) The metals which have been placed above hydrogenare more reactive than hydrogen and these candisplace hydrogen from its compounds like water andacids to liberate hydrogen gas.

(iii) The metals which are placed below hydrogen areless reactive than hydrogen and these cannot displacehydrogen from its compounds like water and acids.

(iv) A more reactive metal (placed higher in the activityseries) can displace the less reactive metal from its saltsolution.

(v) Metals at the top of the series are very reactive and,therefore, they do not occur free in nature. The metalsat the bottom of the series are least reactive and,therefore, they normally occur free in nature. Forexample, gold, present in the reactivity series is foundin free state in nature.

13PAGE # 13

CHEMICAL PROPERTIES OF NON-METALS

Non-metals have usually 4 to 8 electrons in theoutermost shells. They have the tendency to acceptelectrons to complete their octets. By accepting theelectrons, they form negatively charged ions and,therefore, they are electronegative elements. Forexample, nitrogen, oxygen and fluorine can accept 3, 2and 1 electrons respectively to complete their octetsas :

N + 3e� N3�gains 3 electrons

O + 2e� O2�gains 2 electrons

F + e� F� gains 1 electron

Some of the important properties of non-metals arediscussed below :

(a) Reaction of non-metals with oxygen :

Non-metals react with oxygen to form acidic or neutraloxides. These oxides are covalent in nature and areformed by sharing of electrons. The acidic oxidesdissolve in water to give acids.

Acidic oxides

The oxides of carbon, sulphur, phosphorus etc., areacidic and therefore, they turn blue litmus solution red.For example :

(i) Carbon reacts with oxygen of air to form carbondioxide.

C (s) + O2 (g) CO2(g)

Carbon Carbon dioxide

Carbon dioxide dissolves in water to form an acid calledcarbonic acid.

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

Carbonic acid

(an acid)

(ii) Sulphur burns in air to give sulphur dioxide.

S (s) + O2 (g) SO2 (g)

Sulphur dioxide

Sulphur dioxide dissolves in water to form an acid calledsulphurous acid.

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

Sulphurous acid

Sulphur also forms an oxide, sulphur trioxide, whichdissolves in water to give sulphuric acid.

(iii) When phosphorus is burnt in air, it reacts withoxygen of air to form phosphorus pentoxide (P2O5).This is also an acidic oxide and dissolves in water togive an acid, phosphoric acid.

P4 (s) + 5O2 (g) 2P2O5 (s)

Phosphorous pentoxide

P2O5 (s) + 3H2O (l) 2H3PO4 (aq)

Phosphoric acid

Phosphorus also forms an oxide, phosphorus trioxide,which dissolves in water to give phosphorous acid.

P4 (s) + 3O2 (g) 2P2O3 (s)

Phosphorous trioxide

P2O3 (s) + 3H2O (l) 2H3PO3 (aq)

Phosphorous acid

The acidic oxides of non-metals neutralise bases toform salt and water. For example,

SO2 + 2NaOH Na2SO3 + H2O

Sodium sulphite

Neutral Oxides

Some oxides of non-metals are neutral. For example,carbon monoxide (CO), nitric oxide (NO), nitrous oxide(N2O), water (H2O) etc.

For example :

2C (s) + O2 (g) 2CO (g)

Carbon monoxide

(Neutral oxide)

2H2 (g) + O2 (g) 2H2O (l)

Water

(Neutral oxide)

These oxides do not turn blue litmus solution red.

(b) Reaction with hydrogen :

Non-metals react with hydrogen under differentconditions to form corresponding covalent hydrides.For example, H2O, H2S, NH3, HCl, CH4 etc., arecommon hydrides of oxygen, sulphur, nitrogen, chlorineand carbon respectively.

2H2 (g) + O2 (g) sparkElectric 2H2O (l)

H2 (g) + S (s) K 715 H2S (g)

N2 (g) + 3H2 (g)atm200

K775,Mo/Fe 2NH3 (g)

H2 (g) + Cl2 (g) sunlightDiffused 2HCl (g)

(c) Reaction of non-metals with water :

Non-metals do not react with water or steam to givehydrogen gas. This is because non-metals cannot giveelectrons to reduce the hydrogen ions of water intohydrogen gas.

(d) Reaction with acids :

Non-metals do not react with dilute acids and therefore,hydrogen gas is not liberated when non-metals aretreated with dilute acids. Therefore, non-metals do notdisplace hydrogen from dilute acids. For example,carbon, sulphur or phosphorus do not react with diluteacids such as dil HCl or dil H2SO4 to produce hydrogengas. We have seen that hydrogen can only be displacedfrom dilute acids if electrons are supplied to H+ ions ofthe acids.

H2SO4 (aq) 2H+ (aq) + SO42� (aq)

2H+ (aq) + 2e� H2 (g)

But the non-metals are electron acceptors and,

therefore, they cannot give electron to H+ ions of an

acid. Hence, hydrogen gas is not liberated.

14PAGE # 14

(e) Reaction with chlorine :

Non-metals react with chlorine to form covalent

chlorides such as HCl, PCl3, CCl4 etc. For example,

H2 (g) + Cl2 (g) sunlightDiffused 2HCl (g)

Hydrogen chloride

(f) Reaction with salt solution :

, A more reactive non-metal displaces a less reactivenon-metal from its salt. For example, when chlorine is

passed through a solution of sodium bromide, thenbromine is liberated.

2NaBr (aq) + Cl2 2NaCl (aq) + Br2 Sodium chloride

However, bromine cannot displace chlorine from itssalt solution.

2NaCl (aq) + Br2 No reaction.

Therefore, chlorine is a more reactive non-metal thanbromine.

Comparison of reactivities of halogens is : F > Cl > Br > I

Comparison of the Chemical Properties of Metals & Non Metals

Chemical Properties Metals Non-Metals

1. Nature of oxides Metals form basic oxides, someare amphoteric also.

Non-metals form acidic or neutral oxides.

2. Displacement of hydrogen from acids

Metals displace hydrogen from acids and form salts.

Non-metals do not displace hydrogen from acids.

3. Reaction with chlorine Metals react with Cl2 to formelectrovalent chlorides.

Non-metals react with Cl2

to form covalent chlorides.4. Reaction with hydrogen With hydrogen, only a few

metals combine to form electrovalent hydrides.

With hydrogen, non-metals form many stable hydrides which are covalent.

5. Electropositive or electronegative character

Metals are electropositive in nature.

Non-metals are electronegative in nature.

6. Oxidising and reducing agent character

Metals act as reducing agents. Non-metals act as oxidising agents.

OCCURRENCE OF METALS

All metals are present in the earth�s crust either in the

free state or in the form of their compounds. Aluminiumis the most abundant metal in the earth�s crust. The

second most abundant metal is iron and third one iscalcium.

(a) Native and Combined States of Metals :

Metals occur in the crust of earth in the following twostates -

(i) Native state or free state : A metal is said to occur ina free or a native state when it is found in the crust ofthe earth in the elementary or uncombined form.

The metals which are very unreactive (lying at thebottom of activity series) are found in the free state.These have no tendency to react with oxygen and arenot attacked by moisture, carbon dioxide of air or othernon-metals. Silver, copper, gold and platinum aresome examples of such metals.

(ii) Combined state : A metal is said to occur in acombined state if it is found in nature in the form of itscompounds. e.g. Sodium , magnesium etc.

Note :Copper and silver are metals which occur in the freestate as well as in the combined state.

MINERALS AND ORES

The natural substances in which metals or theircompounds occur either in native state or combinedstate are called minerals.The minerals are not pure and contain different typesof other impurities. The impurities associated withminerals are collectively known as gangue or matrix.The mineral from which the metal can be convenientlyand profitably extracted, is called an ore.For example, aluminium occurs in the earth�s crust in

the form of two minerals, bauxite (Al2O

3.2H

2O) and clay

(Al2O

3.2SiO

2.2H

2O). Out of these two, aluminium can

be conveniently and profitably extracted from bauxite .So, bauxite is an ore of aluminium.

Note : Oxygen is the most abundant element in earth�scrust.

(a) Types of Ores :

The most common ores of metals are oxides,sulphides, carbonates, sulphates, halides, etc. Ingeneral, very unreactive metals (such as gold, silver,platinum etc.) occur in elemental form or free state.

(i)Metals which are only slightly reactive occur assulphides (e.g., CuS, PbS etc.).

(ii)Reactive metals occur as oxides (e.g., MnO2, Al

2O

3

(iii)Most reactive metals occur as salts as carbonates,sulphates, halides etc. (e.g., Ca, Mg, K etc.).

15PAGE # 15

Some common ores are listed in the table

Nature of ore

Metal Name of the ore Composition

Aluminium Bauxite Al2O3.2H2O

Copper Cuprite Cu2O

Magnetite Fe3O4

Haematite Fe2O3

Copper pyrites CuFeS2

Copper glance Cu2S

Zinc Zinc blende ZnS Lead Galena PbS

Mercury Cinnabar HgS Calcium Limestone CaCO3

Zinc Calamine ZnCO3

Sodium Rock salt NaCl Magnesium Carnallite KCl.MgCl2.6H2O

Calcium Fluorspar CaF2

Silver Horn silver AgClCalcium Gypsum CaSO4.2H2O

Magnesium Epsom salt MgSO4.7H2O

Barium Barytes BaSO4

Lead Anglesite PbSO4

Halide ores

Sulphate ores

Oxide ores

Carbonate ores

Iron

Sulphide ores

Copper

CORROSION OF METALS

Surface of many metals is easily attacked whenexposed to atmosphere. They react with air or waterpresent in the environment and form undesirablecompounds on their surfaces. These undesirablecompounds are generally oxides.

Thus, corrosion is a process of deterioration of metalas a result of its reaction with air or water (present inenvironment) surrounding it.

(a) Corrosion of Iron :

Iron corrodes readily when exposed to moisture andgets covered with a brown flaky substance called rust.This is also called Rusting of Iron. Chemically, the rustis hydrated iron (III) oxide, Fe

2O

3.xH

2O. Rusting is an

oxidation process in which iron metal is slowly oxidizedby the action of air (in presence of water). Therefore,rusting of iron takes place under the followingconditions:

� Presence of air (or oxygen)

� Presence of water (moisture)

� More the reactivity of the metal, the more will be the

possibility of the metal getting corroded.

(i) Experiment to show that rusting of iron requiresboth air and water -

We take three test tubes and put one clean iron nail ineach of the three test tubes:

(A) In the first test tube containing iron nail, we putsome anhydrous calcium chloride to absorb water (ormoisture) from the damp air present in the test tubeand make it dry.

(B) In the second test tube containing iron nail, we putboiled water. Boiled water does not contain anydissolved air or oxygen in it. A layer of oil is put overboiled water in the test tube to prevent the outside airfrom mixing with boiled water.

(C) In the third test tube containing an iron nail, we putunboiled water so that about two-third of the nail isimmersed in water and the rest is above water exposedto damp air.

After one week, we observe the iron nails kept in all thethree test tubes.

CaCl2Anhydrous Boiled water Water

Rusting of iron

A la

yer

of o

il

Iron nail

Test Tube

Cork

(ii) We will obtain the following observations fromthe experiment :

(A) No rust is seen on the surface of iron nail kept indry air in the first test tube. This tells us that rusting ofiron does not takes place in air alone.

(B) No rust is seen on the surface of iron nail kept in airfree boiled water in the second test tube. This tells usthat rusting of iron does not take place in water alone.

(C) Red brown rust is seen on the surface of iron nailkept in the presence of both air and water in the thirdtest tube. This tells us that rusting of iron takes placein the presence of both air and water together.

(iii) Prevention of rusting :

(A) Corrosion of metals can be prevented by coatingthe metal surface with a thin layer of paint, varnish orgrease.

(B) Iron is protected from rusting by coating it with athin layer of another metal which is more reactive thaniron. This prevents the loss of electrons from ironbecause the active metal loses electrons in preferenceto iron. Zinc is commonly used for covering surface ofiron. The process of covering iron with zinc is calledgalvanization. Iron is also coated with other metalssuch as tin known as tin coating.

(C) By alloying : Some metals when alloyed with othermetals become more resistant to corrosion. Forexample, when iron is alloyed with chromium andnickel, it forms stainless steel. This is resistant tocorrosion and does not rust at all.

(D) To decrease rusting of iron, certain antirustsolutions are used. For example, solutions of alkalinephosphates are used as antirust solutions.

(b) Corrosion of Aluminium :

Due to the formation of a dull layer of aluminium oxidewhen exposed to moist air, the aluminium metal losesits shine very soon after use. This aluminium oxidelayer is very tough and prevents the metal underneathfrom further corrosion (because moist air is not able topass through this aluminium oxide layer). This meanssometimes corrosion is useful.

16PAGE # 16

(c) Corrosion of Copper :

When a copper object remains in damp air for aconsiderable time, then copper reacts slowly withcarbon dioxide and water of air to form a green coatingof basic copper carbonate [CuCO

3.Cu(OH)

2] on the

surface of the object. Since copper metal is low in thereactivity series, the corrosion of copper metal is very,very slow.

(d) Corrosion of Silver :

Silver is a highly unreactive metal, so it does not reactswith oxygen of air easily. But, air usually contains a littleof sulphur compounds such as hydrogen sulphide gas(H

2S), which reacts slowly with silver to form a black

coating of silver sulphide (Ag2S). Silver ornaments

gradually turn black due to the formation of a thin silversulphide layer on their surface and silver is said to betarnished.

ALLOY

An alloy is a homogenous mixture of two or more metalsor a metal and a non-metal.For example, iron is the most widely used metal. But itis never used in the pure form. This is because iron isvery soft and stretches easily when hot. But when it ismixed with a small amount of carbon (about 0.5 to1.5%), it becomes hard and strong. The new form ofiron is called steel.

(a) Objectives of Alloy Making :

Alloys are generally prepared to have certain specific

properties which are not possessed by the constituent

metals. The main objects of alloy-making are:

(i) To increase resistance to corrosion : For example,

stainless steel is prepared which has more resistance

to corrosion than iron.

(ii) To modify chemical reactivity : The chemical

reactivity of sodium is decreased by making an alloy

with mercury which is known as sodium amalgam.

(iii) To increase the hardness : Steel, an alloy of iron

and carbon is harder than iron.

(iv) To increase tensile strength : Magnalium is an

alloy of magnesium and aluminium. It has greater

tensile strength as compared to magnesium and

aluminium.

(v) To produce good casting : Type metal is an alloy of

lead, tin and antimony.

(vi) To lower the melting point : For example, solder is

an alloy of lead and tin (50% Pb and 50% Sn). It has a

low melting point and is used for welding electrical

wires together.

Composition, Properties and uses of some alloys of copper :

Alloy Composition Properties and UsesBrass Cu (60�80%)

Zn (20�40%)

Brass is used for decoration purposes, for making many scientific instruments, telescopes, microscopes, barometers etc.

Bronze Cu (75-90%)Sn (10-25%)

For making statues, cooking utensils and coins.

German silver Cu (30�60%)

Zn (25�35%)

Ni (15�35%)

It is silvery white as silver, malleable and ductile. It is used as imitation silver, in making ornaments and utensils and also for decoration.

Gun metal Cu (88%)Sn (10%)Zn (2%)

It is used for making gears and bearings, and gun barrels.

Bell metal Cu (80%)Sn (20%)

It is used for casting bells.

COMPOSITION, PROPERTIES AND USES OF SOME IMPORTANT ALLOYS OF IRON :Alloy Composition Properties and Uses

Stainless steel Fe (74%)Cr (18%)Ni (8%)

Properties : Stainless steel is hard, tenacious and corrosion resistant.Uses : For making cutlery, utensils, ornamental pieces and other instrument and apparatus.

Nickel steel Fe (96�98%)

Ni (2�4%)

Properties : Nickel steel is hard, elastic and corrosion resistant.Uses : For making electric wire cables, automobile and aeroplane parts, watches, armour plates, propeller shafts, etc.

Alnico Fe (60%)Al (12%)Ni (20-%)Co (8%)

Properties : Highly magnetic.Uses : For making permanent magnets

17PAGE # 17

COMPOSITION, PROPERTIES AND USES OF SOME IMPORTANT ALLOYS OF ALUMINIUM :

Alloy Composition Properties and Uses

Duralium orDuralumin

Al (95%)Cu (4%)Mg (0.5%)Mn (0.5%)

Uses : For making aeroplane, spacecrafts, ships and pressure cookers.Properties : In strength, it is as good as steel but it is very light. It is hard, corrosion-resistant and highly ductile.

Magnalium Al (90�95%)

Mg (5�10%)

Uses : For making light instruments and balance beams.Properties : It is hard and tough.

COMPOSITION, PROPERTIES AND USES OF SOME IMPORTANT ALLOYS OF LEAD

Alloy Composition Properties and Uses

Solder Pb (50%)Sn (50%)

Properties : It has a low-melting point.Uses : Used for soldering purposes.

Type metal Pb (75�80%)

Sb (15�20%)

Sn (5%)

Uses : For making printing type.

AMALGAM

Amalgams are homogenous mixtures of a metal andmercury. For example, sodium amalgam containssodium and mercury.Different amalgams are prepared according to theiruses. For example,(i) Sodium amalgam is produced to decrease thechemical reactivity of sodium metal. It is also used asa good reducing agent.(ii) Tin amalgam is used for silvering cheap mirrors.(iii) The process of amalgamation is used for theextraction of metals like gold or silver from their nativeores.

1818PAGE # 18

ACIDS, BASES AND SALTS

ACIDS

Substances with sour taste are regarded as acids.

Lemon juice, vinegar, grape fruit juice and spoilt milk

etc. taste sour since they are acidic. Many substances

can be identified as acids based on their taste but

some of the acids like sulphuric acid have very strong

action on the skin which means that they are corrosive

in nature. In such cases it would be according to

modern definition -

An acid may be defined as a substance which

releases one or more H+ ions in aqueous solution.

Acids are mostly obtained from natural sources.

(a) Classification of Acids :

(i) Classification of acids on the basis of theirSourceOn the basis of their source, acids can be classified

in two categories :

(A) Organic acids (B) Inorganic acids

(A) Organic acidsThe acids which are usually obtained from

organisms are known as organic acids.

Some Organic Acids with Their Natural Sources

S.No. Organic acid Natural sources S.No. Organic acid Natural sources

1 Acetic acid Vinegar 7 Oleic acid Olive oil

2 Citric acid Citrus fruits (likeorange and lemon)

8 Stearic acid Fats

3 Butyric acid Rancid butter 9 Amino acid Proteins

4 Formic acid Sting of bees and ants 10 Uric acid Urine

5 Lactic acid Sour milk 11 Tartaric acid Tamarind

6 Malic acid Apples 12 Oxalic acid Tomatoes

It may be noted that all organic acids contain carbonas one of their constituting elements. These are weakacids and, therefore, do not ionise completely in theiraqueous solutions. Since these acids do not ionisecompletely in their aqueous solutions, therefore, theirsolutions contains both ions as well asundissociated molecules. For example, formic acid�saqueous solution contains H3O+, HCOO� as well asundissociated HCOOH molecules.

HCOOH + H2O H3O+ + HCOO�

Formic acid Hydronium ion Formate ion

(B) Inorganic Acids. The acids which are usually obtained from mineralsare known as inorganic acids. Since the acids areobtained from minerals, therefore, these acids arealso called mineral acids. Some common examplesof inorganic acids are : Hydrochloric acid (HCl),Sulphuric acid (H2SO4), Nitric acid (HNO3) etc. It may be pointed out that except carbonic acid(H2CO3), these acids do not contain carbon. Acidslike HCl, H2SO4 and HNO3 are strong acids whichionise completely in their aqueous solutions and,therefore, their aqueous solutions do not contain anyundissociated molecules.

(ii) Classification of acids on the basis of theirBasicity :

The basicity of an acid is defined as the number ofhydronium ions [H3O+ (aq.)] that can be produced by

the complete ionisation of one molecule of that acidin aqueous solution.For example, basicity of HCl, H2SO4, H3PO4 is 1, 2

and 3 respectively because one molecule of theseacids, on ionisation, produces 1, 2 and 3 hydroniumions in aqueous solution respectively.

(A) Monobasic Acids :

When one molecule of an acid on complete ionisation

produces one hydronium ion (H3O+) in aqueous

solution, the acid is said to be a monobasic acid.

Examples of Monobasic Acids.

Some examples of monobasic acids are :

(i) Hydrochloric acid (HCl)

(ii) Hydrobromic acid (HBr)

(iii) Nitric acid (HNO3)

(iv) Acetic acid (CH3COOH)

(v) Formic acid (HCOOH)

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1919PAGE # 19

(B) Dibasic Acids :When one molecule of an acid on complete ionisation

produces two hydronium ions (H3O+) in aqueous

solution, the acid is said to be a dibasic acid.

Examples of Dibasic Acids :Some examples of dibasic acids are :

(i) Sulphuric acid (H2SO4)

(ii) Sulphurous acid (H2SO3)

(iii) Carbonic acid (H2CO3)

(iv) Oxalic acid [(COOH)2]

(v) Hydrofluoric acid (HF)

(C) Tribasic Acids :

When one molecule of an acid on complete ionisation

produces three hydronium ions (H3O+) in aqueous

solution, the acid is said to be a tribasic acid.An example of tribasic acids is Phosphoric acid

(H3PO4).

(D) Tetrabasic Acids :

When one molecule of an acid on complete ionisation

produces four hydronium ions (H3O+) in aqueous

solution, the acid is said to be a tetrabasic acid.An example of tetrabasic acids is silicic acid (H4SiO4).

(iii) Classification of acids on the basis of theirstrength :

We know that acids ionise in the aqueous solution to

produce hydronium ions. So, the strength of an acid

depends upon the degree of ionisation, usually

denoted by the letter alpha ().

Degree of ionisation of an acid ()

= Number of molecules of the acid undergoing ionisation

100Total number of acid molecules

More the degree of ionisation () of an acid, more stronger

it will be. Generally, if the degree of ionisation () for

an acid is greater than 30%, it is considered to be a

strong acid. If it is less than 30%,it is considered to be

a weak acid.

On the basis of degree of ionisation, the acids can be

classified as under:

(A) Strong Acids :The acids which undergo almost complete ionisation

in a dilute aqueous solution, thereby producing a high

concentration of hydronium ions (H3O+) are known as

strong acids.

Examples of strong acids :Some examples of strong acids are :

(i) Hydrochloric acid (HCl)

(ii) Sulphuric acid (H2SO4)

(iii) Nitric acid (HNO3)

All these three mineral acids are considered to be

strong acids because they ionise almost completely in

their dilute aqueous solutions.

(B) Weak Acids :

The acids which undergo partial or incomplete

ionisation in a dilute aqueous solution, thereby

producing a low concentration of hydronium ions (H3O+)

are known as weak acids.

Examples of weak acids :

Some examples of weak acids are :

(i) Acetic acid (CH3COOH)

(ii) Formic acid (HCOOH)

(iii) Oxalic acid [(COOH)2]

(iv) Carbonic acid (H2CO3)

(v) Sulphurous acid (H2SO3)

(vi) Hydrogen sulphide (H2S)

(vii) Hydrocyanic acid (HCN)

The aqueous solution of weak acids contain both ions

as well as undissociated molecules.

(iv) Classification on the basis of Concentration of

the Acid :

By the term concentration, we mean the amount of water

present in the given sample of acid solution in water.

(A) Concentrated Acid :

The sample of an acid which contains very small or no

amount of water is called a concentrated acid.

(B) Dilute Acid :

The sample of an acid which contains far more amount

of water than its own mass is known as a dilute acid

It must be mentioned here that concentration of an acid

simply tells the amount of water in the acid. It may not

be confused with strength of an acid, which is a

measure of concentration of hydronium ion it produces

in aqueous solution.

A concentrated acid may not necessarily be a strong

acid while a dilute acid may not necessarily be a weak

acid. A strong acid will remain strong even if it is dilute

because it produces a large concentration of hydronium

ions in aqueous solution. On the other hand, a weak

acid will remain weak even when concentrated because

it will produce lesser concentration of hydronium ions

in aqueous solution.

2020PAGE # 20

CHEMICAL FORMULAE, TYPES AND USES OF SOME COMMON ACIDS

Name Type Chemical Formula Where found or used

Carbonic acid Mineral acid H2CO3 In soft drinks and lends fizz.

Nitric acid Mineral acid HNO3

Used in the manufacture of explosives (TNT, Nitroglycerine) and fertilizers (Ammonium nitrate,

Calcium nitrate, Purification of Au, Ag)

Hydrochloric acid Mineral acid HClIn purification of common salt, in textile industry

as bleaching agent, to make aqua regia, in stomach as gastric juice, used in tanning industry

Sulphuric acid Mineral acid H2SO4

Commonly used in car batteries, in the manufacture of fertilizers (Ammonium

sulphate, super phosphate) detergents etc, in paints, plastics, drugs, in manufacture of artificial

silk, in petroleum refining.

Phosphoric acid Mineral acid H3PO4 Used in antirust paints and in fertilizers.

Formic acid Organic acid HCOOHFound in the stings of ants and bees, used in tanning leather, in medicines for treating gout.

Acetic acid Organic acid CH3COOH Found in vinegar, used as solvent in the manufacture of dyes and perfumes.

Lactic acid Organic acid CH3CH(OH)COOH Responsible for souring of milk in curd.

Benzoic acid Organic acid C6H5COOH Used as a food preservative.

Citric acid Organic acid C6H8O7 Present in lemons, oranges and citrus fruits.

Tartaric acid Organic acid C4H6O6 Present in tamarind.

(b) Chemical Properties of Acids :

(i) Action with metals :Dilute acids like dilute HCl and dilute H

2SO

4 react with

certain active metals to evolve hydrogen gas.

2Na(s) + 2HCl (dilute) 2NaCl(aq) + H2(g)

Mg(s) + H2SO

4 (dilute) MgSO

4(aq) + H

2(g)

Metals which can displace hydrogen from dilute acidsare known as active metals. e.g. Na, K, Zn, Fe, Ca, Mgetc.

Zn(s) + H2SO

4 (dilute) ZnSO

4(aq) + H

2(g)

The active metals which lie above hydrogen in theactivity series are electropositive and more reactive innature. Their atoms lose electrons to form positive ionsand these electrons are accepted by H+ ions of theacid. As a result, H

2 is evolved.

e.g.

Zn(s) Zn2+ (aq) + 2e�

2H+(aq) + SO4

2� (aq) + 2e� H2(g) + SO

42�(aq)

Zn(s) + 2H+(aq) Zn++(aq) + H2(g)

(ii) Action with metal oxides :Acids react with metal oxides to form salt and water.These reactions are mostly carried out upon heating.e.g.

ZnO(s) + 2HCl (aq) ZnCl2(aq) + H

2O()

MgO(s) + H2SO

4(aq) MgSO

4(aq) + H

2O()

CuO(s) + 2HCl(aq.) CuCl2(aq) + H

2O()

(Black) (Bluish green)

(iii) Action with metal carbonates and metalbicarbonates : Both metal carbonates and bicarbonatesreact with acids to evolve CO

2 gas and form salts.

e.g.

CaCO3(s)+ 2HCl(aq) CaCl

2(aq) + H

2O() + CO

2(g)

Calcium Calcium carbonate chloride

2NaHCO3(s) + H2SO4(aq) Na2SO4(aq) + 2H2O(aq) + 2CO2(g)

Sodium Sodiumbicarbonate sulphate

(iv) Action with bases :Acids react with bases to give salt and water.HCl (aq) + NaOH(aq) NaCl + H

2O

Substances with bitter taste and soapy touch areregarded as bases. Since many bases like sodiumhydroxide and potassium hydroxide have corrosiveaction on the skin and can even harm the body, soaccording to the modern definition -A base may be defined as a substance capable ofreleasing one or more OH¯ ions in aqueous solution.

(a) Characteristics of a Base :

(i) A base changes red litmus to blue.

(ii) A base reacts with an acid so that a salt and waterare formed.Base + Acid Salt + Water

(iii) A base combines with carbon dioxide so that acarbonate is formed.

(iv) A base is slippery like soap. It tastes unpleasant

and bitter.

2121PAGE # 21

Distinction Between an Alkali and a Base :

It may be kept in mind that a base which is soluble in

water is called an alkali. On the other hand those bases

which are not soluble in water are termed as bases

only not alkalis. This means that all alkalis are bases

but all bases are not alkalis. For example, ferric

hydroxide [Fe(OH)3] and cupric hydroxide [Cu(OH)2] are

bases but not termed as alkalis because they are

insoluble in water while NaOH and KOH are an

examples of alkalis since these are soluble in water.

(b) Classification of Bases or Alkalis :

Classification of bases or alkalis can be done in

different ways as given below :

(i) Classification on the basis of their strength

(ii) Classification on the basis of their concentration

(iii) Classification on the basis of their acidity.

(i) Classification of the Bases or Alkalis on the Basisof their Strength

We know that alkalis (soluble bases) ionise in

aqueous solution to produce hydroxyl (OH�) ions. So

the strength of an alkali (soluble base) depends upon

its degree of ionisation, usually denoted by the letter

alpha ()

Degree of ionisation of an alkali ()

= moleculesalkaliofnumberTotalionisationundergoingalkalitheofmoleculesofNumber

×100

More the degree of ionisation () of an alkali (or a

soluble base), more stronger it will be. On the basis of

degree of ionisation, the alkalis (or soluble bases)

can be classified as under :

(A) Strong alkalis or bases :

The alkalis or bases which undergo almost complete

ionisation in aqueous solution to produce high

concentration of hydroxyl (OH�) ions are known as

strong alkalis or strong bases.Example of strong alkalis or bases.

Some example of strong alkalis or bases are :

Sodium hydroxide (NaOH), Potassium hydroxide (KOH)

and Barium hydroxide [Ba (OH)2] etc.

)ionisedcompletelyAlmost(

)aq(OH2)aq(Ba)aq()OH(Ba

)aq(OH)aq(K)aq(KOH

)aq(OH)aq(Na)aq(NaOH

22

(B) Weak alkalis or bases :

The alkalis or bases which undergo only partial

ionisation in aqueous solution to produce a relatively

low concentration of hydroxyl (OH�) ions are known as

weak alkalis or weak bases.

Some examples of weak alkalis or bases are : Ammoniumhydroxide (NH4OH), Calcium hydroxide [Ca (OH)2],Magnesium hydroxide [Mg (OH)2] etc.

NH4OH (aq) NH+4

(aq) + OH� (aq)Ca (OH)2 (aq) Ca2+ (aq) + 2OH� (aq)Mg (OH)2 (aq) Mg2+ (aq) + 2OH� (aq)

Since these alkalis are not ionising completely,therefore, there is a dynamic equilibrium between theundissolved alkali and the ions produced by it.

(ii) Classification of Bases or Alkalis on the Basis oftheir Concentration :By the term concentration, we mean the amount ofwater present in the given sample of alkali solution inwater. On the basis of concentration, the alkalis can beclassified as under :

(A) Concentrated alkali :A solution of alkali having a relatively high percentageof alkali in its aqueous solution is known asconcentrated alkali.

(B) Dilute alkali :A solution of alkali having a relatively low percentage ofalkali in its aqueous solution is known as a dilutealkali.If the concentration of alkali in the solution is less than1 mole per litre, then it is considered to be a dilutealkali.

(iii) Classification of Bases or Alkalis on the Basis oftheir Acidity :Before we discuss this classification, let us understandthe meaning of the word �acidity� of an alkali.

The number of hydroxyl (OH�) ions produced by onemolecule of an alkali on complete dissociation in wateror the number of hydrogen ions (of an acid) with whicha molecule of that alkali reacts to produce salt andwater only is known as acidity of an alkali.

For water insoluble hydroxides, acidity of the base isequal to the number of OH� ions present in one moleculeof that base.

On the basis of acidity, the bases can be classified asunder :

(A) Monoacidic Bases (or alkalis) :When one molecule of the base on complete ionisationproduces one hydroxyl (OH�) ion in aqueous solution,the base or alkali is said to be monoacidic

ORA monoacidic base (or alkali) may be defined as onewhose one molecule reacts with one hydrogen (H+)ion completely to form salt and water as the onlyproducts.

Examples of Monoacidic Bases (or alkalis) :

Sodium hydroxide (NaOH), Potassium hydroxide(KOH), Ammonium hydroxide (NH4OH). All thesesubstances produce only one hydroxyl ion on completeionisation in aqueous solution.

NaOH(aq.) Na+ (aq) + OH� (aq)

KOH(aq.) K+ (aq) + OH� (aq)

The dissociation of monoacidic bases or alkalis takes

place in a single step.

2222PAGE # 22

(B) Diacidic Bases (or alkalis) :When one molecule of a base or alkali on completeionisation produces two hydroxyl (OH�) ions inaqueous solution, the base or alkali is said to bediacidic.

ORA diacidic base (or alkali) may be defined as one whoseone molecule reacts with two hydrogen (H+) ionscompletely to form salt and water as the only products.

Examples of Diacidic Bases

Calcium hydroxide [Ca (OH)2] and magnesiumhydroxide [Mg (OH)2]

Ca(OH)2(aq) Ca2+ (aq) + 2OH� (aq)

Mg(OH)2(aq) Mg2+ (aq) + 2OH� (aq)One molecule of both the bases are producing 2OH�

ions in aqueous solution, therefore, these are termedas diacidic bases .

(C) Triacidic Bases :

When one molecule of a base or alkali on complete

ionisation produces three hydroxyl (OH�) ions in aqueous

solution, the base or alkali is said to be triacidic base.

Examples of Triacidic Bases :

Aluminium hydroxide [Al(OH)3], Ferric hydroxide

[Fe (OH)3]

Al (OH)3(aq) Al3+ (aq) + 3OH� (aq)

Al (OH)3 + 3HCl (aq) AlCl3 + 3H2O

In the above equations, one molecule of Al (OH)3 is

producing three OH� ions and one molecule of Al (OH)3is reacting with three hydrogen (H+) ions to form salt

and water only, therefore, it is termed as a triacidic

base.

CHEMICAL FORMULAE, NAMES AND USES OF SOME COMMON BASES

Name Commercial

Name Chemical Formula

Uses

Sodium hydroxide

Causticsoda

NaOHIn manufacture of soap,

paper, pulp, rayon, refining of petroleum etc.

Potassium hydroxide

Causticpotash

KOH

In alkaline storage batteries, manufacture of soap, absorbing

CO2 gas etc.

Calcium hydroxide

Slakedlime

Ca(OH)2In manufacture of bleaching powder,

softening of hard water etc.Magnesium hydroxide

Milk ofmagnesia

Mg(OH)2As an antacid to remove acidity

from stomach.Aluminium hydroxide

� Al(OH)3 As foaming agent in fire extinguishers.

Ammonium hydroxide

� NH4OHIn removing grease stains from

clothes and in cleaning window panes.

(c) Chemical Properties :

(i) Action with metals :Metals like zinc, tin and aluminium react with strongalkalies like NaOH (caustic soda), KOH (causticpotash) to evolve hydrogen gas.

Zn(s) + 2NaOH(aq) Na2ZnO

2(aq) + H

2(g)

Sodium zincate

Sn(s) + 2NaOH(aq) Na2SnO

2(aq) + H

2(g)

Sodium stannite

2Al(s)+ 2NaOH + 2H2O 2NaAlO

2(aq) + 3H

2(g)

Sodium metaaluminate

(ii) Action with non-metallic oxides :Acids react with metal oxides, but bases react withoxides of non-metals to form salt and water.e.g.2NaOH(aq) + CO

2(g) Na

2CO

3(aq) + H

2O()

Ca(OH)2(s) + SO

2(g) CaSO

3(aq) + H

2O()

Ca(OH)2(s) + CO

2(g) CaCO

3(s) + H

2O()

Acids BasesSour in taste. Bitter in taste.Change Colours of indicatorse.g. litmus turns from blue tored, phenolphthalein remains colourless.

Change colours of indicators e.g.litmus turns from red to blue phenolphthalein turns fromcolourless to pink.

Show electrolytic conductivityin aqueous solution.

Show electrolytic conductivity inaqueous solution

Acidic properties disappearwhen react with bases (Neutralization)

Basic properties disappear whenreact with acids (Neutralization)

Acids decompose carbonatesalts.

No decomposition of carbonatesalts by bases

.

DILUTION OF ACIDS AND BASES

Acids and bases are mostly water soluble and can bediluted by adding the required amount of water. Withthe addition of water the amount of acid or base perunit volume decrease and dilution occurs. The processis generally exothermic in nature. When a concentratedacid like sulphuric acid or nitric acid is to be dilutedwith water, acid should be added dropwise to watertaken in the container with constant stirring.

2323PAGE # 23

INDICATORS

An indicator is a chemical compound which indicatesthe presence of acidic, basic or neutral substanceeither by change in colour or odour. For Example, Litmus is an indicator which is red in an acidic solutionbut it has blue colour in a basic solution.

The indicators change colour when the nature of thesolution is changed (from acidic to basic or vice versa).Therefore they are also known as visual indicators.

(a) Types of Indicators :

(i) Olfactory Indicators :Some substances whose odour changes in acidic orbasic solutions are called olfactory indicators. Thecommonly used olfactory indicators are raw onion,vanilla extract and clove oil. Just as visual indicatorschange their colour in response to acidic or basicsolution, an olfactory indicator will change either itsodour or odour intensity with change in acidic or basicnature of solution.

(ii) On the other hand their are some substances whichindicates the behaviour of solution by showing changein colour.

(A) Litmus : Litmus is a purple dye which is extractedfrom a plant �lichen�. A blue litmus strip, when dipped

in an acid solution acquires red colour. Similarly a redstrip when dipped in a base solution becomes blue.

(B) Phenolphthalein :It is also an organic dye. In neutralor acidic solution, it remains colourless while in thebasic solution, the colour of the indicator changes topink.

(C) Methyl Orange :Methyl orange is an orangecoloured dye and keeps its colour in the neutral orbasic medium. In the acidic medium the colour of theindicator becomes red.

(D) Red Cabbage Juice :It is purple in colour in neutralmedium and turns red or pink in the acidic medium. Inthe basic or alkaline medium, its colour changes togreen.

(E) Turmeric juice : It is yellow in colour and remainsas such in the neutral and acidic medium. In the basicmedium its colour becomes reddish or deep brown.

(F) China Rose : Extract of china rose (Gudhal) petalsis of pink colour. It will change into dark pink (magenta)in acidic solution and green in basic solution.

Table of indicators :

Indicator Colour in acidic

mediumColour in basic

medium

Blue litmus Red Blue

Red litmus Red Blue

Turmeric Yellow Reddish-brown

China rose Dark pink (magenta) Green

Methyl orange Red Orange

Phenolphthalein Colourless Pink

NEUTRALISATION

It may be defined as a reaction between acid and

base present in aqueous solution to form salt and

water.

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

Basically neutralisation is the combination between

H+ ions of the acid with OH� ions of the base to form H2O.

e.g.

H+(aq)+ Cl�(aq) + Na+(aq) + OH�(aq) Na+(aq) + Cl�(aq) + H2O()

H+(aq) + OH�(aq) H2O()

Neutralisation reaction involving an acid and base is

of exothermic nature. Heat is evolved in all

neutralisation reactions. If both acid and base are

strong, the value of heat energy evolved remains same

irrespective of their nature.

e.g .

HCl (aq) + NaOH (aq) NaCl (aq) + H2O () + 57.1 KJ

Strong Strong

acid base

HNO3 (aq) + KOH (aq) KNO3 (aq) + H2O () + 57.1 KJ

Strong Strong

acid base

Strong acids and strong bases are completely ionised

of their own in the solution. No energy is needed for

their ionisation. Since the cation of base and anion of

acid on both sides of the equation cancel out

completely, the heat evolved is given by the following

reaction -

H+ (aq) + OH� (aq) H2O () + 57.1 KJ

Applications of Neutralisation :

(i) People particularly of old age suffer from acidity in

the stomach which is caused mainly due to release of

excessive gastric juices containing HCl. The acid is

neutralised by antacid tablets which contain mild bases

like sodium hydrogen carbonate (baking soda),

magnesium hydroxide etc.

(ii) The stings of bees and ants contain formic acid. Its

corrosive and poisonous effect can be neutralised by

rubbing soap which contains NaOH (an alkali).

(iii) The stings of wasps contain an alkali and its

poisonous effect can be neutralised by vinegar (dil.

solution of acetic acid).

(iv) Farmers generally neutralize the effect of acidity in

the soil caused by acid rain by adding slaked lime

(Calcium hydroxide) to the soil.

2424PAGE # 24

pH SCALE

If an aqueous solution has equal concentrations ofhydrogen ions and hydroxide ions in it, it is neitheracidic nor basic, it is said to be neutral. Now, if anaqueous solution has more number of hydrogen ions(and less number of hydroxide ions), it will be an acidicsolution. On the other hand, if an aqueous solutionhas more number of hydroxide ions (and less numberof hydrogen ions), it will be basic (or alkaline) in nature.So, we usually describe the acidic nature or basicnature (alkaline nature) of aqueous solution in termsof hydrogen ion concentration or hydroxide ionconcentration in it.In 1909, Sorensen devised a scale known as pH scaleon which the acidic nature as well as the basic natureof solutions can be expressed only by considering thehydrogen ion concentrations in them. Now, thehydrogen ion concentrations of most of the commonaqueous solutions are in negative powers of 10. Byusing the Sorensen�s pH scale, the hydrogen ion

concentrations of solution having complicated negativepowers of 10 are converted into simple positive figuresof pH values.For example, in neutral water at 25°C (298K) the

concentration of hydrogen ions is 10�7 M and that ofthe hydroxide ions is also 10�7 M. It is not convenient touse such negative powers, therefore a scale has beenfound on which these negative powers of concentrationcan be written as positive figures. This is known as pHscale.

The pH of a solution is the logarithm (to the base 10)of the reciprocal of its hydrogen ion concentration inmoles per litre.

pH = 1log

[H ]

or pH = � log [H+]

For calculating the pH of a solution, we have to use itshydrogen ion concentration [H+] in moles per litre. Inother words, we have to use the molar concentration(M) of the hydrogen ions. By using the above formula,we can calculate the pH of a solution from its hydrogenion concentration. And if we know the pH value, we canfind out the hydrogen ion concentration. We will nowcalculate the pH values of pure water, acidic solutionsand alkaline solutions (or basic solutions) by usingthis formula.

pH of Pure Water :The concentration of hydrogen ions in pure water is10�7 M which means [H+] = 10�7

pH = � log [H+]= � log [10�7]= � [�7] ( log 10�7 = � 7)

or pH = 7

Thus, pH of water is 7. Whenever the pH of a solution

is 7, it will be a neutral solution. Such a solution will

have no effect on any litmus solution or any other

indicator.

pH of an Acidic Solution

Let us add some acid to water so that the hydrogen ion

concentration in water increases to 10�2 M (this is more

than 10�7 M).

Thus, pH of this solution = � log [H+]

= � log [10�2]

= � [�2] (log 10�2 = � 2)

pH of this acidic solution = 2

All the acidic solutions have a pH less than 7.

So, whenever a solution has a pH of less than 7, it will

be acidic in nature and it will turn blue litmus red as

well as methyl orange indicator will be turned red. Acidic

solution do not change the colour of phenolphthalein .

pH of an Alkaline Solution

Let us add some alkali to water so that the

concentration of hydroxide ions in water increases to

10�3 M. Now, our formula for calculating the pH uses

hydrogen ion concentration, therefore, to calculate the

pH of this solution containing 10�3 M hydroxide ions,

we have to calculate the concentration of hydrogen ions

in this alkaline solution. This can be done by using the

formula for the ionic product of water.

Kw = [H+] × [OH�]

or 10�14 = [H+] × [OH�]

or KW = 10�14 M 2 1M = mol / L

Here, the concentration of hydroxide ions,

[OH�] = 10�3 M

Thus, 10�14 = [H+] × 10�3

or [H+]= 14

3

10

10

= 10�11 M

So, the concentration of hydrogen ions in this alkaline

solution is 10�11 M.

pH of this alkaline solution

= � log [H+]

= � log [10�11]

= � [� 11] = 11

All the alkaline solution have a pH of more than 7.

So, whenever a solution has a pH of more than 7, it will

be alkaline in nature and it will turn red litmus blue and

colourless phenolphthalein solution pink.

Note : pH is a pure number. It has no unit.

2525PAGE # 25

THE pH SCALE HAVING pH VALUES FROM 0 TO 14

It should be noted that an acid solution having low pH value is stronger than another solution having higher pH value.

For example, a solution having pH of 2 is a stronger acid than another solution having a pH of 5. It is just opposite in

the case of alkaline solutions. An alkaline solution having higher pH value is a stronger alkali than another solution

having low pH value. For example, a solution having a pH of 12 is a stronger alkali than a solution of pH 10.

The pH values of some of the common substances from our everyday life are given in the following table.

Solution pH Solution pH1. 1M Hydrochloric acid 0 14. Urine 5.5 - 7.52. Battery acid 0.5 15. Saliva 6.5 - 7.53. Dilute hydrochloric acid 1.0 16. Blood 7.3 - 7.54. Gastric juices (Digestive juices in stomach)

1.0 - 3.0 17. Eggs 7.8

5. Lemon juice 2.2 - 2.4 18. Baking soda solution 8.46. Vinegar 2.4 - 3.4 19. Sea-water 8.57. Soft drinks 3.0 20. Washing soda solution 9.08. Wine 2.8 - 3.8 21. Lime-water 10.59. Oranges 3.6 22. Milk of magnesia

[Mg (OH)2 solution]

10.5

10. Tomato juice 4.0 - 4.4 23. 1 M Ammonium hydroxide (Household ammonia)

11.6

11. Beer 4.0 - 5.0 24. Dilute sodium hydroxide 1312. Coffee 4.5 - 5.5 25. 1 M Sodium hydroxide 1413. Milk 6.5

Significance of pH in daily life :

(i) pH in our digestive system : Dilute hydrochloric

acid produced in our stomach helps in the digestion of

food. However, excess of acid causes indigestion and

leads to pain as well as irritation. The pH of the

digestive system in the stomach will decrease. The

excessive acid can be neutralised with the help of

antacids which are recommended by the doctors.

Actually, these are group of compounds (basic in

nature) and have hardly any side effects. A very popular

antacid is �Milk of Magnesia� which is insoluble

magnesium hydroxide. Aluminium hydroxide and

sodium hydrogen carbonate can also be used for the

same purpose. These antacids will bring the pH of the

system back to its normal value. The pH of human

blood varies between 7.36 to 7.42 . It is maintained by

the soluble bicarbonates and carbonic acid present in

the blood.

(ii) pH change leads to tooth decay : The white enamelcoating on our teeth is of insoluble calcium phosphatewhich is quite hard. It is not affected by water. However,when the pH in the mouth falls below 5.5 the enamelgets corroded. Water will have a direct access to theroots and decay of teeth will occur. The bacteria presentin the mouth break down the sugar that we eat in oneform or the other to acids, Lactic acid is one of these.The formation of these acids causes decrease in pH.It is therefore advisable to avoid eating sugary foodsand also to keep the mouth clean so that sugar andfood particles may not be present. The tooth pastescontain in them some basic ingredients and they helpin neutralising the effect of the acids and alsoincreasing the pH in the mouth.

(iii) Role of pH in curing stings by insects : The stingsof bees and ants contain methanoic acid (or formicacid). When stung, they cause lots of pain and irritation.The cure is in rubbing the affected area with soap.Sodium hydroxide present in the soap neutralises acidinjected in the body and thus brings the pH back to itsoriginal level bringing relief to the person who has beenstung. Similarly, the effect of stings by wasps containingalkali is neutralised by the application of vinegar whichis dilute solution of ethanoic acid (or acetic acid)

2626PAGE # 26

(iv) Soil pH and plant growth : The growth of plants in a

particular soil is also related to its pH. Actually, different

plants prefer different pH range for their growth. It is

therefore, quite important to provide the soil with proper

pH for their healthy growth. Soils with high iron minerals

or with vegetation tend to become acidic. The soil pH

can reach as low as 4.The acidic effect can be

neutralised by �liming the soil� which is carried by adding

calcium hydroxide. These are basic in nature and have

neutralising effect. Similarly, the soil with excess of lime

stone or chalk is usually alkaline. Sometimes, its pH

reaches as high as 8.3 and is quite harmful for the plant

growth. In order to reduce the alkaline effect, it is better to

add some decaying organic matter (compost or manure).

The soil pH is also affected by the acid rain and the use

of fertilizers. Therefore soil treatment is quite essential.

SALTS

A substance formed by neutralization of an acid with a

base is called a salt.

e.g.

Ca(OH)2(aq)+ H

2SO

4(aq) CaSO

4(aq) + 2H

2O()

Cu(OH)2(aq) + 2HNO

3(aq) Cu(NO

3)

2(aq) + 2H

2O()

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

(a) Classification of Salts :

It is based on their Mode of Formation :

(A) Normal Salts :The salts which are obtained by complete replacement

of the ionisable hydrogen atoms of an acid by a metallic

or an ammonium ion are known as normal salts. For

example, normal salts NaCl and Na2SO4 are formed by

the complete replacement of ionisable hydrogen atoms

of HCl and H2SO4 respectively

HCl + NaOH NaCl + H2O

Sodium chloride

(Normal salt)

H2SO4 + 2NaOH Na2SO4 + 2H2O

Sodium sulphate

(Normal salt)

(B) Acid Salts : The salts which are obtained by the

partial replacement of ionisable hydrogen atoms of a

polybasic acid by a metal or an ammonium ion are known

as acid salts.

These are usually formed when insufficient amount of

the base is taken for the neutralisation of the acid. For

example, when insufficient amount of NaOH is taken to

neutralise H2SO4, we get an acid salt NaHSO4.

H2SO4 + NaOH NaHSO4 + H2O (Insufficient amount) Sodium hydrogensulphate

(Acid salt)

In this case, only one hydrogen atom out of two hasbeen replaced by sodium atom. Since there is onemore hydrogen atom in NaHSO4 which can bereplaced, therefore, it further reacts with anothermolecule of NaOH to produce Na2SO4 which is anormal salt.

NaHSO4 + NaOH Na2SO4 + H2OSodium Sodium

hydrogensulphate sulphate(Acid salt) (Normal salt)

Acid salts ionise in aqueous solution to producehydronium ions (H3O+), therefore, they exhibit all theproperties of acids.

Some other examples of acid salts are given in Table.

(C) Basic Salts : The salts which are formed bypartial replacement of hydroxyl (�OH) groups of a di-

or a triacidic base by an acid radical are known asbasic salts.

These are usually formed when an insufficientamount of acid is taken for the neutralisation of thebase. For example, when insufficient amount of HClis added to Lead hydroxide, Basic lead chloride[Pb(OH)Cl] is formed

Pb(OH)2 + HCl Pb (OH)Cl + H2OLead hydroxide Basic Lead chloride(Diacidic base) (Basic salt)

Basic salts, for example, Pb(OH)Cl further reacts withHCl to form normal salts

Pb (OH) Cl + HCl PbCl2 + H2OBasic Lead Chloride Lead chloride(Basic salts) (Normal salt)

Some other important examples of basic salts are :

(i) Basic copper chloride, Cu(OH)Cl.(ii) Basic copper nitrate, Cu(OH)NO3(iii) Basic lead nitrate, Pb(OH)NO3.

(D) Double Salts : The salts which are obtained bythe crystallisation of two simple salts from a mixtureof their saturated solutions are known as doublesalts.For example, a double salt potash alum [K2SO4.Al2(SO4)3. 24H2O] is prepared by mixing saturatedsolutions of two simple salts, K2SO4 and Al2(SO4)3and crystallization of the mixture.

K2SO4 + Al2(SO4)3+ 24H2O Crystallisation

K2SO4. Al2(SO4)3. 24

H2O

Potassium Potash alumsulphate (Double salt)

Some other examples of double salts are :

(i) Mohr�s Salt, FeSO4 .(NH4)2SO4. 6H2O,

(ii) Dolomite, CaCO3. MgCO3,

(iii) Carnallite, KCl. MgCl2.6H2O

2727PAGE # 27

(E) Mixed Salts : The salts which contain more than

one type of acidic or basic radicals are called mixed

salts. For example, Sodium potassium carbonate

(NaKCO3) is a mixed salt containing two basic radicals

sodium and potassium. Similarly, calcium oxy chloride,

Ca(OCI)Cl is also a mixed salt containing two acid

radicals OCI� and Cl�.

Some other important examples of mixed salts are :

Sodium potassium sulphate (NaKSO4) (containing two

basic radicals), Disodium potassium phosphate

(Na2KPO4) (containing two basic radicals).

(ii) Classification of salt solutions based on pH

values :

Salts are formed by the reaction between acids

and bases. Depending upon the nature of the acids

and bases or upon the pH values, the salt solutions

are of three types.

(A) Neutral salt solutions : Salt solutions of strong

acids and strong bases are neutral and have pH equal

to 7. They do not change the colour of litmus solution.

e.g. NaCl, KCl, NaNO3, Na

2SO

4 etc.

(B) Acidic salt solutions : Salt solutions of strong ac-

ids and weak bases are of acidic nature and have pH

less than 7. They change the colour of blue litmus

solution to red.

e.g. (NH4)

2SO

4, NH

4Cl etc.

In both these salts, the base NH4OH is weak while the

acids H2SO

4 and HCl are strong.

(C) Basic salt solutions : Salt solutions of strong bases

and weak acids are of basic nature and have pH more

than 7. They change the colour of red litmus solution to

blue.

e.g. Na2CO

3, K

3PO

4 etc.

In both the salts, bases NaOH and KOH are strong

while the acids H2CO

3 and H

3PO

4 are weak.

General formula of alum is [M'SO4. M"(SO

4)

3.24H

2O],

where M' is K+, Na+ or NH4

+ and M" is Al+3, Fe+3 or Cr+3.

�Alum is a double salt of aluminium sulphate and

potassium or sodium sulphate. It is obtained by mixing

solutions of aluminium sulphate and potassium,

ammonium or sodium sulphate and co-crystallising

it. Alum contains large excess of water of crystallisation,

hence on warming, it expands in volume. Alum is used

mainly for softening of water.

K2SO4 + Al2(SO4)3 + 24H2O K2SO4.Al2(SO4)3.24H2O

Potash alum

HYDRATED SALTS - SALTS CONTAINING

WATER OF CRYSTALLISATION:

Certain salts contain definite amount of some watermolecules loosely attached to their own molecules.These are known as hydrated salts and are of crystal-line nature. The molecules of water present are knownas �water of crystallisation�.

In coloured crystalline and hydrated salts, themolecules of water of crystallisation also account fortheir characteristic colours. Thus, upon heating of hy-drated salt, its colour changes since molecules ofwater of crystallisation are removed and the salt be-comes anhydrous.For example, take a few crystals of blue vitriol i.e. hy-drated copper sulphate in a dry test tube or boilingtube. Heat the tube from below. The salt will change toa white anhydrous powder and water droplet will ap-pear on the walls of the tube. Cool the tube and add afew drops of water again. The white anhydrous pow-der will again acquire blue colour.

CuSO4. 5H

2O

CuSO4 + 5H

2O

Copper sulphate Copper sulphate(Hydrated) (Anhydrous)