Chemistry is basically an experimental science. In it we study physical and chemical properties of substance and measure it upto possibility. The results of measurement can we reported in two steps : (a) Arithmetic number, (b) Unit of measurement.

 

Every experimental measurement vary slightly from one another and involves some error or uncertainty depending upon the skill of person making the measurements and measuring instrument. The closeness of the set of values obtained from identical measurement called precision and a related term, refers to the closeness of a single measurement to its true value called accuracy.

 

In the measured value of a physical quantity, the digits about the correctness of which we are surplus the last digit which is doubtful, are called the significant figures. Number of significant figures in a physical quantity depends upon the least count of the instrument used for its measurement.

1. Common rules for counting significant figures : Following are some of the common rules for counting significant figures in a given expression:

Rule 1. All non zero digits are significant.

Example :

x = 1234

has four significant figures. Again

x = 189

has only three significant figures.

Rule 2. All zeros occurring between two non zero digits are significant.

Example :

x = 1007

has four significant figures. Again

x = 1.0809

has five significant figures.

Rule 3. In a number less than one, all zeros to the right of decimal point and to the left of a non zero digit are not significant.

Example :

x = 0.0084

has only two significant digits. Again,

x = 1.0084 has five significant figures. This is on

account of rule 2.

Rule 4. All zeros on the right of the last non zero digit in the decimal part are significant.

Example :

x = 0.00800

has three significant figures 8, 0, 0. The zeros before 8 are not significant again 1.00

has three significant figures.

Rule 5. All zeros on the right of the non zero digit are not significant.

Example :

x = 1000

has only one significant figure. Again

x = 378000 has three significant figures.

Rule 6. All zeros on the right of the last non zero digit become significant, when they come from a measurement. Example : Suppose distance between two stations is measured to be 3050 m. It has four significant figures.

The same distance can be expressed as 3.050 km or 3.050 ´10⁵ cm . In all these expressions, number of significant figures continues to be four. Thus we conclude that change in the units of measurement of a quantity does not change the number of significant figures. By changing the position of the decimal point, the number of significant

digits in the results does not change. Larger the number of significant figures obtained in a measurement, greater is

the accuracy of the measurement. The reverse is also true.

2. Rounding off : While rounding off measurements, we use the following rules by convention:

Rule 1. If the digit to be dropped is less than 5, then the preceding digit is left unchanged.

Example :

x = 7.82

is rounded off to 7.8, again

x = 3.94

is rounded off to 3.9.

Rule 2. If the digit to be dropped is more than 5, then the preceding digit is raised by one. Example : x = 6.87 is rounded off to 6.9, again x = 12.78 is rounded off to 12.8.

Rule 3. If the digit to be dropped is 5 followed by digits other than zero, then the preceding digit is raised by one. Example : x = 16.351 is rounded off to 16.4, again x = 6.758 is rounded off to 6.8.

 

 

Rule 4. If digit to be dropped is 5 or 5 followed by zeros, then preceding digit is left unchanged, if it is even. Example : x = 3.250 becomes 3.2 on rounding off, again x = 12.650 becomes 12.6 on rounding off.

Rule 5. If digit to be dropped is 5 or 5 followed by zeros, then the preceding digit is raised by one, if it is odd. Example : x = 3.750 is rounded off to 3.8. again x = 16.150 is rounded off to 16.2.

(3)Significant figure in calculation

1. Addition and subtraction : In addition and subtraction the following points should be remembered :

1. Every quantity should be changed into same unit.
2. If a quantity is expressed in the power of 10, then all the quantities should be changed into power of 10.
3. The result obtained after addition or subtraction, the number of figure should be equal to that of least, after decimal point.

(ii)Multiplication and division

1. The number of significant figures will be same if any number is multiplied by a constant.
2. The product or division of two significant figures, will contain the significant figures equal to that of least.

 

 

The chosen standard of measurement of a quantity which has essentially the same nature as that of the quantity is called the unit of the quantity. Following are the important types of system for unit,

1. C.G.S. System   :    Length (centimetre), Mass (gram), Time (second)
2. M.K.S. System   :    Length (metre), Mass (kilogram), Time (second)
3. F.P.S. System    :    Length (foot), Mass (pound), Time (second)
4. S.I. System : The 11th general conference of weights and measures (October 1960) adopted International system of units, popularly known as the SI units. The SI has seven basic units from which all other units are derived called derived units. The standard prefixes which helps to reduce the basic units are now widely used.

Dimensional analysis : The seven basic quantities lead to a number of derived quantities such as pressure, volume, force, density, speed etc. The units for such quantities can be obtained by defining the derived quantity in terms of the base quantities using the base units. For example, speed (velocity) is expressed in distance/time.

So the unit is

ms -2 .

m / s

or ms -1 . The unit of force (mass ´ acceleration) is kg ms -2

 

Seven basic SI units

and the unit for acceleration is

 

Length	Mass	Time	Temperature	Electric Current	Luminous Intensity	Amount of substance
metre (m)	Kilogram (kg)	Second (s)	Kelvin (K)	Ampere (A)	Candela (Cd)	Mole (mol)

 

 

Derived Units

 

Physical quantity	Unit	Symbol
Area	square metre	m2
Volume	cubic metre	m3
Velocity	metre per second	ms–1
Acceleration	metre per second square	ms–2
Density	kilogram per cubic metre	kg m–3
Molar mass	kilogram per mole	kg mol–1
Molar volume	cubic metre per mole	m3 mol–1
Molar concentration	mole per cubic metre	mol m–3
Force	newton (N)	kg m s–2
Pressure	pascal (Pa)	N m–2
Energy work	joule (J)	kg m2 s–2, Nm

Standard prefixes use to reduce the basic units

 

Multiple	Prefix	Symbol	Submultiple	Prefix	Symbol
1024	yotta	Y	10–1	deci	d
1021	zetta	Z	10–2	centi	c
1018	exa	E	10–3	milli	m
1015	peta	P	10–6	micro	m
1012	tera	T	10–9	nano	n
109	giga	G	10–12	pico	p
106	mega	M	10–15	femto	f
103	kilo	k	10–18	atto	a
102	hecto	h	10–21	zeto	z
101	deca	da	10–24	yocto	y

 

 

 

Conversion factors

 

1 m = 39.37 inch	1 cal = 4.184 J	1 e.s.u. = 3.3356 ´ 10–10 C	1 mole of a gas = 22.4 L at STP
1 inch = 2.54 cm	1 eV = 1.602 ´ 10–19 J	1 dyne = 10–5 N	1 mole a substance = N0 molecules
1 litre = 1000 mL	1 eV/atom =96.5 kJ mol–1	1 atm = 101325 Pa	1 g atom = N0 atoms
1 gallon (US) = 3.79 L	1 amu = 931.5016 MeV	1 bar = 1 ´ 105 N m–2	t (oF) = 9/5 t (oC) + 32
1 lb = 453.59237 g	1 kilo watt hour = 3600 kJ	1 litre atm = 101.3 J	1 g cm–3 = 1000 kg m–3
1 newton =1 kg m s–2	1 horse power = 746 watt	1 year = 3.1536 ´ 107 s	1Å = 10–10 m
1 J = 1 Nm =1 kg m2 s–2	1 joule = 107 erg	1 debye (D) = 1 ´ 10 –18 esu cm	1nm = 10–9 m

 

 

Matter is the physical material of the universe which occupies space and has mass e.g., water, sugar, metals, plants etc. Matter can be classified as,

 

MATTER

 

Everything that has mass and occupies space

 

Physical classification                                                                         Chemical classification

                                                

Solids          Liquids          Gases                                   MIXTURES                                                 PURE SUBSTANCES

• Variable composition
• Components retain their characteristic properties
• May be separated into pure components by physical methods
• Mixtures of different composition may have widely different properties

• Fixed composition
• Cannot be separated into simpler substances by physical methods
• Properties do not vary

 

                           

Homogeneous              Hetrogeneous                  Elements                     Compounds

• Have same composition throughout
• Components are indistin-guishable for example : a gaseous mixture or a liquid solution

• Do not have same composition throughout
• Components are distin- guishable for example carbon and sulphur mixture (gun powder)

• Can not be decom- posed into simpler substances by chemical changes

 

• Can be decomposed into simpler substances by chemical changes, always at constant composition

 

 

Inorganic          Organic

 

Metals            Metalloids            Non-metals

 

 

 

Each component of a mixture retains its own properties and thus a mixture can be separated into its components by taking advantages of the difference in their physical and chemical properties. The different methods which are employed to separate the constituents from a mixture to purify an impure sample of a substance are,

1. Sedimentation : It is also called gravity separation. It is used for a mixture in which one component is a liquid and the other is insoluble solid heavier than the liquid. Example : Sand dispersed in water.
2. Filtration : It is used for a mixture containing two components one of which is soluble in a particular solvent and the other is not. Example : (i) A mixture of salt and paper using water as solvent (ii) A mixture of sand and sulphur using carbon disulphide as solvent. (iii) A mixture of glass powder and sugar, using water as a solvent in which sugar dissolves but glass does not. (iv) A mixture of sand and sulphur, using carbon disulphide as the solvent in which sulphur dissolves but sand does not.
3. Sublimation : It is used for a mixture containing a solid component, which sublimes on heating from

non-volatile solids. Example : A mixture of sand + naphthalene or powdered glass +

NH 4 Cl

/ camphor / iodine

etc. can be separated by the method of sublimation because substances like naphthalene, etc. form sublimates whereas sand, glass etc. do not.

NH 4 Cl , iodine, camphor

 

 

4. Evaporation : It is used for a mixture in which one component is a non–volatile soluble salt and other is a liquid. Example : Sodium chloride dissolved in sea–water.
5. Crystallization : It is a most common method for a mixture containing solid components and based upon the differences in the solubilities of the components of the mixture into a solvent. For separation, a suitable solvent is first selected. It is of two types :
   1. Simple crystallization : It is applicable when there is a large difference in the solubilities of different components of a mixture.
   2. Fractional crystallization : It is applicable when there is a small difference in the solubilities of different

components of a mixture in the same solvent. Example :

K 2 Cr2 O7

and KCl . Here

K 2 Cr2 O7

is less soluble in

water and hence crystallizes first. A series of repeated crystallization separates the two components in pure form.

6. Distillation : It is the most important method for purifying the liquids. It involves the conversion of a liquid to its vapours on heating (vaporisation) and then cooling the vapours back into the liquid (condensation). It can be used to separate, (i) A solution of a solid in a liquid. e.g., aqueous copper sulphate solution. (ii) A solution of two liquids whose boiling points are different. Several methods of distillation are employed.
   1. Simple distillation : It is used only for such liquids which boil without decomposition at atmospheric pressure and contain non–volatile impurities. Example : (a) Pure water from saline water. (b) Benzene from toluene.
   2. Fractional distillation : It is used for the separation and purification of a mixture of two or more miscible liquids having different boiling points. The liquid having low boiling point vaporises first, gets condensed and is collected in the receiver. The temperature is then raised to the boiling point of second so that the second liquid vaporises and is collected in other receiver. If two liquids present in a mixture have their boiling points closer to each other, a fractionating column is used. Example : (a) Crude petroleum is separated into many useful products such as lubricating oil, diesel oil, kerosene and petrol by fractional distillation. (b) A mixture of acetone and methyl alcohol.
   3. Vacuum distillation or distillation under reduced pressure : It is used for such liquids which decompose on heating to their boiling points. At reduced pressure, the boiling point of liquid is also reduced. Example : Glycerol is distilled under pressure as it decomposes on heating to its boiling point.
   4. Steam distillation : It is used for liquids which are partially miscible with water, volatile in steam. e.g., aniline, oils etc. are purified by steam distillation. The principle involved is of reduced pressure distillation. If P_w and

P_l are vapour pressures of water and liquid at distillation temperature, then

Pw + Pl = P = 76 cm

i.e., atmospheric

pressure. Thus, a liquid boils at relatively low temperature than its boiling point in presence of steam.

Example :    Some solids   like   naphthalene,   o-nitrophenol   which   are   steam   volatile   can   be   purified.

Nitrobenzene, chlorobenzene, essential oils are also extracted or separated by this process

7. Solvent extraction : It is used for the separation of a compound from its solution by shaking with a suitable solvent. The extraction follows Nernst distribution law. The solvent used must be insoluble with other phase in which compound is present as well as the compound should be more soluble in solvent. The extraction becomes more efficient if the given extracting liquid is used for more number of extractions with smaller amounts than using it once with all extracting liquid. Example : (i) Aqueous solution of benzoic acid by using benzene. (ii) Aqueous solution of Iodine by using chloroform or carbon tetrachloride. (iii) Flavour of tea from the tea leaves by boiling with water.

 

 

8. Magnetic separation : It is used for a mixture in which one component is magnetic while the other is non–magnetic. Example : iron ore from the non–magnetic impurities.
9. Chromatography : It is based on the differences in the rates at which different components of a mixture are absorbed on a suitable solvent. It is used in separation, isolation, purification and identification of a substance. It was proposed by a Russian botanist Tswett.
10.  

2

Atmolysis : It is used for separation of the mixture of gases or vapours. It is based on the difference in

 

their rates of diffusion through a porous substance. Example : their hexa–fluorides.

H 2 , SO2 , CH4

and O , U 235 & U 238 in the form of

11. Electrophoresis : It is based upon the differences in the electrical mobility of the different components of the mixture.
12.  

Ultracentrifugation : It is based upon the difference in sedimentation velocity of the components in a centrifugal field.

 

Various chemical reactions take place according to the certain laws, known as the Laws of chemical combination. These are as follows,

1. Law of conservation of mass : It was proposed by Lavoisier and verified by Landolt. According to this law, Matter is neither created nor destroyed in the course of chemical reaction though it may change from one form to other. The total mass of materials after a chemical reaction is same as the total mass before reaction.

Example : A reaction between

AgNO3

solution and KI solution.

AgNO3 (aq)

• KI(aq)

¾¾®  AgI + NaNO3 (aq)  (yellow ppt.)

Mass of

AgNO3 (aq)

• Mass of KI(aq)

= Mass of the ppt. of AgI

• Mass of NaNO3 (aq)

According to the modified statement of the law, The total sum of mass and energy of the system remains constant.

2. Law of constant or definite proportion : It was proposed by Proust. According to this law, A pure chemical compound always contains the same elements combined together in the fixed ratio of their weights whatever its methods of preparation may be.

Example :

CO2

can be formed by either of the following processes:

1. By heating CaCO3   :

Ca CO3

¾¾D ®

Ca O + CO2

2. By heating

NaHCO3    :

2 NaHCO3

¾¾D ®  Na   CO   + H  O  +  CO

 

2         3          2                  2

CO2

 

is collected separately as a product of each reaction and the analysis of

CO2

of each collection reveals

that it has the combination ratio of carbon and oxygen as 12 : 32 by weight.

3. Law of multiple proportion : It was proposed by Dalton and verified by Berzelius. According to this law, When two elements A and B combine to form more than one chemical compounds then different weights of A, which combine with a fixed weight of B, are in proportion of simple whole numbers.

 

Example : Nitrogen forms as many as five stable oxides. The analysis of these oxides (N 2 O, NO, N 2 O3 , N 2 O4

 

and

N 2 O5 ) reveals that for 28 gm. nitrogen, the weight of oxygen that combines is in the ratio 16 : 32 : 48 : 64 :

80 i.e., 1 : 2 : 3 : 4 : 5 in

N 2 O, NO, N 2 O3 , N 2 O4  and

N 2 O5 respectively.

4. Law of equivalent proportion or law of reciprocal proportion : It was proposed by Ritcher. According to this law, The weights of the two or more elements which separately react with same weight of a third element are also the weights of these elements which react with each other or in simple multiple of them.

Example : Formation of

H 2 S, H 2 O

and SO2

can be done as follows,

1. Hydrogen combines with sulphur forming hydrogen sulphide; 2gm. of hydrogen reacts with 32gm of sulphur. (ii) Hydrogen combines oxygen forming water; 2 gm. of hydrogen reacts with 16 gm. of oxygen. (iii) Sulphur combines with oxygen forming sulphur dioxide; 32 gm. of sulphur reacts with 32 gm. of oxygen i.e., in the ratio 32 : 32. This ratio is double of the ratio weights of these elements which combine with 2 gm. of hydrogen. i.e., 32/16 : 32/32 = 2 : 1

Law of Reciprocal proportion can be used to obtain equivalent weights of elements. As elements always combine in terms of their equivalent weights.

5. Gay-Lussac’s Law: It was proposed by Gay–Lussac and is applicable only for gases. According to this law, When gases combine, they do so in volumes, which bear a simple ratio to each other and also to the product

formed provided all gases are measured under similar conditions. The Gay-Lussac’s law, was based on experimental observation.

Example : (i) Reaction between hydrogen and oxygen.

H 2 (g )

• 1 O

2

 

 

2 (g )

¾¾® H 2 O

 

 

(v)

One volume of

H₂ reacts with half volume of O2 to give one volume of

H 2 O .

(ii) Reaction between nitrogen and hydrogen.

N 2 (g )

+ 3 H 2 (g )

® 2 NH 3 (g )

One volume of

N₂ reacts with three volumes of

H 2 to give two volumes of

NH 3 .

 

1. Atomic hypothesis : Keeping in view various law of chemical combinations, a theoretical proof for the validity of different laws was given by John Dalton in the form of hypothesis called Dalton's atomic hypothesis. Postulates of Dalton's hypothesis is as followes,
   1. Each element is composed of extremely small particles called atoms which can take part in chemical combination.
   2. All atoms of a given element are identical i.e., atoms of a particular element are all alike but differ from atoms of other element.
   3. Atoms of different elements possess different properties (including different masses).
   4. Atoms are indestructible i.e., atoms are neither created nor destroyed in chemical reactions.
   5. Atoms of elements take part to form molecules i.e., compounds are formed when atoms of more than one element combine.
   6. In a given compound, the relative number and kinds of atoms are constant.

 

 

2. Modern atomic hypothesis : The main modifications made in Dalton’s hypothesis as a result of new discoveries about atoms are,

1. Atom is no longer considered to be indivisible.
2. Atoms of the same element may have different atomic weights. e.g., isotopes of oxygen O16 , O17 , and O18 .
3. Atoms of different element may have same atomic weights. e.g., isobars Ca 40 and Ar 40 .
4. Atom is no longer indestructible. In many nuclear reactions, a certain mass of the nucleus is converted into energy in the form of a, b and g rays.
5. Atoms may not always combine in simple whole number ratios. e.g., in sucrose (C12 H 22 O11 ) , the elements carbon, hydrogen and oxygen are present in the ratio of 12 : 22 : 11 and the ratio is not a simple whole number ratio.

3. Berzelius hypothesis : “Equal volumes of all gases contain equal number of atoms under same conditions of temperature and pressure”. When applied to law of combining volumes, this hypothesis predicts that atoms are divisible and hence it is contrary to Dalton's hypothesis.
4. Avogadro’s hypothesis : “Equal volumes of all gases under similar conditions of temperature and pressure contain equal number of molecules.” Avogadro hypothesis has been found to explain as follows :

1. Provides a method to determine the atomic weight of gaseous elements.
2. Provides a relationship between vapour density (V.D.) and molecular masses of substances. Vapour density = Volume of definite amount of a gas

Volume of same amount of hydrogen

or   Vapour denstiy =    Mass of 'n' molecule of a gas

Mass of 'n' molecule of hydrogen

 

Vapour density =

 

or

Mass of 1 molecule of a gas Mass of 1 molecule of hydrogen

 

Molecular mass = 2 ´ vapour density

or

 

 

3. It helps in the determination of m