Notes and Study Materials -Some Basic Concepts of Chemistry

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About this unit

UNIT I: Some Basic Concepts of Chemistry Laws of chemical combination, Dalton’s atomic theory: concept of elements, atoms and molecules. Atomic and molecular masses. Mole concept and molar mass; percentage composition and empirical and molecular formula; chemical reactions, stoichiometry and calculations based on stoichiometry.

PHYSICAL QUANTITIES AND SI UNITS

The 11th general conference of weights and measures in 1960 recommended the use of international system of units. Abbreviated as SI Units (after the French expression La System International de units).

FUNDAMENTAL UNITS

The SI system has seven basic units of physical quantities as follows:

 

Physical Quantity
Abbreviation
Name of unit
Symbol
time
t
second
s
mass
m
kilogram
kg
length
l
metre
m
temperature
T
kelvin
K
electric current
I
ampere
A
light intensity
Iv
candela
Cd
amount of substance
n
mole
mol

DERIVED UNITS

The units obtained by combination of basic units are known as derived units e.g. velocity is expressed as distance/time. Hence unit is m/s or ms–1. Some common derived units are:

 

Physical Quantity
Definition
SI Unit
volume
length cube
m3
area
length square
m2
speed
distance travelled
ms–1
per unit time
acceleration
speed changed
ms–2
per unit time
density
mass per unit volume
kg m–3
pressure
force per unit area
kgm–1s–2 or Nm–2 (pressure = Pa)
force
mass times acceleration of object
kgms–2 (Newton N)
energy
force times distance  travelled
kgm2s–2(Joule J)
frequency
cycles per second
s–1 (hertz = Hz)
power
energy per second
kgm2s–3 or Js–1
(Watt = W)
electric charge
ampere times second
As (coloumb = C)
electric potential
energy per unit
JA–1s–1 or kgm2s–3
difference
charge
A–1 (volt = V)

SOME NON-SI UNITS IN COMMON USE

Quantity
Unit
Symbol
SI definition
SI Name
Length
angstrom
Å
10–10 m
0.1 nanometers (nm)
Volume
litre
L
10–3 m3
1 decimeter (dm3)
Energy
calorie
cal
kg m2s–2
4.184 Joule (J)

STANDARD PREFIXES FOR EXPRESSING THE DECIMAL FRACTIONS OR MULTIPLES OF FUNDAMENTAL UNITS

Fraction
Prefix
Symbol
Multiple
Prefix
Symbol
10–1
deci
d
101
Deka
da
10–2
centi
c
102
Hecta
h
10–3
milli
m
103
kilo
k
10–6
micro
m
106
Mega
M
10-9
nano
n
109
Giga
G
10–12
pico
p
1012
Tera
T
10–15
femto
f
1015
Peta
P
10–18
atto
a
1018
Exa
E

PRECISION AND ACCURACY

Precision : It is the closeness of various measurements for the same quantity.
Accuracy : It is the agreement of a particular value to the true value.

 

Example : Let the true weight of a substance be 3.00g. The measurement reported by three students are as follows
Case of A student : It is precision but no accuracy since measurements one close but not accurate.
Case of B student : Measurements are close (precision) and accurate (Accuracy)
Case of C student : Measurement are not close (no precision) and not accurate (no accuracy)

STOICHIOMETRY

Stoichiometry : It is calculation of masses or volumes of reactants and products involved in a chemically balanced reaction. Consider the formation of ammonia.
N2 (g) + 3H2 (g)   2NH3(g)

 

All are gases indicated by letter (g) and coefficients 3 for H2 and 2 for NH3 are called stoichiometric coefficients. The formation of ammonia can be interpreted in many ways:
  • One mole of N2(g) reacts with three moles of H2(g) to give two moles of NH3(g).
  • 28g of N2(g) reacts with 6g at H2(g) to give 34g of NH3(g).
  • 22.4L of N2(g) reacts with 67.2L of H2(g) to give 44.8L of NH3(g)

SIGNIFICANT FIGURES

The weight 7.52 gm of a substance indicates that it is reliable to the nearest hundredth of a gram and may be expressed as 7.52 ± 0.01. It means slightest variation may occur at the second place of decimal or we can say that uncertainty  is ± 0.01 g.
Now consider the weight 6.4234 g. It may correctly be expressed as 6.4234 ± 0.001 g.
In the first case the weight contains three significant figures and in the second case weight contains five significant figures.
  1. Significance of zero : If zero is used to locate the decimal point it is not considered as significant figure. Thus in 0.0072 there are only two significant figures whereas in 70.40, there are four significant figures since zero is after 4. Again in 0.0070 there are two significant figures, since zero after 7 is significant for it has a meaning when written in exponentials. If we compare 7.0 × 10–3 and 7 × 10–3, the first term has uncertainty of one in seventy and second has uncertainty of one in seven. The exponential term does not add to number of significant figures.
  2. Addition and subtraction of quantities : In this case the uncertainty in the result is equal to the sum of the uncertainties of the individual quantities.
  3. Multiplication and division : In this case the uncertainty in the result is equal to the sum of the percentage of individual uncertainties.
  4. Rounding off : The following rules are observed.
    1. If the digit after the last digit to be retained is less than 5, the last digit is retained as such. e.g. 1.752 = 1.75 (2 is less than 5).
    2. If the digit after the last digit to be retained is more than 5, the digit to be retained is increased by 1. e.g. 1.756 = 1.76 (6 is more than 5).
    3. If the digit after the last digit to be retained is equal to 5, the last digit is retained as such if it is even and increased by 1 if odd.
  5. Calculations involving addition and subtraction : In case of addition and subtraction the final result should be reported to the same number of decimal places as the number with the minimum number of decimal places .
  6. Calculations involving multiplication and division : In this case the final result should be reported having same number of significant digits as that of the number having least significant digits.

MATTER

Anything which occupies space, possesses mass and can be felt is called matter.

CLASSIFICATION OF MATTER

ELEMENT

Pure substance consisting of one type of particles in the form of atoms eg. Cu, Na, Fe or molecules eg. H2, O2 etc.

COMPOUND

Pure substance consisting of molecules formed by the combination of atoms of different elements eg. CO2, H2O etc.

MIXTURES

Mixtures are substances made of two or more elements or compounds in any proportion. They may be homogeneous or heterogeneous.

SEPARATION OF MIXTURES

Mixtures can be separated into constituents by following methods:
  1. Filtration can separate those mixtures whose one component is soluble in a particular solvent and other is not.
  2. Distillation can be used to separate constituents of mixtures having different boiling points.
  3. Extraction dissolves one out of several components of mixture.
  4. Crystallisation is a process of separating solids having different solubilities in a particular solvent.
  5. Sublimation separates volatile solids which sublime on heating from non-volatile solids.
  6. Chromatography is the technique of separating constituents of a mixture which utilises the property of difference of adsorption on a particular adsorbent.
  7. Gravity separation separates constituents having different densities.
  8. Magnetic separation can separate magnetic components from non magnetic ones.

PHYSICAL AND CHEMICAL CHANGES

A change which does not affect chemical composition and molecular structure is a physical change and the one that involves alteration of chemical composition and molecular structure is a chemical change.
  1. Chemical Combination is reaction between two or more elements or compounds to form a single substance.
H2 + I2   2HI
  1. Displacement means replacement of one element of compound by another.
  2. Decomposition involves splitting of a compound to form two or more substances.
  1. Combustion is a complete and fast oxidation of a substance.
  1. Neutralisation is the reaction between acid and base to form a salt.
  1. Polymerisation is the combination of molecules of same or different substances to form a single molecule called polymer

  2. Photochemical changes occur in presence of visible or ultraviolet light.
  3. Double decomposition or metathesis is the exchange of oppositely charged ion on mixing two salt solutions.
  4. Hydrolysis involves reaction of salts with water to form acidic or basic solutions.

LAWS OF CHEMICAL COMBINATIONS

  1. Law of conservation of mass : This law was given by French chemist A. Lavoisier (1774) which states that “during any physical or chemical change, the total mass of products is equal to the total mass of reactants”. It is also called law of indistinctibility. It does not hold good for nuclear reaction.
  2. Law of definite proportions : This law was given by Proust (1799) and states that “a chemical compound always contains some elements combined together in same proportion by mass”. For example different samples of pure CO2 always have carbon and oxygen in 3 : 8 ratio by mass.
  3. Law of multiple proportions : This law was given by John Dalton (1803) and states that “when two elements combine to form two or more compounds, the different mass of one of the elements and the fixed mass of the one with which it combines always form a whole number ratio”. This law explains the concept of formation of more than one compound by two elements.
  4. Law of reciprocal proportions : This was given by Richter (1792) and states that “when two elements combine separately with a fixed mass of third, the ratio of masses in which they do so is same or whole number multiple of the ratio in which they combine with each other.” This law is also called law of equivalent proportions and is helpful in determining equivalent weights.
  5. Gay Lussac’s law of combining volumes : This law states that when gases react with each other, their volumes bear a simple whole no. ratio to one another and to volume of products (if gases) and similar conditions of pressure and temperature.
  6. Dalton’s atomic theory :
    • Proposed by John Dalton in 1808. Main points are :
    • Matter is made up, by indivisible particles called atoms
    • Atoms of same elements are identical in physical and chemical properties.
    • Atoms of different substances are different in every respect
    • Atoms always combine in whole numbers to form compounds
    • Atoms of resultant compounds possess similar properties
Drawbacks of Dalton’s theory :
  • Does not explain structure of atom.
  • Fails to explain binding forces between atoms in compounds.
  • Does not explain Gay Lussac’s law.
  • Does not differentiate between atom and molecule.
  1. Avogadro’s law :
It states that “equal volumes of all gases, under similar conditions of temperature and pressure contain equal number of molecules”. Applications are:
  • Deducing atomicity of elementary gases
  • Deriving relationship between molecular mass and vapour density
  • Deriving formula of substances
  • Determining molecular wt. of a gas
  • Deducing the gram molecular volume.

ATOM

Atom is the smallest particle of element which might not be able to exist independently.

MOLECULE

Molecule is the smallest particle of the substance which can exist independently. It can be subdivided as
  1. Homoatomic molecules are molecules of same element and can be further divided as monoatomic, diatomic and polyatomic molecules depending upon number of atoms. eg: He, O2, P4 etc.
  2. Heteroatomic molecules are molecules of compound. They can be diatomic and polyatomic. eg: H2O, PCl5, H2SO4, NO etc.

ATOMIC MASS UNIT (A.M.U.)

It is the unit of representing atomic masses. 1 a.m.u. = th the mass of C-12.

MOLE

It is a unit which represents 6.023 × 1023 particles. The number 6.023 × 1023 is called Avogadro’s number and is represented by N0 or NA. Avogadro’s number of gas molecules occupy a volume of 22400 cm3 at N.T.P. Number of molecules in 1 cm3 of gas at NTP is Loschmidt N0. With value 2.688 × 1019.

ATOMIC MASS

“It is the number of times the atom of the element is heavier than H atom” was the first proposed definition. Later on oxygen was preferred as standard. In 1961 C-12 was chosen as standard and thus “the number of times the atom of an element is heavier than 12th part of C-12 is called atomic mass of the element.
Atomic mass = 

AVERAGE ATOMIC MASS

It is the mass of each isotope determined separately and then combined in ratio of their occurrence. Suppose a and b are two isotopes of an element with their occurence ratio p : q then
Average atomic mass = 

DETERMINATION OF ATOMIC MASS

 

  1. Dulong and petit’s rule : It is based on experimental facts. “At ordinary temperature, product of atomic mass and specific heat for solid elements is approximately 6.4 and this product is known as atomic heat of the element”
Atomic mass × specific heat = 6.4
The law is valid for solid elements except Be, B, Si and C.
Correct At. mass = Eq. mass × valency
  1. Specific heat method : This method is for gases.  , where Cp = specific heat at constant pressure and Cv = specific heat at constant volume. the ratio g is a constant = 1.66 for monoatomic, 1.40 for diatomic, 1.33 for triatomic gas and atomic mass of gaseous element
=.
  1. Chloride formation method : This method converts the element (whose mass is to be determined) into volatile chloride whose vapour density is found by Victor Mayer method.
Molecular mass = 2 × V.D.
  1. Vapour density method is suitable for elements having volatile chlorides.
Atomic mass = Eq. mass of metal × valency.
  1. Mitscherlich’s law of isomorphism : It states that isomorphous substances have similar chemical constitution. Isomorphous substances form crystals of same shape and valencies of elements forming isomorphous salts are also same. eg: ZnSO4. 7H2O, MgSO4.7H2O and FeSO4.7H2O are isomorphous.

GRAM ATOMIC MASS (GAM)

Is the mass of an atom expressed in gms.
No. of Gm-atoms of element = 

MOLECULAR MASS :

It is the average relative mass of the molecule as compared with mass of C-12 atom.
Molecular mass = 

CALCULATION OF MOLECULAR MASS :

  1. Graham’s law of diffusion : It states that rate of diffusion of two gases is inversely proportional to the square root of ratio of their molecular weights.
  1. Victor meyer method : This method can determine the molecular mass as
Molecular mass = × 22400
where W is the mass of liquid in gm. occupying a volume V ml at STP.
  1. Vapour density  method : Vapour density is the ratio of volume of a gas to the mass of same volume of hydrogen under identical conditions.
or 
Thus molecular mass = 2 × V.D.
  1. Colligative properties method : This method can be helpful in determining molecular mass as
elevation in boiling point 
Where Tb is elevation in b.p., Kb is  molal elevation constant w is wt. of solute W is wt. of solvent
Depression in freezing point 

GRAM MOLECULAR MASS OR MOLAR MASS :

That amount of substance whose mass in grams is equal to its molecular mass or the equivalently molecular mass of a substance expressed in grams is called gram molecular mass. Gram molecular mass is also called one gram molecule. thus
No. of gm molecules = 

EQUIVALENT MASS :

It is the number of parts by weight of the substance that combines or displaces, directly or indirectly, 1.008 parts by mass of hydrogen or 8 parts by mass of oxygen or 35.5 parts by mass of chlorine. It can be calculated as-
  • Equivalent mass for elements = 
  • Equivalent mass for acids = 
  • Equivalent mass for bases = 
  • Equivalent mass for salts = 
  • Equivalent mass for oxidising agents = 
  • Equivalent mass for reducing agents =
  • Equivalent weight of radicals  = 

FORMULA MASS :

It is obtained by adding atomic masses of various atoms present in the formula and this term replaces molecular mass in ionic compounds.

ACIDITY :

It is the number of OH– ions that can be displaced from one molecule of a substance.

BASICITY :

It is the number of H+ ions that can be displaced from one molecule of a substance.

GRAM EQUIVALENT MASS (GEM) :

It is the mass of a substance expressed in grams or equivalently the quantity of substance whose mass in grams is equal to its equivalent mass is called one gram equivalent or gram equivalent mass.
No. of gm equivalents = .

METHODS OF DETERMINING EQUIVALENT MASSES :

  1. Hydrogen displacement method :  It is for metals which can displace H2 from acids.
Equivalent mass of metal
  1. Metal displacement method : It utilises the fact that one GEM of a more electropositive metal displaces one GEM of a less electropositive metal from its salt.
  2. Conversion method : When one compound of a metal is converted to another compound of similar metal then
where E is the eqv. mass of the metal.
  1. Electrolytic method :
    It states that the quantity of substance that reacts at electrode when  Faraday of electricity is passed is equal to its GEM.
GEM = Electrochemical equivalent × 96500
and ratio of weights deposited by equal amount of electricity is in ratio of their equivalent masses.
  1. Oxide method :
Equivalent mass of metal = 

 

  1. Double decomposition :
  1. Neutralisation method for acids and bases :
Equivalent mass of acid (base)  
  1. Silver salt is method commonly used for organic acids.
Eqv. mass of acid = 
Mol. mass of acid = Eqv. mass of acid × Basicity
  1. Platinichloride method for bases :
Eqv. mass of base  
=
Mol. mass of base = Eqv. mass of base × Acidity
  1. Chloride method :
Eqv. mass of metal = 
           
  1. Volatile chloride method :
Valency of metal

CHEMICAL EQUATION :

It is the equation representing chemical change in terms of formula of reactants and products.
  1. An equation which has not been equalised in terms of number of atoms of reactants and products is called a skeleton equation.
  2. An equation having equal number of atoms of various kinds on both sides is a balanced equation.

EMPIRICAL FORMULA :

It is the simplest formula of a compound giving simplest whole number ratio of atoms present in one molecule. e.g. CH is empirical formula of benzene.

MOLECULAR FORMULA :

It is the actual formula of a compound showing the total number of atoms of constituent elements e.g. C6H6 is molecular formula of benzene.
Molecular formula = n × empirical formula, where n is simple whole number.

SOLUTION :

It is a homogenous mixture of two or more substances. The component of solution having larger proportion is solvent and others are solute.

MOLE FRACTION :

It is the ratio of moles of a constituent to the total number of moles in a solution.
Let A be solute & B is solvent then mole fraction of solute (xA)
= , where n is the number of moles.
Mole fraction of solution  

MASS PERCENTAGE :

It is the number of parts by mass of solute per hundred parts by mass of solution. If WA is mass of solute and WB the mass of solvent, then
Mass percentage of A = . 

VOLUME PERCENTAGE :

It is the number of parts by volume of solute per hundred parts by volume of solution. If VA is volume of solute and VB is the volume of solvent then
Volume percentage of A = 

PARTS PER MILLION (PPM) :

It is the mass of solute present in one million parts by mass of solution.

NORMALITY :

It is the number of gram equivalents of a solute present in one litre of solution.
                  
Normality depends on temperature. Also if strength is given in normalities, N1 of A & N2 of B
Then N1V1 = N2V2.

MOLARITY :

It is the number of moles of solute present in one litre of solution.
  
and millimoles = M × V(in ml).
      Molarity and mass percentage have the relation M
   =  , where d = density
If a solution of molarity M1 and volume V1 adds up with a solvent to a final volume V2, then molarity M2 is given by
If two different solutions (M1, V1) and (M2, V2) are mixed then molarity of resulting solution is
M = 
Also, Molarity × GMM of solute = Normality × GEM of solute

MOLALITY :

It is the number of moles of solute in 1 kg of solvent.
Molality (m) = 
                     
Molality is independent of temperature.

FORMALITY (F) :

It is the number of gram formula mass of ionic solute dissolved in 1 litre of solution.
Formality =  

LIMITING REAGENT :

It is the reactant which is completely consumed during the reaction.

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