Matter:
Anything that exhibits inertia is
called matter.
The quantity of matter is its mass.
Classification of Matter: -
Based on chemical composition of
various substances.
Elements:
·
It is the simplest
form of the matter.
·
Smallest unit of an
element is known as atom.
·
Total number of the
known elements is 118 out of which 98 elements occur naturally and 20 are
formed by artificial transmutation.
·
Examples: Na, K,
Mg. Al, Si, P, C, F, Br etc.
Compound:
·
It is a
non-elemental pure compound.
·
Formed by chemical
combination of two or more atoms of different elements in a fixed ratio.
·
Examples: H2O,
CO2, C6H12O6 etc.
Mixture:
·
Formed by physical
combination of two or more pure substances in any ratio.
·
Chemical identity
of the pure components remains maintained in mixtures.
·
Homogeneous
mixtures are those whose composition for each part remains constant.
·
Example, Aqueous
and gaseous solution.
·
Heterogeneous
mixtures are those whose composition may vary for each and every part.
·
Example, Soil and
concrete mixtures.
Physical Quantities and Their
Measurement:
Fundamental Units:
-
These units can
neither be derived from one another nor can be further resolved into any other
units. Seven fundamental units of the S.I. system
|
Physical
quantity
|
Name
of the unit
|
Symbol
of the unit
|
|
Time
|
Second
|
S
|
|
Mass
|
Kilogram
|
kg
|
|
Length
|
Meter
|
m
|
|
Temperature
|
Kelvin
|
K
|
|
Electric
current
|
Ampere
|
A
|
|
Luminous
intensity
|
Candela
|
Cd
|
|
Amount
of substance
|
Mole
|
Mol
|
Derived Units: -
These units are
the function of more than one fundamental unit
|
Quantity
with Symbol
|
Unit
(S.I.)
|
Symbol
|
|
Velocity
(v)
|
Metre
per sec
|
ms-1
|
|
Area
(A)
|
Square
metre
|
m2
|
|
Volume
(V)
|
Cubic
metre
|
m3
|
|
Density
(r)
|
Kilogram
m-3
|
Kg
m-3
|
|
Energy
(E)
|
Joule
(J)
|
Kg
m2s-2
|
|
Force
(F)
|
Newton
(N)
|
Kg
ms-2
|
|
Frequency
(n)
|
Hertz
|
Cycle
per sec
|
|
Pressure
(P)
|
Pascal
(Pa)
|
Nm-2
|
|
Electrical
charge
|
Coulomb
(C)
|
A-s
(ampere – second)
|
Measurement of Temperature
Three scales of
temperature
·
Kelvin scale (K)
·
Degree Celsius
scale (oC)
·
Degree Fahrenheit scale
(oF)
Relations between
the scales:
·
oF = 9/5(oC) + 32
·
K = oC +
273
0 K temperatures is
called absolute zero.
Dalton’s Atomic Theory:
·
Every matter
consists of indivisible atoms.
·
Atoms can neither
be created nor destroyed.
·
Atoms of a given
element are identical in properties
·
Atoms of
different elements differ in properties.
·
Atoms of different
elements combine in a fixed ratio to form molecule of a compound.
Precision and
Accuracy:
·
Precision: Closeness of outcomes of different
measurements taken for the same quantity.
·
Accuracy: Agreement of experimental value to the
true value
Significant figures:
Rules:
·
All non-zero digits
are significant.
·
Zeroes preceding
the first non-zero digit are not significant.
·
Zeroes between two
non-zero digits are significant.
·
Zeroes at the end
of a number are significant when they are on the right side of the decimal
point.
·
Counting numbers of
objects have infinite significant figures.
Scientific Notation:
Numbers are
represented in N × 10n form.
Where,
·
N = Digit
term
·
n = exponent having
positive or negative value.
·
Examples,
12540000 = 1.254 × 107
0.00456 = 4.56 ×10-3
12540000 = 1.254 × 107
0.00456 = 4.56 ×10-3
Mathematical Operations of Scientific
Notation:
Multiplication and
Division:
Follow the same
rules which are for exponential number.
Example: (7.0 ×103 ) ×
(8.0×10-7 ) = ( 7.0×8.0) × ( 10[3 + (-7)] ) = 56.0 × 10-4
Result cannot have
more digits to the rite of decimal point than either of the original numbers
(7.0 ×103 ) /
(8.0×10-7 ) = ( 7.0/8.0) × ( 10[3 - (-7)] ) = 0.875 ×1010 =
0.9 ×1010
Addition and
Subtraction:
Numbers are written
in such way that they have same exponent and after that coefficients are added
or subtracted.
(5 ×103 ) +
(8×105 ) = (5 ×103 ) + (800×103 ) = (5+800) ×103 =
805×103
Result must be
reported with no more significant figures as there in the original number with
few significant figures.
Rules for limiting
the result of mathematical operations:
·
If the rightmost
digit to be removed is more than 5, the preceding number is increased by one.
·
If the rightmost
digit to be removed is less than 5, the preceding number is not changed.
·
If the rightmost
digit to be removed is 5, then the preceding number is not changed if it is an
even number but is increased by one if it is an odd number.
Dimensional Analysis: -
·
This is based on
the fact that ratio of each fundamental quantity in one unit with their
equivalent quantity in other unit is equal to one.
·
Derived unit first
expressed in dimension and each fundamental quantities like mass length time
are converted in other system of desired unit to work out the
conversion factor
·
Original Quantity ×
Conversion factor = Equivalent Quantity
(In former unit)
(In final Unit)
Example: - (1 kg/2.205 pound) =
1=(1kg/1000gm)
So 1 kg = 2.205 pound = 1000 gm
Laws of Chemical Combination:
Law of conservation
of mass:
“For any chemical
change total mass of active reactants are always equal to the mass of the product
formed”
Law of constant
proportions:
“A chemical
compound always contains same elements in definite proportion by mass and it
does not depend on the source of compound”.
Law of multiple
proportions:
“When two
elements combine to form two or more than two different compounds then the
different masses of one element B which combine with fixed mass of the other
element bear a simple ratio to one another”
Law of reciprocal
proportion:
“If two elements B
and C react with the same mass of a third element (A), the ratio in which they
do so will be the same or simple multiple if B and C reacts with
each other”.
Gay Lussac’s law of
combining volumes:
“At given
temperature and pressure the volumes of all gaseous reactants and products bear
a simple whole number ratio to each other”.
Atomic and Molecular Masses:
Atomic Mass:
·
Mass of an atom.
·
Reported in atomic
mass unit “amu” or unified mass “u”
·
One atomic mass
unit i.e. amu, is the mass exactly equal to one-twelfth the mass of one
carbon-12 atom.
Molecular Mass:
·
Mass of a molecule
of covalent compound.
·
It is equal to the
sum of atomic masses of all the elements present in the molecule.
Formula Unit Mass
·
Mass of a molecule
of an ionic compound
·
It is also equal to
the sum of atomic masses of all the elements present in the molecule
Mole Concept:
Mole:
·
Unit of amount of
substance.
·
One mole amount of
substance that contains as many particles or entities as there are atoms
in exactly 12 g of the 12C isotope.
Molar mass:
·
Mass of one mole of
a substance in gram
·
Molar mass in gram
in numerically equal to atomic/molecular/formula mass in amu or u.
? Percentage
composition:
Mass percentage of
an element in a compound = (Mass of that element in the compound /
Molecular mass of the compound)×100
Percentage yield:
·
It is the ratio of
actual yield of the reaction to the theoretical yield multiplied by 100.
·
% yield = (Actual
yield /Theoretical yield) ×100
Empirical formula and molecular
formula:
Molecular Formula:
-
Represents the actual number of each
individual atom in any molecule is known as molecular formula.
Empirical Formula:
-
Expresses
the smallest whole number ratio of the constituent atom within the
molecule.
Molecular formula = (Empirical formula)n
Molecular weight = n × Empirical weight
also,
Molecular weight = 2 × Vapour
density
Limiting Reagent:
The reactant which
is totally consumed during the course of reaction and when it is consumed
reaction stops.
For a balanced
reaction reaction:
A +B → C + D
B would be a
limiting reagent if nA / nB>nB/nA
Similarly, A is a
limiting reagent if nA / nB<nB/nA
Concentration of the solutions:
Mass by Mass
Percentage:-
Amount of solute in gram present per
100 gm of the solution.
Mass percentage of solute = [(Mass of
solute)/ (Mass of solution)] x100
Mass by Volume
Percentage:-
Amount solute in gram present per 100
mL of the solution.
Volume by Volume
Percentage:-
Volume of solute per 100 mL of the
solution
Volume by volume percentage of solute =
[(Volume of solute)/(volume of solution)] x100
Parts per million (
ppm) :-
The
amount of solute in gram per million (106) gram of the solution.
ppm = [(mass of solute/mass of
solution)]x 106
Mole fraction:-
Ratio of the moles of one component of
the solution to the total number of moles of solution
Total mole fraction of all the
components of a solution is equal to 1.
For binary solutions having two
components A and B
Mole fraction of A
XA = (nA)/(nA+nB)]
Mole fraction of B
XB =
(nB)/(nA+nB)]
or XB =
1- XA
Molarity(M):-
Number of moles of
solute per 1000 mL of the solution.
M = (Number
of moles of solute)/(Volume of solution in L)
Molality(m): -
number of moles of
solute per 1000 gram of the solvent.
m = (Number of
moles of solute)/(Weight of solvent in kg)
0 comments:
Post a Comment