Anything that exhibits inertia is called matter.
The quantity of matter is its mass.
Classification of Matter:
Based on chemical composition of various substances.
PHYSICAL QUANTITIES AND THEIR MEASUREMENT
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
Name of the unit
Symbol of the unit
Amount of substance
These units are the function of more than one fundamental unit
Quantity with Symbol
Metre per sec
Cycle per sec
A-s (ampere – second)
MEASUREMENT OF TEMPERATURE
Three scales of temperature
Relations between the scales:
0 K temperatures is called absolute zero.
Dalton’s Atomic Theory:
Precision and Accuracy:
Numbers are represented in N × 10n form.
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:
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
Formula Unit Mass
Mass percentage of an element in a compound = (Mass of that element in the compound/Molecular mass of the compound)×100
EMPIRICAL FORMULA AND MOLECULAR FORMULA
Represents the actual number of each individual atom in any molecule is known as molecular formula.
Expresses the smallest whole number ratio of the constituent atom within the molecule.
Molecular formula = (Empirical formula)n
Molecular weight = n × Empirical weight
Molecular weight = 2 × Vapour density
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:
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
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
Number of moles of solute per 1000 mL of the solution.
M = (Number of moles of solute)/(Volume of solution in L)
number of moles of solute per 1000 gram of the solvent.
m = (Number of moles of solute)/(Weight of solvent in kg)