Basic Concepts of General Chemistry
Anything that exhibits inertia is called matter.
The quantity of matter is its mass.
Matter exists in three different physical forms i.e, solid, liquid and gas.
Property 
Solid 
Liquid 
Gas 
Tightness 
Very tightly packed 
Tightly packed 
Loosely packed 
Intermolecular space 
Minimum 
Intermediate 
Maximum 
Force of attraction 
Maximum 
Intermediate 
Minimum 
Kinetic Energy 
Minimum 
Intermediate 
Maximum 
Density 
Maximum 
Intermediate 
Minimum 
Volume 
Fixed 
Fixed 
Variable 
Shape 
Fixed 
Variable 
Variable 
Compressibility factor 
Minimum 
Intermediate 
Maximum 
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 nonelemental pure compound.
 Formed by chemical combination of two or more atoms of different elements in a fixed ratio.
 Examples: H_{2}O, CO_{2}, C_{6}H_{12}O_{6} 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 
m^{2} 
Volume (V) 
Cubic metre 
m^{3} 
Density (r) 
Kilogram m^{3} 
Kg m^{3} 
Energy (E) 
Joule (J) 
Kg m^{2}s^{2} 
Force (F) 
Newton (N) 
Kg ms^{2} 
Frequency (n) 
Hertz 
Cycle per sec 
Pressure (P) 
Pascal (Pa) 
Nm^{2} 
Electrical charge 
Coulomb (C) 
As (ampere – second) 
Measurement of Temperature
Three scales of temperature
 Kelvin scale (K)
 Degree Celsius scale (^{o}C)
 Degree Fahrenheit scale (^{o}F)
Relations between the scales:
 ^{ o}F = 9/5(^{o}C) + 32
 K = ^{o}C + 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 nonzero digits are significant.
 Zeroes preceding the first nonzero digit are not significant.
 Zeroes between two nonzero 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 × 10^{n} form.
Where,
 N = Digit term
 n = exponent having positive or negative value.
 Examples,
12540000 = 1.254 × 10^{7}
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 ×10^{3 }) × (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 ×10^{3 }) / (8.0×10^{7 }) = ( 7.0/8.0) × ( 10^{[3  (7)] }) = 0.875 ×10^{10 }= 0.9 ×10^{10 }
Addition and Subtraction:
Numbers are written in such way that they have same exponent and after that coefficients are added or subtracted.
(5 ×10^{3 }) + (8×10^{5 }) = (5 ×10^{3 }) + (800×10^{3 }) = (5+800) ×10^{3 }= 805×10^{3}
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.
Laws of Chemical Combination:
Law of conservation of mass
 This law was put forth by Antoine Lavoisier in 1789.
 Law of conservation of mass states that mass can neither be created nor destroyed in a chemical reaction.
 100 grams of ice cubes in a glass will become liquid when allowed to melt. According to this law, the mass of both the starting and ending substance will be equal.
Law of definite proportions
 This law was given by, a French chemist, Joseph Proust.
 He stated that a given compound always contains exactly the same proportion of elements by weight.
 Proust worked with two samples of cupric carbonate — one of which was of natural origin and the other was synthetic. He found that the composition of elements present in it was same for both the samples as shown below:

%of copper
%of carbon
%oxygen
Natural Sample
51.35 9.74 38.91 Synthetic Sample
51.35 9.74 38.91  Thus, he concluded that irrespective of the source, a given compound always contains same elements combined together in the same proportion by mass.
 It is sometimes also referred to as Law of Definite Composition.
 In both Sample A and B, the ratio of the mass of hydrogen to the mass of oxygen is always 1:8.
 Thus, if 9 g of water is decomposed, 1 g of hydrogen and 8 g of oxygen are always obtained.
Law of Multiple Proportions
 This law was proposed by Dalton in 1803.
 According to this law, if two elements can combine to form more than one compound, the masses of one element that combine with a fixed mass of the other element, are in the ratio of small whole numbers.
 For example, hydrogen combines with oxygen to form two compounds, namely, water and hydrogen peroxide.
 Here, the masses of oxygen (i.e., 16 g and 32 g), which combine with a fixed mass of hydrogen (2g) bear a simple ratio, i.e., 16:32 or 1: 2.
Gay Lussac’s Law of Gaseous Volumes
 This law was given by Gay Lussac in 1808.
 When gases combine or are produced in a chemical reaction they do so in a simple ratio by volume, provided all gases are at the same temperature and pressure.
 Thus, 100 mL of hydrogen combine with 50 mL of oxygen to give 100 mL of water vapour.
Avogadro’s Law
 Avogadro proposed that equal volumes of all gases at the same temperature and pressure should contain equal number of molecules
Law of Reciprocal Proportions –
 The law of reciprocal proportions was proposed by Jeremias Ritcher in 1792.
 It states that, "If two different elements combine separately with the same weight of a third element, the ratio of the masses in which they do so are either the same or a simple multiple of the mass ratio in which they combine."
Example –
 Methane CH_{4}  Ratio of weight of carbon and hydrogen is 12:4 or 3:1
 Water, H_{2}O  Ratio of weight of Oxygen to Hydrogen is 16:2 or 8:1.
 So, methane and water both contain a hydrogen and one other element. According to this law if we combine carbon and oxygen (the other element in both compounds) it should be in a ratio of 3:8, or a simple multiple of that ratio.
Atomic and Molecular Masses:
Atomic mass
 An atomic mass unit (symbolized AMU or amu) is defined as precisely 1/12 the mass of an atom of carbon12.
 Atomic mass refers to the mass of an atom.
 It depicts how many times an atom of an element is heavier than 1/12^{th} the mass of one atom of carbon12.
 The relative atomic masses of all elements have been established with reference to an atom of carbon12.
 Atomic mass is equal to the sum of number of protons and neutrons in an atom.
NAME 
SYMBOL 
ATOMIC NUMBER 
PROTONS 
NEUTRONS 
MASS NUMBER 
Hydrogen 
H 
1 
1 
0 
1 
Helium 
He 
2 
2 
2 
4 
Lithium 
Li 
3 
3 
4 
7 
Berylium 
Be 
4 
4 
5 
9 
Boron 
B 
5 
5 
6 
11 
Carbon 
C 
6 
6 
6 
12 
Nitrogen 
N 
7 
7 
7 
14 
Oxygen 
O 
8 
8 
8 
16 
Fluorine 
F 
9 
9 
10 
19 
Neon 
Ne 
10 
10 
10 
20 
Molecular mass
 The molecular mass of a substance is the sum of the atomic masses of all the atoms in a molecule of the substance.
 It is expressed in atomic mass units (u).
Formula unit mass:
 The formula unit mass of a substance is a sum of the atomic masses of all atoms in a formula unit of a compound.
 For example, sodium chloride has a formula unit NaCl.
 Its formula unit mass can be calculated as– 1 × 23 + 1 × 35.5 = 58.5 u
Equivalent weight
Equivalent weight = molecular weight or atomic weight / n –factor
n Factor for Acids ie., Basicity
 Number of ionisable H^{+} per molecule is the basicity of acid.
 Basicity of HCL=1
 Basicity of H_{2}SO_{4}=2
 Basicity of H_{3}PO_{4}=3
Example – equivalent weight of H_{2}SO_{4} = 98 / 2 = 49 g / mol
n Factor for Bases ie., Acidity
 Number of ionizable per molecule is the acidity of a base.
 Acidity of NaOH=1
 Acidity of Mg(OH)_{2}=2
 Acidity of Al(OH)_{3}= 3
Example  equivalent weight of NaOH = 40/1 = 40 g/mol
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
Average Atomic mass
When an element exist as isotopes, i.e. atoms of same element with different atomic mass, we need to calculate average atomic mass.
 Chlorine occurs in nature in two isotopic forms of masses and in the ratio of respectively.