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 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

m

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

 Mathematical Operations of Scientific Notation:

 Multiplication and Division:

Follow the same rules which are for exponential number.

Example: (7.0 ×10) × (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) / (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 ×10) + (8×10) = (5 ×10) + (800×10) = (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.

 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 CH4  - Ratio of weight of carbon and hydrogen is 12:4 or 3:1
  • Water, H2O - 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 carbon-12.
  • Atomic mass refers to the mass of an atom. 
  • It depicts how many times an atom of an element is heavier than 1/12th the mass of one atom of carbon-12.
  • The relative atomic masses of all elements have been established with reference to an atom of carbon-12. 
  • 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 H2SO4=2
  • Basicity of H3PO4=3

Example – equivalent weight of H2SO4 = 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.