Mole Concepts

Dalton’s atomic theory

1. Matter consists of individual atoms.

2. All the atoms of a given element have identical properties including identical mass. Atoms of different elements differ in mass.

3. Compounds are formed when atoms of different elements combine in different ratio.

4.Chemical reactions involve reorganization of atoms. These are neither created nor destroyed in a chemical reaction.

Laws of chemical combination

1. (Lavoisier) the law of conservation of matter or mass states that, in a chemical reaction, the total amount of matter of the reaction compounds remains constant.

2. Proust's law states that the elements in a compound are all present in a fixed proportion by weight, regardless of how the compound is prepared.

3.the law of multiple proportions: When two elements combine to form more than one compound, the element whose mass varies combines with a fixed mass of the second element weights in a simple whole-number ratio such as 2:1, 3:1, or 3:2.

4.Gay-Lussac's law of combining volumes states that the volumes of the gases involved in a chemical reaction (both reactants and products) are in the ratio of small whole numbers.

Mole Concept

Mole (measurement), a base unit of the International System of Units, is defined as the amount of a substance that contains as many elementary particles (atoms, molecules, ions, electrons, or other particles) as the number of carbon atoms in 0.012 kg (12 g) of carbon-12.

1 mole= collection of 6.02 × 1023species (Avogadro's number) = NA = 1 gm-specie

Atomic weight

An element’s atomic weight is given in grams. It represents the mass of one mole (6.02 × 1023 atoms) of that element. Numerically, the atomic mass and the atomic weight of an element are the same, but the first is expressed in grams and the second is in atomic mass units. So, the atomic weight of hydrogen is 1 gram and the atomic mass of hydrogen is 1 amu.

The chemical atomic weight (atomic wt.) of an element is the weighted average of the individual atomic weights, or mass numbers, of the isotopes. For example, chlorine, atomic wt. 35.457, is composed of chlorine-35 and chlorine-37, the former occurring with an abundance of 76% and the latter of 24%.

?Molecular formula = empirical formula × n

?Molecular weight is twice the vapour density (ratio of density of vapour to the density of H2 at similar P & T).

?Limiting reagent = moles/stoichiometric coefficient

?Percentage yield = (actual yield/the max theoretical yield) × 100

?Specific gravity = density of any substance/density of water at 4?C

?Molarity, molality, mole fraction, mole %, % w/v, normality

Typical concentration terms

1.Oleum: e.g. 102% =max amount of H2SO4 (102 g) that can be obtained from 100 g of this sample by adding 2 g water.

2. Hydrogen peroxide (H2O2): e.g. 20 V H2O2 implies 20 L O2 can be obtained from 1 L of this sample.

(H2O2→ H2O + ½ O2)

Eudiometry

Solvent gas(es) absorbed

KOH CO2, SO2, Cl2

Ammon Cu2Cl2 CO

Turpentine oil O3

Alkaline pyrogallol O2

Water NH3, HCl

CuSO4/CaCl2 H2O

EQUIVALENT CONCEPT

No. of equivalents = moles × n

Eq wt = M.M./n

n factor of ionic compounds = total +ve/-ve charge

n factor of some compounds

1. KMnO4 acidic= 5

Strongly Basic = 1

Neutral/basic = 3

2. K2CrO4 = 6

3. C2O42- = 2