This article provides you the revision notes on Class 11 Chemistry: Chapter- Some Basic Concepts of Chemistry, to give you a quick glance of the chapter. This article is a continuation of the revision notes on Class 11 Chemistry, Chapter- Some Basic Concepts of Chemistry, Part-I. In Part-I you got to learn about matter and its classification, basic physical quantities and their measurement.
Topics covered in this part of chapter notes for ‘Some Basic Concepts of Chemistry’, are:
- Laws of Chemical Combinations
- Dalton's Atomic Theory
- Atomic and Molecular Masses
- Mole Concept
- Empirical Formula for Molecular Formula
- Concentration of the solutions
The key notes of the chapter are as follows:
Laws of Chemical Combinations
There are 5 basic laws governing combination of elements to form compounds.
1. Law of Conservation of Mass (Antoine Lavoisier in 1789)
This law states that matter (mass) can neither be created nor destroyed.
2. Law of Definite Proportions (Louis Proust in 1799)
This law states that a chemical compound always consists of the same elements combined together in the same ratio, irrespective of the method of preparation or the source from where it is taken.
3. Law of Multiple Proportions (Dalton in 1803)
This law states that when two elements combine to form two or more compounds, then the different masses of one element, which combine with a fixed mass of the other, bear a simple ratio to one another.
4. Gay Lussac’s Law of Gaseous Volumes (Lussac in 1808)
This law states that under similar conditions of temperature and pressure, whenever gases react together, the volumes of the reacting gases as well as products bear a simple whole number ratio.
H2 (g) + O2(g) → H2O (g)
100 mL 50 mL 100 mL
Here, the volumes of hydrogen and oxygen which combine together (i.e. 100 mL and 50 mL) bear a simple ratio of 2:1.
5. Avogadro Law (Avogadro in 1811)
According to this law equal volumes of gases at the same temperature and pressure should contain equal number of molecules.
Dalton's Atomic Theory:
Basic postulates of Dalton's Atomic Theory are:
- All substances are made up of tiny, indivisible particles called atoms.
- Atoms of a given element are identical in properties like shape, size, mass and other properties.
- Atoms of different elements differ in properties.
- Atom is the smallest unit that takes part in chemical combinations.
- Atoms can neither be created nor destroyed during any physical or chemical change.
- Atoms combine with each other in simple numerical ratios to form compound atoms called molecules.
It is the smallest particle of an element, which may or may not have independent existence is called an atom. For example, oxygen (O), hydrogen (H), etc.
It is the smallest particle of a substance which is capable of independent existence. For example, H2O, O2, etc.
Atomic and Molecular Masses
- It is the mass of an atom.
- It is represented by 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. And 1 amu = 1.66056×10–24 g.
Gram Atomic Mass:
Atomic mass of an element expressed in grams is the gram atomic mass or gram atom
- It is the 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:
- It is the 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
- It is a unit of amount of substance.
- One mole amount of a substance contains the same number of chemical units (atoms, molecules, ions or electrons) as there are atoms in exactly 12 grams of pure carbon-12.
- A mole represents a collection of 6.022 x1023( Avogadro's number) chemical units.
- It is the mass of one mole of a substance in gram.
- It is the volume occupied by one mole of a substance.
- It is the mass percentage of each constituent element present in any compound.
Empirical Formula for Molecular Formula
- It represents the smallest whole number ratio of the constituent atom within the molecule.
- For example, CH is the empirical formula of benzene.
- It represents the actual number of each individual atom in any molecule.
- For example, C6H6 is the molecular formula of benzene.
Relationship between empirical and molecular formulae:
- Molecular formula = n × Empirical formula
The representation of a chemical change in terms of symbols and formulae of the substances involved in the reaction is called chemical equation.
It is the reactant which gets consumed first or limits the amount of product formed.
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
- It is the mass of the solute in grams per 100 grams of the solution.
- It is the volume of the solute per 100 units of the volume of solution.
Parts per million ( ppm):
- It is the amount of the solute in gram per million (106) gram of the solution.
- It is the 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.
- It is the number of moles of solute dissolved per litre (dm3) of the solution.
- It is the number of moles of solute present in 1 kg of solvent.
Try the following questions:
1. Determine the empirical formula of an oxide of iron which has 69.9% iron and 30.1% dioxygen by mass.
2. Calculate the concentration of nitric acid in moles per litre in a sample which has a density, 1.41 g mL–1 and the mass per cent of nitric acid in it being 69%.
3. What is the symbol for SI unit of mole? How is the mole defined?
4. The density of 3 molal solution of NaOH is 1.110 g mL–1. Calculate the molarity of the solution.
5. Calculate the mass of ammonia produced if 2.00 × 103 g dinitrogen reacts with 1.00 × 103 g of dihydrogen according to the following chemical equation:
N2(g) + H2(g) → 2NH3(g)