Introduction to atoms and molecules — lesson. Science TNSB Mentoring, Class 10 Crash course.

Importance:

The chapter "Atoms and molecules" is one of the fundamental chapters in chemistry and carries a weightage of \(5\) to \(8\) marks.

It lays the foundation for understanding chemical calculations, laws of chemical combination, and the quantitative nature of chemical reactions. A clear grasp of this chapter is essential for solving numerical problems and for advanced topics in chemistry.

Question distribution:

Part I (\(1\) mark) - One question
Part II (\(2\) mark) - One question
Part III or IV (\(4\) or \(7\) marks) - One question

(Note: The exact mark distribution may vary slightly across examinations.)

Learning objectives:

Atoms and molecular masses: Understand atoms, molecules, atomic mass, and molecular mass.
Gram mass concepts: Learn gram atomic mass and gram molecular mass.
Avogadro’s hypothesis: Understand its meaning, applications, and gas–volume relationships.
Mole concept: Apply the mole concept to solve numerical problems and determine atomicity.
Chemical calculations: Calculate vapour density and percentage composition of compounds.

Introduction to atoms and molecules:

Matter is present everywhere and is made up of tiny particles called atoms.
Atoms combine to form molecules, which constitute all substances.
The idea of atoms was first proposed by Greek philosophers in the \(5th\) century \(BCE\).
John Dalton gave the first scientific atomic theory.
Later discoveries proved that atoms contain electrons, protons, and neutrons.
Scientists like J. J. Thomson, Rutherford, Bohr, and Schrödinger contributed to the development of the modern atomic theory.
The modern atomic theory explains the structure of atoms and forms the basis of chemical reactions and calculations.

Postulates of modern atomic theory:

Atoms are divisible and consist of subatomic particles.
Atoms of the same element may have different masses (isotopes).
Atoms of different elements may have the same mass (isobars).
Atoms can be converted from one element to another in nuclear reactions (discovery of artificial transmutation).
Atoms combine in fixed but not always simple whole-number ratios.
An atom is the smallest particle that participates in a chemical reaction.
According to \(E = mc^2\), mass can be converted into energy.

The modern atomic theory forms the basis for understanding chemical reactions and calculations.

PYQ - Modern atomic theory

Atomic mass:

Mass number = No. of protons + no. of neutrons

Atomic mass unit (amu) is one-twelfth of the mass of a carbon - \(12\) atom; an isotope of carbon, which contains \(6\) protons and \(6\) neutrons.

Relative atomic mass (RAM):

Relative atomic mass (A_{r}) = \frac{Average mass of the isotopes of the element}{\frac{1}{12} th of the mass of one C - 12 atom}

Relative atomic mass is only a ratio; hence, it does not have a unit. If the atomic mass of an element is expressed in grams, it is called gram atomic mass.

Example:

Gram Atomic Mass of hydrogen \(= 1\) \(g\)
Gram Atomic Mass of carbon \(= 12\) \(g\)
Gram Atomic Mass of nitrogen \(= 14\) \(g\)
Gram Atomic Mass of oxygen \(= 16\) \(g\)

Average atomic mass (AAM):

The average atomic mass of an element is the weighted average of the masses of its naturally occurring isotopes.

Average atomic mass = [\begin{array}{l} (Mass of 1^{st} isotope \times % abundance of 1^{st} isotope) + \\ (Mass of 2^{nd} isotope \times % abundance of 2^{nd} isotope) \end{array}]

To calculate the average atomic mass of carbon, the following formula is used:

\begin{array}{l} Average atomi mass of C \\ = (12 \times \frac{98.9}{100}) + (13 \times \frac{1.1}{100}) \\ = (12 \times 0.989) + (13 \times 0.011) \\ = 11.868 + 0.143 \\ = 12.011 amu \end{array}

It is important to remember that when the atomic mass of carbon is given as \(12\) amu, it refers to the average atomic mass of carbon isotopes, not the mass of individual carbon atoms.

Molecules:

A molecule is a combination of two or more atoms held together by strong chemical forces of attraction, i.e. chemical bonds.
A molecule can be made up of two atoms of the same kind, whereas a compound must contain at least two different elements.
If the molecule is made of a similar kind of atoms, it is called a homoatomic molecule.
The molecule that consists of atoms of different elements is called a heteroatomic molecule. A heteroatomic molecule is known as a compound.

The number of atoms present in the molecule is called its ‘atomicity’.

Atomicity	Number of atoms present	Name
\(1\)	\(1\)	Monoatomic
\(2\)	\(2\)	Diatomic
\(3\)	\(3\)	Triatomic
\(>3\)	\(>3\)	Polyatomic

PYQ - Atomicity

Relative molecular mass (RMM):

The relative molecular mass of a molecule is the ratio between the mass of one molecule of the substance to

\frac{1}{12^{th}}

mass of an atom of carbon-\(12\).

Relative molecular mass is only a ratio. Hence, it has no unit. If the molecular mass of a compound is expressed in grams, it is called gram molecular mass.

Example:

Gram Molecular Mass of \(H_2O\) \(=\) \(18\) \(g\)
Gram Molecular Mass of \(CO_2\) \(=\) \(44\) \(g\)
Gram Molecular Mass of \(NH_3\) \(=\) \(17\) \(g\)
Gram Molecular Mass of \(HCl\) \(=\) \(36.5\) \(g\)

PYQ - Gram molecular mass

Calculating relative molecular mass of sulphuric acid (\(H_2SO_4\)):

Sulphuric acid contains, hydrogen (\(H\)) \(=2\) atoms, sulphur (\(S\)) \(=1\) atoms and oxygen (\(O\)) \(=4\) atoms.

Relative molecular mass of sulphuric acid,

\begin{array}{l} = (2 \times Mass of hydrogen) + (1 \times Mass of sulphur) + (4 \times Mass of oxygen) \\ = (2 \times 1) + (1 \times 32) + (4 \times 16) \\ = 98 \end{array}

i.e., one molecule of \(H_2SO_4\) is \(98\) times as heavy as

\frac{1}{12^{th}}

of the mass of a carbon – \(12\).

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1. Introduction to atoms and molecules

Theory: