Class 9 Ch- 3 Atoms and Molecules

 

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Ch- 3 Atoms and Molecules

Atoms 

Ancient Greek philosophers - Democritus and Lucius, said that when matter is divided and divided, after some time particles are found which cannot be further divided. The democrats gave the name to this particle an atom.

According to English scientist John Dalton’s atomic theory, all matter, whether an element, or a compound or a mixture is composed of small particles called atoms. 

Atoms are the smallest indivisible particles of an element that take part in chemical reaction.

The postulates of this theory may be stated as follows:

(i) All matter is made of very tiny particles called atoms.

(ii) Atoms are indivisible particles, which cannot be created or destroyed in a chemical reaction.

(iii) Atoms of a given element are identical in mass and chemical properties.

(iv) Atoms of different elements have different masses and chemical properties.

(v) Atoms combine in the ratio of small whole numbers to form compounds.

(vi) The relative number and kinds of atoms are constant in a given compound. Atoms are composed of fundamental or subatomic particles which are protons, neutrons and electrons.

 About atoms:

• Atoms are composed of fundamental or subatomic particles which are protons, neutrons and electrons.
• Protons are positively charged particles, electrons are negatively charged particles and neutrons are neutral in charge.
• Central part of the atom called nucleus contains both protons and neutrons. Electrons are present around the nucleus. 
• Atoms are very small in size. Atomic radius is measured in nano-metres.

1/109 m = 1 nm

1 m = 109 nm. 

• Atoms normally exist in a combined form which gives us various forms of matter.

Two important laws of chemical combination given by Antoine L. Lavoisier and Joseph L. Proust as follows:

1.Law of constant proportion - This law was stated by Proust as “In a chemical substance the elements are always present in definite proportions by mass”.

Example- Water is a compound in which the ratio of the mass of hydrogen to the mass of oxygen is always 1:8.

2H₂O = 2H₂ + O₂ 

2H₂ = 2 (2 x 1) = 4  ( atomic mass of hydrogen = 1)

O₂ =16 x 2 = 32 (atomic mass of oxygen = 16)

Ratio of hydrogen and oxygen in 2 molecules of water

4 : 32 or   2 : 16 or  1: 8

1.Law of conservation of mass- Mass can neither be created nor destroyed in a chemical reaction. 

Example:  A + B → C + D

Reactants → Products 

Mass of reactants = Mass of products

N₂ + 3H₂ = 2NH₃

(2 X 14) + 3 x (1 x 2) = 2 ( 14 + 1 x 3 )   [ At. mass of N = 14 and H= 1]

28 + 6 = 2 x 17

34 = 34

Total mass of reactants= Total mass of products

The Modern day symbols of atom of different elements:

Dalton was the first scientist to use these symbols for elements.

Symbols of some elements are given from the first letter of the name and a letter, appearing later in the name, some other symbols have been taken from the names of elements in Latin.

Symbols for some elements:

 

Element	Symbol	Element	Symbol
Aluminium	Al	Iodine	I
Argon	Ar	Iron	Fe
Barium	Ba	Lead	Pb
Boron	B	Magnesium	Mg
Bromine	Br	Neon	Ne
Calcium	Ca	Nitrogen	N
Carbon	C	Oxygen	O
Chlorine	Cl	Potassium	K
Cobalt	Co	Silicone	Si
Copper	Cu	Silver	Ag
Fluorine	F	Sodium	Na
Gold	Au	Sulpher	S
Hydrogen	H	Uranium	U

 

Atomic number – The number of protons present in the nucleus of an atom is called its atomic number.

In an atom Number of protons = Number of electrons.

Therefore,

Atomic number = Number of protons = Number of electrons

 Atomic mass - Atomic mass of an element is the total mass of one atom of given element.

 Relative Atomic mass: The relative atomic mass of the atom of an element is the average masses of the atom, as compared to 1/12th the mass of one carbon-12 atom.

 Molecules- An atom of most elements does not exist independently, they are mostly found in the form of molecules and ions and these molecules or ions combine to form the matter or compound.

Molecules are a group of two or more atoms that are chemically bonded together by the attractive forces.

Definition - Molecule is the smallest particle of an element or a compound that is capable of an independent existence and shows all the properties of that substance.

Example: 

H₂ is a 1 molecule of element hydrogen in which two atoms of hydrogen combine together by the attractive forces.

H or 2H are not molecules; they are one or two atoms of hydrogen.

Similarly H₂O is a 1 molecule of compound of water.

O₂ is a molecule of oxygen consists of two atoms of oxygen and it

is known as a diatomic molecule, O₂. 

O₃ is a molecule of 3 atoms of oxygen which is known as ozone. 

• A molecule is formed when two or more atoms bonded chemically.
• A molecule of a compound has two or more elements.

 Atomicity 

The number of atoms constituting a molecule is known as its atomicity.

• Atomicity of O₂ is 2 and is called diatomic element, 
• Atomicity of O₃ is 3, and atomicity of  H₂O is 3 (2 atoms of hydrogen and 1 atom of oxygen) etc. and are called tri-atomic elements.
• 1 molecule of P₄ phosphorus contains 4 atoms and is called tetra-atomic.
• Sulfur (S₈) contains more than 4 atoms are called poly-atomic. 

 Compound 

Compound is a substance that has two or more elements of different atoms in a fixed ratio.

Compounds are formed by the combination of two or more elements.

Example: CaCO₃ is a compound of 1 atom of calcium, 1 atom of carbon and 3 atoms of oxygen.

• A compound is formed when different elements bonded chemically.
• All compounds are molecules but all molecules are not compounds.

 Ion

Compounds are composed of metals and non- metals in which metals are positively charged particles and non- metals are negatively charged particles. These charged particles are called ions.

Example: Common salt is a compound of sodium and chlorine, where sodium (metal) is a positively charged particle and chlorine (non- metal) a negatively charged particle. Sodium and chlorine are the ions. They are represented by Na⁺ and Cl^-. 

 Anion & Cation - A negatively charged ion is called an ‘anion’ and the positively charged ion, a ‘cation’.

Example: Na⁺ is an anion and is a Cl^- cation.

Poly-atomic ion - A group of atoms carrying a charge is known as a poly-atomic ion.

Example: Ammonium- NH₄⁺, Nitrate- NO₃^-, Sulphate- SO₄^2-, Hydrogen carbonate- HCO₃^-. 

 Chemical formula

Chemical formula of a compound is a symbolic representation of the number of atoms and type of atoms of molecules.

Example: Chemical formula of compound sodium hydroxide is NaOH. This formula represents that sodium hydroxide composed of one atom of sodium, one atom of oxygen and one atom of hydrogen.

 Valency

The combining capacity of an element with other is known as its valency. 

Valency of an element depends on the number of electrons present in the outermost shell of their atoms.

In metals, number of electrons in outermost shell is less than 4 and valency is equal to the number of electrons in the outermost shell of their atoms.

Example: Valency of Na

Atomic number of Na = 11

Electronic configuration  = 2, 8, 1  No. of electron in outermost shell is 1.

Therefore valency = 1

In non-metals number of electrons is more than 4 and valency is equal to eight minus the number of electrons in the outermost shell.

Example: Valency of Chlorine Cl 

Atomic number of chlorine = 17

Electronic configuration = 2, 8, 7

Valency of Cl = 8 - 7 = 1

The rules for the writing a chemical formula:

• The valencies or charges on the ion must balance.

• When a compound consists of a metal and a non-metal, the symbol of the metal is written first.
• The poly-atomic ion in the compound should be enclosed in the bracket.

Formula of simple compound

1.    Formula of Calcium carbonate 

                Ca            CO₃

Valency   +2              -2

       ( Exchanges of valency )    

              (Ca)₂    (CO₃)₂   

   Molecular formula of calcium carbonate is CaCO₃

2.              Formula of Water

    H                O

    1                -2

      (Exchanges of valency)   

                 H₂   O₁

    Molecular formula of water is H₂O   

The 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).

Example: 

Calculate the relative molecular mass of the water (H₂O).

Solution: Water contains the elements of hydrogen and oxygen.

Atomic mass of hydrogen = 1u,

Atomic mass of oxygen = 16 u

Therefore the molecular mass of water = 1 x 2 + 16 = 2 + 16 = 18 u.

To calculate the molecular mass of CaCO₃

Calcium carbonate contains calcium Ca, Carbon C and oxygen O.

Atomic mass of Ca = 20 u

Atomic mass of C = 12 u

Atomic mass of O = 16 u

Therefore the molecular mass of CaCO₃ = 40 + 12 + 16 x 3 = 40 + 12 + 48= 100 

Formula Unit Mass

 The formula unit mass of a substance that contains ions is the sum of the atomic masses of all atoms in a formula unit of a compound.

Example: Ions contained substance NaOH

Formula unit mass of NaOH

Atomic mass of Na = 23 u

Atomic mass of O = 16 u

Atomic mass of H = 1 u

Formula unit mass = 23 + 16 + 1 = 40 u

Mole concept 

 The mole concept is a method of expressing the amount of a substance and the mole is the amount of a substance that contains 6.022 x 10²³ atoms or ions or molecules.

1 mole of any substance = 6.022 x 10²³.

6.022 x 10²³ number is called Avogadro Constant.

Therefore the number of particles or atoms, or molecules or ions present in 1 mole of any substance is 6.022 × 10²³. This number is called the Avogadro Constant or Avogadro Number.

Avogadro's number represented by N₀.

It was named in honour of the Italian scientist Amedeo Avogadro.

The mass of 1 mole of a substance is equal to its relative atomic or molecular mass in grams.

1 mole = 6.022 × 10²³ number = Relative mass in grams.

Example: 

The mass of 1 mole of NaOH = Molecular mass of NaOH = 23 + 16 + 1= 40 g.

[Molecular mass of Na = 23g, O = 16g , H = 1g ]

 The mass of 1 mole of NaOH = 40g

Mathematical expressions to find the mole:

Given mass = m

Molar mass = M

Number of moles = n

Given number of particles = N

Avogadro's number = N₀

Then 

I.  Number of moles = Given mass in g / molar mass

 Example: Calculate the number of moles of 69 g of Sodium.

Solution: 

Molar mass of sodium Na (m) = 23 g

Given mass of sodium Na (M) = 69 g

Therefore 

Number of moles = Given mass in g / molar mass = 69 / 23 = 3

    Number of moles = 3 moles

 II. Number of moles = Given number of particles/Avogadro's number

Example: Calculate the number of mole of 12.044 × 10²³ particles of Mg.

Solution:

Given 

Number of particles = 12.044 × 10²³

Avogadro's number (1 mole)   = 6.022 × 10²³

Number of moles = N / N₀= 12.044 × 10²³ / 6.022 × 10²³

Number of moles = 2 moles.

 

III. Calculate the mass of the followings:

(i) 0.5 mole of O₂ gas 

(ii) 0.5 mole of O atoms 

(iii) 3.011 × 10²³ number of O atoms

(iv) 6.022 × 10²³ number of O₂ molecules.

 

Solution: 

(i) We have 

1 Mole = Molar mass

Molar mass of O₂= 2 X 16 = 32 g

Mass = Number of moles X Molar mass

Mass = n X M = 0.5 X 32 = 16 g.

 

(ii) 1 Mole = Molar mass

Molar mass of O= 16 g

Mass = Number of moles X Molar mass

Mass = n X M = 0.5 X 16 = 8 g.

 

(iii) Number of moles = Given number of particles/Avogadro's number

And    N / N₀= 3.011 × 10²³ / 6.022 × 10²³   = 1/2

Number of moles = 1/2  mole = 0.5 moles.

Molar mass of O = 16 g

Now mass = n X M = 0.5 X 16 = 8 g.

 

(iv) Number of moles = Given number of particles/Avogadro's number

     N / N₀= 3.011 × 10²³ / 6.022 × 10²³   = 1/2

Number of moles = 1/2 mole = 0.5 moles.

Molar mass of O₂ = 16 X 2 = 32 g

Now mass = n X M = 0.5 X 32 = 16 g.

 

IV. Calculate the number of particles in each of the following:

(i) 80 g of Ca atoms 

(ii) 7 g N₂ molecules 

(iii) 0.1 moles of carbon atoms 

Solution:

(i) Given mass of calcium Ca = 80 g

Molar mass of Ca = 40 g

Avogadro's number = 6.022 × 10²³

Number of particles 

N= Avogadro's number X Given mass / Molar mass = (N₀ X m) / M 

Number of particles N  =  (6.022 × 10²³ x 80 g) / 40

Number of particles N = 6.022 × 10²³ X 2 = 12.044 X 10²³

(ii) Given mass of calcium Ca = 80 g

Molar mass of N = 7 g

Avogadro's number = 6.022 × 10²³

Number of particles 

N= Avogadro's number X Given mass / Molar mass = (N₀ X m) / M

Number of particles N  = (6.022 × 10²³ X 7) / 14 

Number of particles N = (6.022 × 10²³) /2 = 3.011 X 10²³

(iii)  Given number of mole of carbon (C)  n = 0.1 mole

Avogadro's number= 6.022 × 10²³

Number of particles N = Given number of moles × Avogadro’s number =  n × N₀

Number of particles N  = 0.1 × 6.022 × 10²³

Number of particles N = 0.6022 × 10²³.