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CBSE NOTES CLASS 9 SCIENCE CHATER 3

ATOMS AND MOLECULES

Law of conservation of mass:

Law of constant proportions

Dalton’s Atomic Theory

Atom

Names of elements

Symbols of Elements

Atomic Mass Unit (amu or u)

Atomic mass of an element

Molecule:

Atomicity

Molecules of Compounds

Ion

Valency

Formulae of simple compounds

Gram atomic mass

Molecular mass

Formula unit mass

Molecular Mass and Mole Concept

Avogadro’s constant or Avogadro’s number

Mole

Formulae for mole concept

CBSE NOTES CLASS 9 SCIENCE CHATER 3

ATOMS AND MOLECULES

1. Law of conservation of mass:

This law was stated by Lavoisier in 1789. It states that

“In all physical and chemical changes, the total mass of reactants is equal to the total mass of products.”

Or

Mass can neither be created nor be destroyed in a chemical reaction.

2. Law of constant proportions (or constant composition):

This law was first stated by Proust in 1797. According to the law

“In a chemical substance the elements are always present in definite proportions by mass”.

For example the ratio of hydrogen and oxygen in pure water is always 1:8 by mass.

This law is also called law of definite proportions.

3. Dalton’s Atomic Theory : Postulates of Dalton Atomic Theory

(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.

4. Atom: It is the smallest particle of an element which can take part in a chemical change. It may or may not be capable of independent existence. It cannot be seen by naked eye.

Atomic radius is measured in anometers.

1/109 m =10-9 m = 1 nm & 1 m = 109 nm

Relative Sizes

Radii (in m)

Example

10–10

Atom of hydrogen

10–9

Molecule of water

10–8

Molecule of haemoglobin

5. Names of elements: Names of elements in the past were derived from the name of the place where they were found for the first time. For example, the name copper was taken from Cyprus.

Some names were taken from specific colours. For example, the name of gold was taken from the English word meaning yellow.

Sometimes the names are derived from the names of their discoverers, e.g., Rutherfordium.

Now-a-days, IUPAC (International Union of Pure and Applied Chemistry) approves names of elements.

6. Symbols of Elements: The abbreviations used for lengthy names of elements are termed as their symbols. The symbol of an element is the first letter or the first and another letter from English name or Latin name of the element. While writing a symbol, the first letter is always capital and the second is always small.

For example

(i) Hydrogen, H

(ii) Aluminum, Al and not AL

(iii) Cobalt, Co and not CO

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

Silicon

Si

Copper

Cu

Silver

Ag

Chromium

Cr

Sodium

Na

Fluorine

F

Sulphur

S

Gold

Au

Uranium

U

Hydrogen

H

Zinc

Zn

7. Atomic Mass Unit (amu or u) - One atomic mass unit is a mass equal to exactly one-twelfth (112th) the mass of one atom of carbon-12.

8. Atomic mass of an element: The atomic mass of an element is the number which indicates how many times an atom of an element is heavier than 112th of mass of an atom of carbon-12.

Atomic masses of a few elements

Name of Element

Symbol

Atomic Number

Atomic mass

Hydrogen

H

1

1

Helium

He

2

4

Lithium

Li

3

7

Beryllium

Be

4

9

Boron

B

5

11

Carbon

C

6

12

Nitrogen

N

7

14

Oxygen

O

8

16

Fluorine

F

9

19

Neon

Ne

10

20

Sodium

Na

11

23

Magnesium

Mg

12

24

Aluminium

Al

13

27

Silicon

Si

14

28

Phosphorus

P

15

31

Sulphur

S

16

32

Chlorine

Cl

17

35.5

Argon

Ar

18

40

Potassium

K

19

39

Calcium

Ca

20

40

8. Molecule: It is the smallest particle of an element or compound that is capable of independent existence and shows all the properties of that substance.

[The molecule of an element is made up of same type of atoms, while the molecule of a compound is made up of different types of atoms]

9. Atomicity: The number of atoms present in a molecule of an element or a compound is known as its atomicity, e.g. the atomicity of oxygen is 2 while atomicity ozone is 3.

Atomicity of some elements

Carbon dinoxide

Type of Element

Name

Atomicity

Non-Metal

Argon

Monoatomic

Helium

Monoatomic

Oxygen

Diatomic

Hydrogen

Diatomic

Nitrogen

Diatomic

Chlorine

Diatomic

Ozone

Triatomic

Phosphorus

Tetra-atomic

Sulphur

Poly-atomic

Metals

Sodium

Monoatomic

Iron

Monoatomic

Aluminium

Monoatomic

Copper

Monoatomic

Compounds

Hydrochloric Acid

Diatomic

Carbon monoxide

Diatomic

Triatomic

Sulphuric Acid

Poly-atomic

10. Molecules of Compounds: Atoms of different elements join together in definite proportions to form molecules of compounds.For example we can tabulate the diffrent atoms in water as follows,

Element

Ratio by mass

Atomic mass (u)

Mass ratio /atomic mass

Simplest ratio

H

1

1

11

2

O

8

16

816=12

1

11. Ion: It is an electrically charged atom or group of atoms. It is formed by the loss or gain of electrons by an atom. Ions are of two types :

(i) Cation: It is positively charged ion and is formed by the loss of electron from an atom e.g. H+, Na+, Ca2+, Al3+, NH4+ etc.

(ii) Anion: It is negatively charged ion and is formed by the gain of electrons by an atom, e.g. Cl-, O2-, F-, CO32-, PO43- etc.

A compound made by combination of ions is called ionic compound.

Some common, simple and polyatomic ions

Val

Metalic ion

Symbol

Non-metallic ions

Symbol

Polyatomic ions

Symbol

1.

Sodium

Na+

Hydrogen

H+

Ammonium

NH4+

Potassium

K+

Hydride

H-

Hydroxide

OH

Silver

Ag+

Chloride

Cl-

Nitrate

NO3

Copper (I)*

Cu+

Bromide

Br-

Hydrogen carbonate

HCO3

Iodide

I

Nitrite

NO2-

2.

Magnesium

Mg2+

Oxide

O2-

Carbonate

CO32–

Calcium

Ca2+

Sulphide

S2-

Sulphite

SO32–

Zinc

Zn2+

Sulphate

SO42–

Iron(II)*

Fe2+

Copper(II)*

Cu2+

3.

Aluminium

Al3+

Nitride

N3-

Phosphate

PO43–

Iron(III)*

Fe3+

13. Valency: The combining power (or capacity) of an element or ion is known as its valency.

14. Formulae of simple compounds: Binary compounds are those compounds which are made up of two different elements e.g. NaCl, KBr, and CaO etc.

Following rules are to be followed for writing the formula.

(i) The valencies or charges on the ions must be balanced.

(ii) For a compound made up of a metal and a non-metal, the symbol of metal is written first.

(iii) In compounds formed with polyatomic ions, the ion is enclosed in a bracket before writing the number to indicate the ratio.

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15. Gram atomic mass: The atomic mass of an element expressed in grams is known as gram atomic mass. (Gram atomic mass is also known as gram atomic weight).

16. Molecular mass: The number of times a molecule of a compound is heavier than 112th of the mass of C-12 atom, is known as its molecular mass.

The molecular mass is equal to the sum of the atomic masses of all atoms present in one molecule of the substance. For example H2S contains two atoms of hydrogen and one atom of S, so molecular mass of H2S is 2×1 + 32 = 34.

Gram Molecular mass is the molecular mass expressed in grams.

17. Formula unit mass: It is equal to the sum of atomic masses of all the atoms in a formula unit of ionic compounds. Formula unit mass of NaCl is 23 + 35.5 = 58.5.

18. Mole Concept: One mole of any species (atoms, molecules, ions or particles) is that quantity in number having a mass equal to its atomic or molecular mass in grams. Irrespective of the substance under consideration 1 mole is equal to 6.022 × 1023 a specie.

The mass of 1 mole of particles (or simply is equal to its mass in grams.

1 mole = 6.023 × 1023 particles

1 mole atoms = 6.023 × 1023 atoms

1 mole electrons = 6.023 × 1023 electrons

1 mole protons = 6.023 × 1023 protons

1 mole ions = 6.023 ×1023 ions

1 mole molecules = 6.023 × 1023 molecules

19. Avogadro’s constant or Avogadro’s number: The number of particles present in one mole (i.e. 6.023 × 1023) is called Avogadro’s number or Avogadro’s constant.

20. Mole

1 mole = 6.023 × 1023 particles.

= mass of 1 mole particles in grams

= 22.4L of a gas at N.t.p.

1 mole atoms = gram atomic mass

1 mole molecules = gram molecular mass

N.t.p. stands for normal temperature (0°C) and normal pressure (1 atmosphere or 760 mm of Hg)

21. Formulae for mole concept :

Number of moles = Given mass in gramsGram molecular mass 

Number of moles =  Volume of gas in litres at STP22.4

Number of moles = Number of particles in the given amountAvogadro's number 

Number of atoms (or molecules or ions)

= given massmolar mass × Avogadro number

Given mass = molar mass × number of moles