Well structured WBBSE Class 9 Physical Science MCQ Questions Chapter 5 Work, Power and Energy can serve as a valuable review tool before exams.

Work, Power and Energy Class 9 WBBSE MCQ Questions

Multiple Choice Questions :

Question 1.
What type of energy a raised hammer possesses ?
(i) Kinetic energy
(ii) Potential energy
(iii) Gravitational energy
Answer:
Potential energy

Question 2.
What is the absolute unit of force in SI system ?
(i) 1 kg-m
(ii) 1 joule
(iii) 1 erg
Answer:
1 Joule

Question 3.
What is the relation between power and force ?
(i) power = force × velocity
(ii) velocity = force × power
(iii) force = power × velocity
Answer:
power = force × velocity.

Question 4.
What is the amount of work performed by a person on the weight w that he carries and moves horizontally through a distance d ?
(i) No work is done
(ii) Work done = w × d
(iii) work done = w ÷ d
Answer:
No work is done

Question 5.
The dimension of energy is :
(i) [M L² T^-2]
(ii) [M² L² T³]
(iii) [M L² T³]
(iv) [M L³ T²]
Answer:
ML² T-²

Question 6.
1 Horse power = ______.
(i) 7.46 watt
(ii) 74.6 watt
(iii) 746 watt
(iv) 7460 watt
Answer:
746 watt

Question 7.
All energies are transformed ultimately into :
(i) heat energy
(ii) kinetic energy
(iii) sound energy
(iv) electric energy
Answer:
Heat energy

Question 8.
Which is a scalar quantity :
(i) Force
(ii) Displacement
(iii) Work
(iv) Velocity
Answer:
Work

Question 9.
If the velocity of a body is doubled then kinetic energy will be :
(i) Same
(ii) Doubled
(iii) Tour times
(iv) Half times
Answer:
Four times

Question 10.
Work done in unit time is known as :
(i) Power
(ii) Energy
(iii) Force
Answer:
Power

Question 11.
1 kilowatt = ________.
(i) \(\frac{2}{3}\) HP
(ii) \(\frac{4}{3}\) HP
(iii) \(\frac{7}{3}\) HP
Answer:
\(\frac{4}{3}\) HP

Question 12.
What is the unit of kinetic energy in SI system :
(i) erg
(ii) HP
(iii) joule
Answer:
joule

Question 13.
1 mega-watt = _________.
(i) 10³ watt
(ii) 10⁴ watt
(iii) 10⁵ watt
(iv) 10⁶ watt
Answer:
10⁵ watt

Question 14.
What is the practical unit of power in CGS system :
(i) erg
(ii) watt
(iii) joule
Answer:
watt

Question 15.
Kilogram-meter is the unit of :
(i) energy
(ii) power
(iii) work
Answer:
Work

Question 16.
The dimension of work is :
(i) [M L² T^-2]
(ii) [M² L² T²]
(iii) [M L³ T^-2]
Answer:
[ML² T^-2]

Question 17.
The dimension of power is :
(i) [M L² T^-3]
(ii) [M² L² T²]
(iii) [M L³ T³]
Answer:
[ML² T^-3]

Question 18.
1 watt = ______.
(i) 10⁵ erg/sec
(ii) 10⁶ erg/sec
(iii) 10⁷ erg/sec
Answer:
10⁷ erg/sec.

Question 19.
SI unit of energy is :
(i) Joule
(ii) erg
(iii) dyne
(iv) newton
Answer:
Joule

Question 20.
The potential energy of a boy is maximum when he is :
(i) sitting on the floor
(ii) standing on the floor
(iii) sleeping on the floor
Answer:
Standing on the floor.

Question 21.
Water stored in a dam possesses :
(i) KE
(ii) PE
(iii) elecric energy
(iv) no energy
Answer:
Electric energy

Question 22.
The kinetic energy of an object is K. If its mass is reduced to half, them its kinetic energy will be :
(i) K
(ii) 2 K
(iii) \(\frac{K}{2}\)
Answer:
\(\frac{K}{2}\)

Question 23.
The work done on an object does not depend upon the :
(i) force applied
(ii) initial velocity
(iii) displacement
Answer:
Initial velocity

Question 24.
In case of negative work, the angle between the force and displacement is :
(i) 0°
(ii) 45°
(iii) 180°
Answer:
180°

Question 25.
In a tug of war, work done by a losing team is :
(i) Zero
(ii) Positive
(iii) negative
Answer:
Negative

Question 26.
A stone is thrown vertically upward. It comes to rest momentarily at the highest point. What happens to its kinetic energy ?
(i) It convents into elastic potential energy
(ii) It convents into gravitational potential energy
(iii) It converts into chemical energy.
Answer:
It convents into gravitational potential energy.

Question 27.
A force of 100 N acting on a body does 100 j work. The distance through which the body is displaced –
(i) 5 m
(ii) 10 cm
(iii) 10 m
Answer:
10 m

Question 28.
When electric current passes through an electric bulb, electric energy is converted into :
(i) heat energy only
(ii) light energy only
(iii) both heat and light energy
Answer:
Both heat and light energy.

Fill in the blanks :

1. The gravitational unit of work in the _____ system is newton-meter.
Answer:
SI

2. Power is the ratio between work and _____
Answer:
time

3. Just before touching the ground, a freely falling body possesses only ______ energy.
Answer:
kinetic energy

4. Watt is the unit of ______ in SI system.
Answer:
power.

5. A raised hammer possesses ______ energy.
Answer:
potential

6. Water current or air current mainly possesses ______ energy.
Answer:
kinetic energy

7. 1 horse power = ______ watt.
Answer:
746

8. Both power and energy are ______ quantities.
Answer:
scalar

9. 1 joule = ________ erg.
Answer:
10⁷

10. Potential energy of body is taken as ________
Answer:
mgh

11. 1 kilowatt = _______ Horse power.
Answer:
134

12. Power × time = ________
Answer:
work done

13. Work is a _______ quantity.
Answer:
scalar

14. newton × meter = _______.
Answer:
joule

15. The capacity of a body or a person to do work is its ________
Answer:
energy

16. kinetic energy = _______ × m × v²
Answer:
\(\frac{1}{2}\)

17. When a force acts _______ to the direction of displacement of the particle, the force does not do any work.
Answer:
perpendicular

18. If there is no _______ of the point of application of the force, then the force is called no work force.
Answer:
displacement

19. If the direction of displacement (S) makes an angle θ with the direction of the applied force (F), component of the force along the direction of displacement will be ________
Answer:
FS\cos\θ

20. The ______ work is done, when the displacement is in the direction of the force.
Answer:
positive

21. In SI, the absolute unit of work is ______.
Answer:
joule

22. work done per unit time is called _______.
Answer:
power

23. One horse power is equal to _______ watt.
Answer:
746

24. The capacity of a body to do work is ______.
Answer:
energy

25. A stretched spring has potential energy due to its ______.
Answer:
shape.

26. The gravitational potential energy of body when raised from ground to certain height is _______ of the path followed.
Answer:
independent

27. Larger the mass of a body ______ is the kinetic energy of the body.
Answer:
greater

28. Total energy of a system is ______ , when energy is changed from one form to other.
Answer:
conserved

29. The mechanical energy is of _______ types.
Answer:
two

Detailed explanations in West Bengal Board Class 9 Physical Science Book Solutions Chapter 4.1 Atomic Structure offer valuable context and analysis.

WBBSE Class 9 Physical Science Chapter 4.1 Question Answer – Atomic Structure

Very short answer type questions

Question 1.
Which atom has no neutron?
Answer:
Ordinary hydrogen atom has no neutron.

Question 2.
Which force binds the nucleons?
Answer:
The nuclear force, a short ranged attractive force binds the nucleons.

Question 3.
Which fundamental particle is responsible to produce isotopes?
Answer:
Neutron is responsible to produce isotopes, variation of it causes variation of mass number.

Question 4.
Name the element of the same mass number and atomic number?
Answer:
Ordinary hydrogen has the same mass number and atomic number, each being 1.

Question 5.
What are nucleons ?
Answer:
The constituents of the nucleous of an atom, proton and neutron are called nucleons.

Question 6.
Name the heaviest isotope of hydrogen.
Answer:
Tritium is the heaviest isotope of hydrogen.

Question 7.
State one difference between H and H⁺.
Answer:
H is an unstable hydrogen atom having no net electric charge. H⁺ is a positive hydrogen ion carrying unit positive charge.

Question 8.
What is deuteron and what is triton ?
Answer:
The nucleous of deuterium is called deuteron and that of tritium is called triton.

Question 9.
What is isotopic weight ?
Answer: The weight of an isotope of an element is known as isotopic weight.

Question 10.
Can there be a doubly ionised hydrogen ion ?
Answer:
Hydrogen has only one electron, so it can be ionised singly by losing the electron, so there cannot be a doubly ionised hydrogen ion.

Question 11.
What is the relation between atomic number and mass number ?
Answer:
The relation is A = Z + N, where A = mass number, Z = atomic number, N = number of neutrons.

Question 12.
Name two particies of an atom which are always equal in number.
Answer:
Proton and electron.

Question 13.
What is the magnitude of charge of a neutron ?
Answer:
Neutron has no charge.

Question 14.
Which part of an atom contains protons?
Answer:
Nucleus.

Question 15.
Why an atom always remains neutral inspite of the existence of electrons and protons which are charged particles ?
Answer:
As electrons and protons of an atom are always equal in number, the positive charges of the protons and negative chargy of the electrons cancell each other. As a result, an atom always remains neutral.

Question 16.
Who discovered neutron?
Answer:
J. Chadwick (1932) discovered neutron.

Question 17.
Who discovered nucleus of the atom ?
Answer:
Rutherford discovered nucleus of the atom.

Question 18.
What are electron-shells ?
Answer:
The electrons outside the nucleus of the atom rotate in definite number in certain specified circular paths around the nucleus. These con-centric circular paths of the rotating electrons are called electronshells.

Question 19.
Write the electronic configuration of chlorine.
Answer:
The electronic configuration of chlorine is :
K L M
2 8 7

Question 20.
What are ions ?
Answer:
Electrically charged atoms (or, group of atoms) are called ions.

Question 21.
How are ions formed from neutral atom ?
Answer:
The formation of ions from neutral atoms involves loss or gain of one or more than one electrons.

Question 22.
What are the fundamental particles that constitute an atom ?
Answer:
Electron, proton and neutron are the fundamental particles that constitute an atom.

Question 23.
Among the different fundamental particles, which one is positively charged and which one is negatively charged ?
Answer:
Proton is positively charged particle and electron is negatively charged particle.

Question 24.
What is cation ?
Answer:
Positively charged atom or radical is callled a cation.

Question 25.
What is anion?
Answer:
Negatively charged atom or radical is called an anion.

Question 26.
Among electron, proton and neutron, which one is the heaviest and which one is the lightest particle?
Answer:
Neutron is the heaviest particle and electron is the lightest particle.

Question 27.
Calculate the number of protons and neutrons in \({ }_{92}^{235} \mathrm{U}\).
Answer:
\({ }_{92}^{235} \mathrm{U}\) → protons : 92
neutrons : (235-92)=143

Question 28.
What is the relation between \({ }_{92}^{235} \mathrm{U}\) and \({ }_{92}^{238} \mathrm{U}\) ?
Answer:
These are two isotopes of uranium.

Question 29.
Give example where mass number and atomic numbers are equal.
Answer:
Mass number and atomic numbers are equal in those elements which do not have any isotope. e.g. Na, F etc.

Question 30.
Between K and K⁺ which one is more stable?
Answer:
K⁺is more stable than K as in the earlier case the electronic configuration is just like inert gas argon.

Question 31.
What is heavy water ?
Answer:
Oxide of deuterium (D₂O) is called heavy water.

Question 32.
What is the maximum capacity of L shell to accommodate electrons ?
Answer:
8 electrons.

Question 33.
What is the mass of an electron ?
Answer:
Mass of an electron is 9.11 × 10-28 g.

Question 34.
What is the charge of an electron ?
Answer:
Charge of an electron is 1.602 × 10^-19 coulomb or 4.8 × 10^-10 esu}

Question 35.
What is amu?
Answer: It is an unit used for measuring the atomic mass of an atom. 1 amu} = 1.66.3 × 10²⁴ g

Question 36.
What is the mass of a neutron ?
Answer:
Mass of a neutron is 1.675 × 10^-24 g.

Question 37.
What is the radius of an electron ?
Answer:
The radius of an electron is 2.8 × 10^-13 cm.

Question 38.
Which is the smaliest partices present in all atoms ?
Answer:
Electrons.

Question 39.
What is cathode ray ?
Answer:
When electric discharge passes from high by potential source through a gas at very low pressure (0.01 mm of Hg) in a glass tube provided with two metallic electrodes, a stream of invisible rays is emitted from the surface of the cathode which moves in straight line towards the anode with a high velocity. These rays are called cathode rays as they originate at the cathode.

Question 40.
Who discovered radio-activity ?
Answer:
Madam curie discovered radio-activity.

Short answer type questions

Question 1.
What are the postulates of Dalton’s atomic theory ?
Answer:
Postulates of Dalton’s atomic theory :
(i) Every element is composed of infinite number of very small indivisible particles. These smallest particles are known as ‘atoms’.
(ii) Atoms cannot be created or destroyed by a chemical reaction.
(iii) Atoms of different elements are different. But atoms of a definite element are same.

Question 2.
What are the fundamental particles? Why are they called ‘fundamental’ ?
Answer:
Fundamental particles: The sub-atomic particleselectron, proton and neutrons are known as fundamental particles.

Reason : Experimentally it is found that these particles are the primary components of all atoms of all elements, except ordinary hydrogen, the nucleus of which does not contain any neutron. That is why they are known as fundamental particles.

Question 3.
State some other sub-atomic particles other than electron, proton and neutron.
Answer:
Other sub-atomic particles are :

1. Positron
2. Antiproton
3. Messon
4. Neutrino
5. Antineutrino
6. V-particle
7. Deuteron etc.

Question 4.
What makes electrons move round the nucleus ?
Answer:
Explanation : The negatively charged electrons in each shell are attracted by Coulombian electrostatic force towards the positive nucleus that contains positively charged protons. To counter balance this inward force the electrons rotate about the nucleus, as a revolving body is always accompanied with an outward force.

Question 5.
What is nuclear force?
Answer:
Nuclear force: In the nucleus of an atom, the protons and neutrons are strongly held together by nuclear forces which result from the attraction between the constituent protons and neutrons. This force is so strong that it is very difficult to separate the nucleons.

Question 6.
Define atomic number. Atomic number is the fundamental property of an element – Explain.
Answer:
Atomic number : The atomic number of an atom of an element is the number of protons in its nucleus.
Atomic number is the fundamental property : Number of protons determine the nature of the element. Change in the number of protons changes its chemical properties. As a result, a new element is formed. So, atomic number of an element is its intrinsic property.

Question 7.
What do you mean by mass number of an element ?
Answer:
Mass number : The mass number of an atom of an element is the total number of protons and neutrons present in the núcleus of the atom.
Mass number of atom = Number of protons + Number of neutrons. Mass number is written on the upper left-hand side of the symbol of the element.

Question 8.
What is the relation between mass number and atomic number ?
Answer:
Relation between mass number and atomic number :
We know that,
Mass number of an atom = Number of protons + Number of neutrons

Suppose, number of protons in the nucleus of an atom = Z, number of neutrons = N, and mass number = A
Then we can write,
A = Z + N
As number of protons = atomic number
∴ Atomic number = mass number – number of neutrons

Question 9.
Define isotope.
Answer:
Isotope : The different atomic species of the same element, which have necessarily the same atomic number but different mass number are called isotopes.
e.g. Hydrogen has three isotopes :
\({ }_1^1 \mathrm{H}\) (Ordinary hydrogen) ; \({ }_1^2 \mathrm{H}\) (Deuterium) ; \({ }_1^3 \mathrm{H}\) (Tritium)

Question 10.
What is meant by \({ }_8^{16} \mathrm{O}\) ? Show its electronic configuration.
Answer:
\({ }_8^{16} \mathrm{O}\) is meant by : \({ }_8^{16} \mathrm{O}\) stands for an isotope of oxygen that has 8 protons and 8 neutrons in the nucleus and 8 electrons surrounding the nucleus.
The electronic configuration is :
K – shell has 2 electrons, L-shell has 6 electrons.

Question 11.
How do the physical and chemical properties differ in the isotopes of an element? Have they any common property?
Answer:
Difference between physical and chemical properties of isotopes of an element : Isotopes of an element have the same chemical property since their atomic number is same i.e. number of valence electrons is same. Physical property of the isotopes of an element slightly differ due to the increase of mass number which results to slight increase of density, melting point and boiling point.
Common property : As isotopes have same atomic number, so they have same chemical properties.

Question 12.
Define atomic weight of an element on the basis of carbon-12.
Answer:
Atomic weight of an element on the basis of carbon-12 : The relative weight of one atom of carbon 12_c is taken as the atomic weight of the element. Here, one-twelvth part of the weight of a carbon atom (C = 12) has been taken as the unit.

Question 13.
Distinguish between atomic number and mass number.
Answer:
Difference between atomic number and mass number:

Atomic Number	Mass Number
(i) It is equal to the number of protons present in the nucleus of an atom.	(i) It is equal to the sum of the numbers of protons and neutrons present in the nucleus of an atom.
(ii) From atomic number, number of valence electrons can be determined which in turn gives the idea of chemical combination of the atom.	(ii) Mass number gives an idea about the atomic mass of the element concerned but it does not give any idea about chemical activity of the element until the number of protons or the number of neutrons is indicated.

Question 14.
How is an atom transformed to an ion? What are cations and anions ?
Answer:
Ions : Electrically charged atoms or radicals are called ions. Cation : Positively charged atom or radical is known as cation. e.g. NH₄ ⁺, Ca²⁺, Na⁺etc.
Anion: Negatively charged atom or radical is called an anion. e.g. SO₄^-2, NO₃^–, Cl^–etc.

Question 15.
What are the differences between an atom and ion ?
Answer:
Differences between an atom and an ion:

Atom	Ion
(i) An atom is electrically neutral.	(i) An ion is electrically charged atom. It is formed when an atom either gains or loses one or more electrons.
(ii) An atom may or may not have free existence. e.g. Na atom reacts immediately when it comes in contact with water. 2 Na + 2 H2O = 2 NaOH + H2↑	(ii) An ion exists freely in solution.

Question 16.
What are the differences between atomic weight and actual weight of an atom ?
Answer:
Distinction between atomic weight and actual weight of an atom :

Atomic weight	Actual weight of an atom
(i) It is a number which represents how many times an atom of an element is heavier than 1/12 weight of carbon atom (C-12)	(i) It is a number which represents how many times an atom is heavier than one atomic mass unit (a.m.u.) which is equivalent to 1.6603 × 10-24 g.
(ii) It is a simple ratio and therefore has no unit.	(ii) It represents the actual weight of an atom and therefore has unit.

Question 17.
Mention two similarities between atomic structure and the structure of the solar system.
Answer:
Similarities between atomic structure and the structure of the solar system :

Atomic Structare	Solar system
(i) Nucleus is at the centre, electrons revolve round it.	(i) The sun is at the centre, planets revolve round it.
(ii) Mass of the nucleus is much greater than that of an electron moving round it.	(ii) Mass of the sun is much greater than that of any planet revolving round it.

Question 18.
Mention three dissimilarities betwoen atomic structure and the structure of the solar system.
Answer:
Dissimilarities between atomic structure and the structure of the Solar system :

Atomic Structure	Solar system
(i) More than one electron may be held in the same orbit.	(i) More than one planet never be held in the same orbit.
(ii) Force between the nucleus and electrons is electrostatic in nature.	(ii) Force between the sun and the planets is gravitational in nature.
(iii) In the atom, the protons and electrons are electrically charged bodies.	(iii) The sun and all other heavenly planets are uncharged.

Broad answer type questions

Question 1.
What conclusions are found from Rutherford’s experiment ?
Answer:
Conclusions from Rutherford’s experiment :
(i) The atom consists of an extremely small, centrally located region called the nucleus in which the whole mass and the entire positive charge carried by the protons remain confined.
(ii) The major portion of the space in an atom is empty.
(iii) Negatively charged electrons present in the atom remain outside the nucleus at relatively large distances and they revolve round the nucleus at some definite circular path known as orbit.

Question 2.
What is the fundamental difference between Rutherford’s model and Bohr’s model?
Answer:
Difference between Rutherford’s model and Bohr’s model : The fundamental difference between the two models is that Bohr’s model is based on the concept of quantisation of energy and angular momentum of electron. Electrons can move only is certain permitted orbits with definite amount of energy and angular momentum. Rutherford’s model does not give an idea about the permitted orbits.

Question 3.
What are the uses of Radio-isotopes ?
Answer:
Uses of Radio-isotopes :
(i) In medicine : Radio isotopes and proved to be very useful in medical diagnosis. Radio-isotope of iodine is used by the patients with thyroid disorder.
(ii) In agriculture : By introducing radioactive phosphorous fertiliser, it has been possible to enhance growth of the plant and crop.
(iii) In Industry : In research and process control radio-isotopes are used as a tracer.
(iv) Radio-carbon dating : With the help of radioactive isotope of carbon \({ }_6^{14} \mathrm{C}\) it has been possible to estimate the age of our earth which is about 4-5 billion years.

Question 4.
What is valence electron ? What is radioactivity ? Give the name of the natural radioactive elements and artificial radioactive elements.
Answer:
Valence Electron : Electron or electrons present in the outermost orbit of the electronic shell of an atom is known as valence electron.

Radioactivity : It is a neuclear phenomenon in which the elements having higher atomic number and mass number (n / p>1.54) disintegrate spontaneously emitting energy from the nucleus in the form of radioactive rays and this emission cannot be stopped by any physical or chemical process which means this emission is independent of external agencies and conditions.
Natural Radioactive element :
Radium (Ra), Thorium (Th), Uranium (U) etc.
Artifical Radioactive element :
Neptunium (Np), Eistenium (Es) etc.

Question 5.
‘Atomic number is the intrinsic property of an element’-explain.
Answer:
Explanation: Atomic number, the number of protons in the nucleus of an atom of an element is unique for an element. It cannot

change even during or after any chemical or physical change that the element undergoes. No two elements have the same atomic number, on the other hand, different elements have different atomic numbers. The atomic number of an element indicates the ability of chemical reactivity of the element. So, atomic number is the intrinsic property of the element.

Question 6.
Atomic weight or relative atomic mass of most of the elements is not a whole number. Why?
Answer:
Explanation : It is found that the atomic weights of most of the naturally occuring elements are fractional although their mass numbers (n+p) are whole numbers. This happens because, most of the natural element consist of a mixture of two or more of their isotopes in various proportions but almost in constant composition.

For example : The atomic weight of chlorine as found by chemical methods is 35.457 . Such fractional atomic weight is due to the fact that chlorine exists in nature as a mixture containing 75.4 % of the ligher \({ }_{17}^{35} \mathrm{Cl}\) and 24.6 % of the heavier \({ }_{17}^{37} \mathrm{Cl}\) isotopes. The composition of the mixture is always found to be constant.
∴ Atomic weight of chlorine
= \(\frac{75.4 × 35+24.6 × 37}{100}\) = 35.457

Numerical Problems

Working formula

(i) 1 gram-molecular weight of a substance contains molecules = 6.022 × 10²³
(ii) 1 gram-atom of an element contains atoms = 6.022 × 10²³
(iii) At NTP, the volume of 1 gram-molecular gas = 22.4 lit.
(iv) At NTP, 22.4 lit. volume of gas has molecules = 6.022 × 10²³
(v) Atomic number = Number of protons
(vi) Mass number = Number of protons + Number of neutrons
(vii) Maximum number of electron in an Bohr’s orbit = 2 n²(n = 1,2,3 … etc.)
(viii) Atomic weight = \(=\frac{\text { Weight of one atom of an element }}{\text { Weight of one atom of carbon }} \times 12\)
(ix) 1 mole = 1 gram-molecular weight of a substance.

Example 1: Calculate the number of moles in 50 g CaCO 3.
Answer:
Formula mass of CaCO₃ = 40+12+3 × 16 = 100
Gram formula mass of CaCO₃ = 100 g
Agin we know, gram formula mass of CaCO₃ = 1 mole ∴ 100 g CaCO₃ = 1 mole
∴ 50 g CaCO₃ = (\(\frac{1}{100}\) × 50) mole = 0.5 mole

Example 2 : Calculate the mass of a silver atom.
Answer:
Atomic mass of silver = 108
∴ Gram-atomic mass of silver = 108 g
Now, 6.022 × 10²³ atoms of silver = 1 mole of silver atoms = 108 g
∴ Mass of 6.022 × 10²³ atoms of silver = 108 g
∴ Mass of 1 atom of silver = \(\frac{108}{6.022 \times 10^{23}}\) = 17.93 × 10^-23 g

Example 3 : Calculate the mass of water in gram of 5 gram-molecular water.
Answer:
We know,
molecular weight of water = 2 × 1+16 = 18
1 gram-molecular mass of water = 18 g water
5 gram-molecular mass of water = (18 × 5) g = 90 g water

Example 4 : What is the volume of oxygen at NTP of 4 gram-molecular weight of Oxygen ?
Answer:
We know, at NTP,
1 gram-molecular weight of oxygen gas has volume 22.4 lit.
∴ 4 gram-molecular weight of oxygen gas has volume = (22.4 × 4) lit. = 89.6 lit.

Example 5 : Calculate the number of hydrogen molecules in 1 grammolecular weight of hydrogen and 1 gram of hydrogen.
Answer:
We know,
1 gram-molecular weight of hydrogen gas has molecule = 6.022 × 10²³.
Molecular weight of hydrogen gas = 2
∴ Gram-molecualr weight of hydrogen gas = 2 g.
∴ 2 g hydrogen has molecule = 6.022 × 10²³
1 g hydrogen has molecule = \(\frac{6.022 \times 10^{23}}{2}\)
= 3.0115 × 10²³

Example 6 : At NTP, 1.12 lit of a gas has weight 2.2 g. Calculate its molecular weight.
Answer:
At NTP 1.12 lit. of a gas weight = 2.2 g
At NTP, 22.4 lit. of a gas has weight = \(\frac{2.2}{1.12}\) × 2.2 .4 g
= 44 g
∴ The molecular weight of the gas = 44

Example 7 : What are the numbers of protons, neutrons, and electrons in \({ }_{28}^{65} \mathrm{Cu}\) atom ?
Answer:
We know from the symbol \({ }_{29}^{65} \mathrm{Cu}\), the atomic number of Cu atom is 29 and the mass number is 65 .
So, the number of protons in Cu atom is 29 ; the number of electrons in Cu atom is also 29.
∴ Number of neutrons = (65-29) = 36

Example 8: What is the mass number and what is the atomic number of \({ }_{92}^{235} \mathrm{X}\) ? What will be its symbol if the number of neutrons are three more?
Answer:
In \({ }_{92}^{235} \mathrm{X}\), the number of protons are 92 and the sum total of protons and neutrons are 235 .
So, the atomic number is 92 and mass number is 235 .
If the number of neutrons are three more, then the mass number will be (235+3) = 238.
In that case the symbol will be \({ }_{92}^{238} \mathrm{X}\).

Example 9: What is the relation between \({ }_{35}^{17} \mathrm{A}\), and \({ }_{37}^{17} \mathrm{B}\) ? Is there any difference in their chemical properties?
Answer:
\({ }_{35}^{17} \mathrm{A}\), and \({ }_{37}^{17} \mathrm{B}\), both have same atomic number but different mass numbers. So they are isotopes of each other. As the chemical property depends on number of protons, so they have same chemical properties.

Detailed explanations in West Bengal Board Class 9 Physical Science Book Solutions Chapter 4.4 Acids, Bases and Salts offer valuable context and analysis.

WBBSE Class 9 Physical Science Chapter 4.4 Question Answer – Acids, Bases and Salts

Very short answer type question

Question 1.
Which element must an acid contain?
Answer:
Hydrogen

Question 2.
The aqueous solution of HCl shows acidic character. Which ion is responsible for it?
Answer:
Hydronium ion (H₃O⁺)

Question 3.
Give the name and formula of anion present in the aqueous solution of bases.
Answer:
Flydroxyl ion (OH^–)

Question 4.
Which indicator is used in the titration of strong acid and weak base?
Answer:
Methyl orange

Question 5.
Name an acid salt.
Answer:
Sodium bisuiphate (NaHSO₄)

Question 6.
Which ions disappear in neutralization?
Answer:
W ions generated by acid and OH^– ions generated by alkali disappear in a neutralization process.

Question 7.
What is an oxide?
Answer:
An oxide is a compound of oxygen formed with another element.

Question 8.
Name a tribasic acid.
Answer:
Phosphoric acid (HPO₄)

Question 9.
Give an example of an organic acid.
Answer:
Formic acid (HCOOH)

Question 10.
Give an example of a normal salt.
Answer:
Sodium chloride (NaCl)

Question 11.
What is the nature of aqueous solution of carbon dioxide?
Answer:
Acidic

Question 12.
What type of salt is this — NaHCO₃?
Answer:
Acidic.

Question 13.
What is the use of methyl orange?
Answer:
As an indicator

Question 14.
What is neutralisation?
Answer:
Neutralisation: It is the process in which acids and alkalis in equivalent quantities in their aqueous solutions react to produce salt and the neutral substance water.

Question 15.
Give an example of a neutral oxide.
Answer:
Carbon monoxide (CO)

Question 16.
What will be the colour of the solution if few drops of phenolphthalein are added in aqueous solution of ammonium hydroxide?
Answer:
Pink colour

Question 17.
What will be the colour of the solution if few drops of phenolphthalein are added in aqueous solution of sodium carbonate?
Answer:
Pink colour

Question 18.
Which Indicator is used in the titration of weak acid and strong base?
Answer:
Phenolphthalein.

Question 19.
What is the nature of ZnO?
Answer:
Amphoteric oxide.

Question 20.
Give an example of an acid which has oxidising property.
Answer:
Nitric acid (HNO₃)

Question 21.
What will be the colour of the solution if few drops of phenolphthalein are added in aqueous solution of sodium hydrogen carbonate?
Answer:
Colourless

Question 22.
Which one is used in vanishing colour— NH₄OH or NaOH?
Answer:
NH₄OH

Question 23.
What is the basicity of CH₃COOH?
Answer:
Basicity of CHCOOH is 1.

Question 24.
What is the acidity of NaOH?
Answer:
Acidity of NaOH is 1.

Question 25.
Give an example of an oxide of metal which is not basic in nature.
Answer:
Mn₃O₄.

Question 26.
Give an example of a double salt.
Answer:
K₂SO₄ ,Al₂(SO₄)₃, 24H₂O (Potassium alum)

Question 27.
How many replaceable hydrogen atoms are present in sulphuric acid?
Answer:
Two replaceable hydrogen atoms are present in sulphuric acid.

Question 28.
Give an example of an organic base.
Answer:
Methylamine (CH₃NH₂)

Question 29.
Give an example of an inorganic gaseous compound which has basic property.
Answer:
Ammonia (NH₃)

Question 30.
What are the colour of methyl red in acidic and alkaline medium?
Answer:
The colour of methyl red in acidic medium is red and that in alkaline medium is yellow.

Question 31.
What is an Indicatlor?
Answer:
An indicator: It is some weak organic acid or base which indicates distinctive colours in acid, alkali and neutral solutions.

Question 32.
Why the farmers add slaked lime in the soil?
Answer:
Farmers add slaked lime (Calcium hydroxide) to reduce acidity of the soil.

Question 33.
What should we take in case of acidity and why?
Answer:
We should take antacid in case of acidity because magnesium hydroxide is present in antacid which neutralises excess HCl produced in the stomatch.

Question 34.
Give one example of basic salt.
Answer:
Basic lead nitrate [Pb(OH)NO₃

Short answer type questions 

Question 1.
What are oxides ? What are the types of oxides?
Answer:
Oxides : Binary compounds of oxygen with any element (metallic and non-metallic) are called oxides.
Types of oxides :

• Acidic oxide
• Basic oxide
• Neutral oxide
• Amphoteric oxide
• Peroxide
• Mixed oxide
• Poly-oxide
• Sub-oxide
• Super-oxide

Question 2.
Define acidic oxide.
Answer:
Acidic oxide : An acidic oxide is an oxide of a non-metal usually. It reacts with an alkali or a base to form salt and water. Example : CO₂, SO₂, P₂O. etc.

Question 3.
Define basic oxide.
Answer:
Basic oxide: A basic oxide is usually an oxide of a metal usually. It reacts with an acid to produce salt and water. Example: CaO, MgO. CuO etc.

Question 4.
Define neutral oxide.
Answer:
Neutral Oxide : The oxides of certain non-metals which neither react with acid nor with base are called neutral oxide. Example : CO, H₂O, NO etc.

Question 5.
Define amphoteric oxide.
Answer:
Amphoteric oxide : The oxides of certain weak electropositive metals which react with both acids and bases to form salts and water are called amphoteric oxides. Example : Al₂O₃, ZnO, PbO etc.

Question 6.
What are acids ?
Answer:
Acids : Ordinarily, an acid is a compound, the molecules of which contain one or more hydrogen atoms replaceable partially or completely, directly or indirectly by a metal or a group of elements behaving like a metal to form salt. Example : HCl, H₂SO₄, HNO₃ etc.

Question 7.
What are bases ?
Answer:
Bases : A base in general is an oxide or hydroxide of a metal and when reacts with an acid produces salt and water. Example : Al₂O₃, Al(OH)₃, CaO etc.

Question 8.
What are alkalis ?
Answer:
Alkalis : Water-soluble hydroxide of metals are called alkalis. Example : NaOH, KOH etc.

Question 9.
What are salts ? What are the different types of salt ?
Answer:
Salt : The replaceable hydrogen atoms in an acid when replaced by metal or basic radical partially or fully then the compound so produced is called salt.
Types of salt:

•  normal salt
• acid salt
• basic salt
• double salt and
• complex salt

Question 10.
Define normal salts.
Answer:
Normal salts : The salts produced by complete displacement of all the replaceable hydrogen atom or atoms present in the molecule of an acid by a metal or a radical acting like metal are called normal salts. Example : NaCl, Na₂SO₄, Na₃PO₄ etc.

Question 11.
Define acid salts or bi-salts.
Answer:
Acid salts or bi-salts : The salts produced by the partial displacement of the replaceable hydrogen atoms present in the molecule of an acid by a metal or a radical acting like metal are known as acid salts or bi-salts. Example : NaHSO₄, NaHCO₃, Na₂HPO₄ etc.

Question 12.
Define basic salt.
Answer:
Basic salt: Salts formed by the partial displacement of oxide or hydroxide of alkalis by an acid is known as basic salt. Example: Pb(OH)Cl, Pb(OH)NO₃ etc.

Question 13.
Define double salt.
Answer:
Double salt : It is formed by the association of two or more normal salts.
Example : K_BSO4, Al₂(SO₄)₃, 24H₂O; (NH₄)₂ SO₄, FeSO₄, 6H₂O etc.

Question 14.
Define complex salt.
Answer:
Complex salt : Salts that dissociate in water to give one simple ion and one complex ion are called complex salts.
Example : K₄[Fe(CN)₆], [Cu(NH₃)₄]SO₄, [Ag(NH₃)2Cl] etc.

Question 15.
What is neutralisation reaction ?
Answer:
The reaction in which equivalent amount of an acid reacts with equivalent amount of a base and therefore the properties of acid and base are completely lost by forming salt and water is called neutralisation reaction.

Question 16.
What do you mean by titration ?
Answer:
Titration : The process by which bases are neutralised with the help of acids or vice-versa is known as titration.

Question 17.
What is an indicator ?
Answer:
The substances which indicate the completion moment or end point of a titration reaction by changing their own colour are known as indicators.
Example : Methyl yellow, Methyl orange.

Question 18.
Which ions disappear in neutralisation ?
Answer:
Ions disappear in neutralisation: H⁺ ions generated by acid and OH^– ions generated by alkali disappear in a neutralisation process.

Question 19.
What is vanishing colour ? How is it prepared ?
Answer:

•  Vanishing colour : It is a coloured solution which becomes colourless on exposure to air for a passage of time.
• Preparation of vanishing colour : It is a dilute solution of ammonia in water with a few drops of phenolpthalein. Its colour is deep pink.

Question 20.
How does vanishing colour act ?
Answer:
Action of vanishing colour : When the pink solution of vanishing colour is spread over a white linen, the linen turns pink. On allowing the pink-coloured wet linen to dry in air, the ammonia from the solution evaporates and therefore the linen regains its original white colour.

Question 21.
Can a vanishing colour be prepared with dilute NaOH solution ?
Answer:
Vanishing colour cannot be prepared with dilute NaOH solution due to the fact that sodium hydroxide (NaOH) does not evaporate.

Question 22.
Why is acid called a proton donor ?
Answer:
Acid is called a proton donor: Acids produce cation (H⁺) in aqueous solution. H⁺ is known as proton. So acid is called a proton donor.

Question 23.
Why is aikali called a proton acceptor ?
Answer:
Alkali is called a proton acceptor: Alkali produces hydroxyl ion (OH^–) in aqueous solution, hydroxyl accepts proton (H⁺) coming from aqueous solution of acid to produce water. So, alkali is called a proton acceptor.

Question 24.
Aqueous solution of HCl turns blue litmus red but HCl vapour does not, why ?
Answer:
Reason: HCl vapour is a covalent compound, so it does not ionises in vapour state. In aqueous solution, HCl produces H₃O⁺ (hydronium ion) and acts as an electrovalent compound which turns blue litmus red.

Question 25.
Sodium carbonate is a neutral salt, but its aqueous solution is alkaline in nature, why ?
Answer:
Reason: Sodium carbonate reacts with water and produces caustic soda which is a strong alkali, and a weak acid carbonic acid. Due to the production of strong alkali, after neutralisation of H⁺, there is excess OH^– (hydroxyl ion) in the solution. So the solution will be alkaline in nature.
Equation : Na₂CO₃ + H₂O → (Na⁺ + OH^–) + H₂CO₃

Question 26.
Why CaH₂ and CH₄ are not acids ?
Answer:
CaH₂ and CH₄ are not acids : Calcium hydride (CaH₂) and methane (CH₄) do not produce H⁺ ions in aqueous solution, so these are not acids, for every acid produces H⁺ ions in aqueous solution.

Broad answer type questions 

Question 1.
What is Arrhenius concept of acid ?
Answer:
Arrhenius concept of acid : According to Arrhenius, a Swedish scientist, an acid is a chemical compound that dissociates in water, producing H⁺ ions (cations) as the only positive ions. H⁺ ion is called proton. So an acid is known as proton donor. But H⁺ (proton) does not remain in free state in solution, it attains stability being attached to a molecule of the solvent. In water solution proton combines with a water molecule to produce hydronium ion (H₃O⁺).
Example:

Question 2.
What is Arrhenius concept of base?
Answer:
Arrhenius concept of base: According to Arrhenius, a base is a chemical compound which produces OH^– ions (hydroxyl ions) as the only anions when they are dissolved in water.
Example:

Question 3.
Acids are hydrogen compounds but all hydrogen compounds are not acids — Explain.
Answer:
Explanation : It is to be noted that any acid contains hydrogen but any compound containing hydrogen is not an acid. A compound containing one or more than one hydrogen atom in its molecule will be termed as acid only when its hydrogen atom or atoms are replaceable by a metal and the product thus obtained must be a salt.

Example : None of the four hydrogen atoms of methane (CH₄) can be replaced by a metal. The metals like sodium, potassium displace hydrogen from water (H₂O) but the product obtained in each case is not a salt. So, methane and water, though they contain hydrogen atoms in their molecules are not acids.

Question 4.
What are the important properties of acids ?
Answer:
Important properties of acids :

• Taste : Generally, the aqueous solutions of acids have a sour taste.
• pH value : pH value of an acid is generally less than 7.
• In aqueous solution, the acids conduct electricity.
• Reactions with metals : The solutions of acids in water react with many metals like zinc, magnesium, iron etc. which are more electropositive than hydrogen with the liberation of hydrogen gas and formation of corresponding salts.
• Reactions with oxides and hydroxides of metals: Acids react with metallic oxides or hydroxides producing salts and water.
• Reactions with carbonates and bicarbonates : Acids are also characterised by their tendency to react with metallic carbonates and bicarbonates evolving carbon dioxide.
• According to Arrhenius theory of electrolytic dissociation, the acids in aqueous solutions produce hydrogen ions as the positive ions or cations (H⁺).
• The aqueous solution of an acid turns blue litmus red.

Question 5.
What are the important properties of alkalis ?
Answer:
Important properties of alkalis :

• Alkalis react vigorously with acids to produce salts and water,
• Their aqueous solution turn red litmus blue.
• The aqueous solutions of alkalis are soapy to touch and conduct electricity.
• The alkalis in aqueous solution ionise to produce hydroxyl ions (OH^–).

Question 6.
State the properties of acids and bases with respect to indicators.
Answer:
Properties of acids with respect to indicators : The aqueous solution of acids turns —

• blue litmus to red.
• orange coloured methyl organge to pinkish red :
• phenolphthalein remains colourless in aqueous solution of an acid.

Properties of bases with respect to indicators : The aqueous solution of bases turns  —

• red litmus to blue.
• orange coloured methyl orange to yellow.
• olourless phenolphthalein solution to pink.

Question 7.
Discuss the role of indicators in case of titration.
Answer:
Uses of indicators in titration reaction : Indicators are able to determine the end-point of a titration reaction. List of some well-known indicators are given below

Indicator		Colour changes in
		Neutralsolution	Acidic solution	Alkaline solution
1.	Litmus	Violet	Red	Blue
2.	Methyl-orange	Orange	Red	Yellow
3.	Phenolphthalein	Colourless	Colourless	Pink

The selection of a suitable indicator to determine the correct end point of a reaction :

• Strong acid and weak base: Methyl orange
• Weak acid and strong base: Phenolphtyalein
• Strong acid and strong base: Any indicator
• Weak acid and weak base: No suitable indicator.

Question 8.
How to ascertain whether a given colourless solution is acidic or alkaline by a simple test?
Answer:
Test: A strip of filter paper is dipped in red litmus solution. As a result, the colour of the paper turns red. The piece of paper is dried and the dry piece of paper is dipped in the given solution. If its colour remains red, the given solution is acidic. But if the colour of the litmus paper turns blue, the given solution is alkaline.

Question 9.
What is basicity of an acid and acidity of a base?
Answer:
Basicity of an acid : Basicity of an acid is expressed by the number of replaceable hydrogen atoms present in each molecule of the acid.
Example: Monobasic acid : HCl, HNO.

(i.e. basicity is 1)
Dibasic acid : H₂
Tribusic acid : H₃PO₄
(i.e. basicity is 3)

Acidity of a base : Acidity of a base is expressed by the number of hydroxyl groups present in each molecule of the base.
Example: Monoacidic base : NaOH, KOH
Diacidic base : Ca(OH)₂, Mg(OH)₂
Triacidic base : Al(OH)₃

Question 10.
All alkalis are bases but all bases are not alkalis – Explain.
Answer:
Explanation : Bases are the oxides or hydroxides of metals and react with acids to produce salts and water. The bases which are soluble in water are known as alkalis. Ferric hydroxide (Fe(OH₃), Zinc hydroxide (Zn(OH₂), Aluminium hydroxide (Al(OH₃), Sodium hydroxide (NaO), Potassium hydroxide (KOH) all are bases but not alkalis. Among them only water soluble metallic hydroxides NaOH and KOH are alkalis. From the above examples of bases and alkalis, it is clear that all alkalis are bases but all bases are not necessarily alkalis.

Well structured WBBSE Class 9 Physical Science MCQ Questions Chapter 4.1 Ideas of History can serve as a valuable review tool before exams.

Atomic Structure Class 9 WBBSE MCQ Questions

Multiple Choice Questions :

Question 1.
Atoms of which element have no neutron?
(i) Oxygen
(ii) Hydrogen
(iii) Carbon
(iv) Nitrogen
Answer:
Hydrogen

Question 2.
Which force binds the nucleons ?
(i) Gravitational force
(ii) Electrostatic force
(iii) Nuclear force.
(iv) Electromagnetic force
Answer:
Nuclear force.

Question 3.
Which scientist first enunciated the atomic concept of matter ?
(i) Newton
(ii) Einstein
(iii) Dalton
(iv) Rutherford
Answer:
Dalton.

Question 4.
The atomic number is –
(i) mass of an atom
(ii) total number of protons and neutrons
(iii) number of protons
(iv) number of neutrons
Answer:
Total number of protons and neutrons.

Question 5.
The maximum possible number of electrons in the outermost shell of an atom is –
(i) 16
(ii) 18
(iii) 8
(iv) 10
Answer:
8.

Question 6.
The magnitude of the negative charge an electron carries is–
(i) 1.602 × 10^-2 coulombs
(ii) 1.602 × 10^-9 coulombs
(iii) 1.602 × 10^-19 coulombs
(iv) 1.602 × 10^-25 coulombs
Answer:
1.602 × 10^-19 coulombs.

Question 7.
If two isotopes of a certain element are known, it follows that the atoms of the element :
(i) differ chemically from each other
(ii) have different number of electrons surrounding their nuclei.
(iii) have different number of neutrons in their nuclei.
(iv) have the same mass number.
Answer:
Have different number of neutrons in their nuclei.

Question 8.
which of the following are example of isotopes ?
(i) \({ }_6^{14} \mathrm{C}\) and \({ }_7^{14} \mathrm{N}\)
(ii) \({ }_{19}^{40} \mathrm{K}\) and \({ }_{20}^{40} \mathrm{Ca}\)
(iii) \({ }_8^{16} \mathrm{O}\) and \({ }_8^{18} \mathrm{O}\)
(iv) \({ }_6^{14} \mathrm{C}\) and \({ }_8^{16} \mathrm{O}\)
Answer:
\({ }_8^{16} \mathrm{O}\) and \({ }_8^{18} \mathrm{O}\).

Question 9.
Sodium atom forms a cation by losing one electron. The cation will be :
(i) S⁺
(ii) Na⁺
(iii) K⁺
(iv) H⁺
Answer:
Na⁺

Question 10.
A deutron contains –
(i) a neutron and a positron
(ii) a neutron and a proton
(iii) a neutron and two protons
(iv) a proton and two neutrons
Answer:
A neutron and a proton.

Question 11.
The nucleus of an atom contains –
(i) electrons
(ii) protons alone
(iii) neutrons alone
(iv) protons and neutrons
Answer:
Protons and neutrons.

Question 12.
The neutron was discovered by –
(i) J. J. Thomson
(ii) G.T. Seaborg
(iii) E. Rutherford
(iv) James Chadwick
Answer:
James Chadwick.

Question 13.
The number of electrons in the nucleus of \({ }_6^{12} \mathrm{C}\) is –
(i) 6
(ii) 12
(iii) 0
(iv) 3
Answer:
0.

Question 14.
Positron is – (i) \({ }_{-1}^0 \mathrm{e}\)
(ii) \({ }_{+1}^0 \mathrm{e}\)
(iii) \({ }_{+1}^1 \mathrm{H}\)
(iv) None of these
Answer:
\({ }_{+1}^0 \mathrm{e}\).

Question 15.
The average distance of an electron in an atom from its nucleus is in the order of :
(i) 10⁶ m
(ii) 10^-6 m
(iii) 10^-10 m
(iv) 10^-15 m
Answer:
10^-10 m.

Question 16.
A neutral atom (atomic no. > 1) contains
(i) Proton oniy
(ii) Neutron + Proton
(iii) Neutron + Electron
(iv) Neutron + Proton + Electron
Answer:
Neutron + Proton + Electron

Question 17.
The radius of an atom is in the order of –
(i) 10^-16 cm
(ii) 10^-13 cm
(iii) 10^-15 cm
(iv) 10^-8 cm
Answer:
10^-8 cm}.

Question 18.
Chlorine atom differs from chlorine ion in the number of which of the following :
(i) Protons
(ii) Neutrons
(iii) Electrons
(iv) Both Protons and Neutrons
Answer:
Electrons.

Question 19.
Neutrons are present in the nuciei of all elements except –
(i) Hydrogen
(ii) Oxygen
(iii) Deuterium
(iv) Chlorine
Answer:
Hydrogen.

Question 20.
When electrons revolve in stationary orbits –
(i) there is no change in energy level
(ii) they become stationary
(iii) they are gaining kinetic energy
(iv) there is increase in energy
Answer:
There is no change in energy level.

Question 21.
Which of the following isoelectronic species has less electrons than protons ?
(i) O^2-
(ii) F^–
(iii) Na^*
(iv) Mg²⁺
Answer:
O^2-.

Question 22.
As we move away from the nucleus, the energy of an orbit –
(i) decreases
(ii) increases
(iii) remains unchanged
(iv) none of these
Answer:
increases.

Question 23.
If one electron is added to the outermost shell of a chlorine atom, it produces a –
(i) new atom
(ii) anion
(iii) cation
(iv) there will be no change
Answer:
anion

Question 24.
Neucleons are –
(i) only protons
(ii) only neutrons
(iii) protons and neutrons
(iv) protons, electrons and neutrons
Answer:
protons and neutrons.

Question 25.
Rutherford’s α-particle scattering experiment led to the discovery of
(i) electrons
(ii) protons
(iii) neutrons
(iv) atomic neucleus
Answer:
atomic neucleus

Question 26.
The electronic configuration of sodium atom is –
(i) 2,8,3
(ii) 2,8,8
(iii) 2,8,1
(iv) 2,5,3
Answer:
2,8,1

Question 27.
The valency of the element having electronic configuration 2,8,8,1 is-
(i) 1
(ii) 2
(iii) 3
(iv) 7
Answer:
1

Fill in the blanks :

1. Most part of an atom is _____.
Answer:
vacant or empty.

2. The number of protons in an atom of an element is the ______ number.
Answer:
atomic.

3. Atom is the smallest part of an _______ that participates in chemical reactions but does not usually exist freely in nature.
Answer:
element.

3. Chemical combination takes place by the union of ______ number of atoms of the elements in simple ratios 1: 3,2: 3,1: 2 etc.
Answer:
integral.

4. Only the nucleus of ordinary _______ does not contain any neutron.
Answer:
hydrogen.

5. Constituents of the nucleus are called ______.
Answer:
nucleons.

6. An ion is an atom or a group of atoms that carries ______ charge.
Answer:
electric.

7. Since the mass of ______ is negligibly small, the whole mass of an atom is supposed to be concentrated at its nucleus.
Answer:
electrons.

8. isotope are the atoms of the same element which have the same _______ number but different mass number.
Answer:
atomic.

9. The tota! number of electrons present in all the shells of an atom is equal to the number of ______ present in its nucieus.
Answer:
protons.

10. Round the nucleus, negatively charged particles called _______ revolve in different paths.
Answer:
electrons.

11. Nuclear force is a _______ attractive torce active within the range of 2 × 10-15 metre that acts between the nucieons.
Answer:
short range.

12. The nucleus of an atom is _______ charged.
Answer:
positively.

13. The outermost orbit of an atom can possess maximum _______ electrons.
Answer: eight.

14. To calculate the number of neutrons in an atom, we substract its ________ from its mass number.
Answer:
atomic number.

15. The whole mass and positive charges of an atom remain confined to its _______.
Answer:
nucleus.

16. For the isotope \({ }_6^{13} \mathrm{C}\), the number of neutrons is ______
Answer:
7.

17. The atomic number of an atom of an element is the number of ______ in its nucleus.
Answer:
protons.

18. The maximum capacity of a shell to accommodate electrons is given by the general rule _______.
Answer:
2 n².

19. The atomic number of potassium is _______.
Answer:
19.

20. Hydrogen has ______ isotopes.
Answer:
three.

21. _______ recognised that an element might have atoms of identical chemical properties but of different atomic weights.
Answer:
Soddy.

22. An electron has wave as well as ________ nature.
Answer:
particle.

23. A proton is ______ times heavier than an electron.
Answer:
1837.

24. An ______ is a well-defined circular path in which the electron revolves.
Answer:
orbit.

25. The lowest energy level in an atom is _______ level.
Answer:
K.

26. Electrically charged atoms are called _______.
Answer:
ions.

27. The particles, which Thomson called corpuscles, later came to known as ______.
Answer:
electrons.

28. Becquerel put small crystals of _______ upon the black paper.
Answer:
Potassium uranyl sulphate [K(UO₂)(SO₄)₃.3H₂ O]

29. The diameter of the atom is about while that of the nucleus is -.
Answer:
10^-8 cm}, 10^-13 cm}

30. 114 Be + 42 He ______ + 10n.
Answer:
\({ }_6^{14} \mathrm{C}\) →

31. The definite small quantity of energy is known as energy –
Answer:
quanta.

32. The radius of nucleus is – time less than that of the atom.
Answer:
10⁵.

33. The particles in the nucleus are called
Answer:
Nucleons

Well structured WBBSE Class 9 Physical Science MCQ Questions Chapter 7 Sound can serve as a valuable review tool before exams.

Sound Class 9 WBBSE MCQ Questions

Multiple Choice Questions :

Question 1.
What should be the minimum frequency of a vibrating body for producing sound?
(i) 2000 Hz
(ii) 100 Hz
(iii) 20 Hz
(iv) 150 Hz
Answer:
20 Hz

Question 2.
How does pitch of sound depend on frequency ?
(i) pitch increases with decrease of frequency
(ii) pitch decreases with increase of frequency
(iii) pitch increases or decreases accordingly with increase or decrease of frequency
(iv) none of these.
Answer:
Pitch increases or decreases accordingly with increase or decrease of frequency.

Question 3.
The unit of frequency is :
(i) ohm
(ii) mho
(iii) Hz
(iv) torr
Answer:
Hz

Question 4.
Migraine is caused by –
(i) sound pollution
(ii) air pollution
(iii) water pollution
(iv) soil pollution.
Answer:
sound pollution

Question 5.
When one end of a long iron pipe is struck –
(i) two
(ii) three
(iii) four sounds are heard.
Answer:
Two

Question 6.
The relation between velocity (v) of a sound with its density (d) is :
(i) v ∝ \(\frac{1}{d}\)
(ii) v ∝ \(\frac{1}{\sqrt{d}}\)
(iii) v = \(\frac{1}{d}\)
(iv) v = \(\frac{1}{\sqrt{d}}\)
Answer:
v ∝ \(\frac{1}{\sqrt{d}}\)

Question 7.
The relation between velocity (v) of a sound with the absolute temperature (T) is –
(i) v ∝ √T
(ii) v ∝ T
(iii) v = T
(iv) v = √T
Answer:
v ∝ √T

Question 8.
If the size of the sounding body be bigger then intensity of sound will be :
(i) increased
(ii) decreased
(iii) remains same
Answer:
increased

Question 9.
The velocity of sound through iron is –
(i) 4 times
(ii) 10 times
(iii) 15 times
(iv) 16 times than the velocity of sound through air.
Answer:
15 times

Question 10.
The sensation due to short sound remains for about –
(i) \(\frac{1}{20}\) second
(ii) \(\frac{1}{10}\) second
(iii) \(\frac{1}{30}\)
(iv) \(\frac{1}{5}\) second and is known as persistence of hearing.
Answer:
\(\frac{1}{10}\) second.

Question 11.
Relation among velocity (v), wavelength (λ) and frequency (v) is :
(i) v = v λ
(ii) v = v λ
(iii) λ = v\v
(iv) v = vλ
Answer:
v = v\λ

Question 12.
If the velocity of a wave havingwave length 1.7 m is 340 m / s then what will be the frequency of the wave?-
(i) 20 Hz
(ii) 200 Hz
(iii) 2000 Hz
(iv) 100 Hz
Answer:
200 Hz

Question 13.
If the tones present in a note are of frequencies 100,150,200,280, 300,450 and 500 , what is the fundamental tone ?-
(i) 100
(ii) 280
(iii) 500
(iv) 450
Answer:
100

Question 14.
The sound of lightning is heard after 6 sec. it is seen. If the velocity of sound is 330 m/s at that time what will be the distance of the cloud?
(i) 1980 m
(ii) 2000 m
(iii) 330 m
(iv) 350 m
Answer:
1980 m

Question 15.
Wave motion transfers –
(i) matter
(ii) energy
(iii) momentum
(iv) all of these
Answer:
energy

Question 16.
Distance between two successive compressions is –
(i) \(\frac{\lambda}{2}\)
(ii) \(\frac{\lambda}{4}\)
(iii) λ
(iv) 2 λ
Answer:
λ

Question 17.
Sound wave can travel –
(i) only in solid
(ii) only in liquid
(iii) only in gas
(iv) in all of these
Answer:
in all of these

Question 18.
The speed of sound in air-
(i) decreases with the increase in temperature
(ii) remains the same with the increase in temperature
(iii) remains the same with the decrease in temperature
(iv) increases with the increase in temperature.
Answer:
decreases with the increase in temperature.

Question 19.
Sound wave is –
(i) transverse in nature
(ii) longitudinal in nature
(iii) both transverse and longitudinal in nature
(iv) none of these
Answer:
longitudinal in nature

Question 20.
Distance between a compression and a rarefaction is
(i) \(\frac{\lambda}{4}\)
(ii) \(\frac{\lambda}{2}\)
(iii) λ
(iv) 2λ
Answer:
\(\frac{\lambda}{2}\)

Question 21.
Which of the following waves is a mechanical wave ?
(i) light wave
(ii) Sound wave
(iii) X-rays
(iv) Ultra-violet ray
Answer:
Sound wave

Question 22.
The speed of sound is maximum in –
(i) air
(ii) hydrogen
(iii) water
(iv) iron
Answer:
iron

Question 23.
In SONAR, we use –
(i) infrasonic
(ii) radio waves
(iii) audible sound
(iv) ultrasonic
Answer:
ultrasonic

Question 24.
Infra sound can be heard by –
(i) bat
(ii) rhinoceros
(iii) dolphin
(iv) human being
Answer:
rhinoceros.

Fill in the blanks :

1. Sound moves ______ through solid than through liquid.
Answer:
faster

2. Sound of frequency above is known as ultrasonic ______ sound.
Answer:
20,000 Hz

3. The sound of frequency below 20 Hz is called ______ sound.
Answer:
subsonic

4. Sound originates from a mechanically ______ body.
Answer:
vibrating

5. We can hear sound only if the frequency of a vibrating body lies within the limits of ______ to 20,000 Hz.
Answer:
20

6. The maximum displacement of a particle in a medium on either side of its mean position is known as ______.
Answer:
amplitude.

7. Time period is the time required by a particle to make ______ complete oscillation.
Answer:
one

8. The number of complete oscillations of a body per second is known as its ______ .
Answer:
frequency

9. To perceive the initial and the reflected sounds clearly, the minimum time interval between these should be at least ______ second.
Answer:
\(\frac{1}{10}\)

10. Pitch is the characteristic of a musical sound that distinguishes a sharp sound from a ______ sound.
Answer:
flat

11. Pitch increases corresponding to ______ in wavelength.
Answer:
decrease

12. Pitch of sound produced by a vibrating air column increases with the ______ of the air column.
Answer:
length

13. Note is analogous to white or polychromatic light and ______ analogous to light of a single colour.
Answer:
tone

14. In a note, the sound of least frequency is called the ______ tone.
Answer:
fundamental

15. A ______ continuous noise of level may cause tension, headache, migraine, neurological diseases.
Answer:
high

16. A ______ is necessary for the propagation of sound.
Answer:
medium

17. If the frequency of sound is high, pitch of the sound will also be ______.
Answer:
high

18. The particular harmonic whose frequency is double the fundamental frequency is called ______ of the fundamental.
Answer:
octave

19. Sound is a form of ______ .
Answer:
energy

20. A ______ medium is required for the propagation of sound.
Answer:
material

21. The velocity of sound in a gas is proportional to the square root of the _______ of the gas.
Answer:
density

22. The source of sound in laboratory is _______.
Answer:
sonometer

23. _______ is the safe intensity of sound according to WHO.
Answer:
45 decibel (or 45 dB)

24. Sound originates from a body that ______
Answer:
vibrates

25. The maximum displacement of particle in vibratory motion is called ______.
Answer:
amplitude

26. Light moves as _____ wave
Answer:
transverse

27. Sound travels through gases as ______ waves
Answer:
longitudinal

28. A wave transmits ______ through the medium
Answer:
energy

29. the unit of frequency is ______.
Answer:
Hertz.

30. For a echo to be heard the minimum distance between the source and the reflector should be ______ m
Answer:
16.6

31. ultrasonic waves have frequency greater than ______ Hz.
Answer:
20,000

32. Loudness, pitch and quality are the _____ of musical sound
Answer:
characteristics.

Comprehensive WBBSE Class 9 Physical Science Notes Chapter 7 Sound can help students make connections between concepts.

Sound Class 9 WBBSE Notes

Wave motion and wave : When a disturbance produced in one portion of space travels to another portion, without involving the transfer of any material with it, the motion of the disturbance is called wave motion.

The disturbance itself is called a wave.

• Longitudinal wave : If the particles of a medium move along the direction of the motion of the wave, the wave is called a longitudinal wave.
• Transverse wave : If the particles of a medium move perpendicular to the direction of the motion of the wave, the wave is called a transverse wave.

Characteristics of sound :

• Sound is produced by a vibrating source.
• Sound travels as a longitudinal wave through a material medium.
• A vibrating source produces compression and rarefaction pulses, one after the other in the medium. These pulses travel one behind the other as a sound wave.
• In sound propagation, it is the energy of the sound that travels and not the particles of the medium.
• Sound cannot travel in vacuum.

Oscillation : The change in density from one maximum value to the minimum value and again to the maximum value makes one complete oscillation.

Periodic motion : Any motion which repeats itself after a fixed interval of time is called periodic motion.

All oscillatory motions are periodic motions but all periodic motions are not oscillatory motions.

Examples of oscillatory motion :

• Motion of a pendulum
• To and fro motion of a loaded spring.
• Vibrations of the wire of a stringed musical instrument.
• Motion of liquid contained in U-tube when it is compressed once in one limb and left to itself etc.

Compression pulses correspond to regions of high density and pressure, whereas rarefaction pulses correspond to regions of low density and pressure in the medium.

Some definitions :
(a) Displacement : The displacement of a particle in oscillatory motion at any instant is the distance of the particle from its equilibrium position at that instant.

(b) Amplitude : The maximum displacement of a particle in oscillatory motion an either side of its equilibrium position is called amplitude of its motion.

(c) Time period : The time required for a complete oscillation by a particle in oscillatory motion is called its time period. It is usually denoted by T and is measured in second (s).

(d) Wavelength : The distance between two consecutive compressions or two consecutive rarefactions is called the wavelength (A,). It is measured either in centimetre (cm) or in metre (m).

(e) Frequency : The number of oscillations of the density of the medium at a place per unit time is called the frequency (v).
\(v=\frac{1}{T}\) (V = frequency, T = time period)

Frequency is measured in per second (s^-1), which is also known as hertz (Hz).
Relation among wavelength, frequency and velocity of a wave : In a medium, the distance travelled by the wave in unit time is called the velocity of the wave in the medium.

The number of complete vibrations made per unit time by a particle in the path of wave is called the frequency of the wave.
Velocity of sound = Frequency × wave length
i..e v = vγ

Reflection of sound : When sound waves hit on the boundary separating two homogeneous media, a portion of sound changes its direction from the surface of separation and returns to the first medium. This is known as reflection of sound.

Laws of reflection of sound :

• The incident sound, reflected sound and the normal on the surface of separation through the point of incidence remain on the same plane.
• The angle of incidence is equal to the angle of reflection.

Mach number : The ratio of speed of a body to the speed of sound is called mach number of the body. If the mach number of a body is greater than 1, then the velocity of the body is supersonic velocity.

Echo : When a sound after reflection from some reflector is again heard separated from the original sound, then this reflected sound is known as an echo.

Condition for echo : For occurrence of an echo the minimum distance between the source of sound and a reflector of sound producing echo should be around 16.6m.

Reverberation : The persistence of sound due to repeated reflection and its gradual fading away is called reverberation of sound.

Subsonic sound : We cannot hear the sound generated from the sources having frequency less than 20Hz. These sounds are known as subsonic sound or infrasonic sound.

Ultrasonic sound : We can not hear sound generated from the sources having frequency 20,000 Hz (20 kHz) or more. These sounds are known as ultrasonic sound.

Audible sound : We can hear only the sounds of the sources producing about 20 Hz to 20,000 Hz in frequencies. The sounds are known as audible sound.

Applications of ultrasound :

• Ultrasound is generally used to clean parts located in hard-to- reach places, for example spiral tube, odd shaped parts, electronic components etc.
• Ultrasound is used for drilling holes or making cuts of desired shape.
• Ultrasounds can be used to detect cracks and flaws in megtal blocks.
• Ultrasonic waves have given doctors powerful and safe tools for imaging human organs.

(a) Echocardiography is a technique in which ultrasonic waves, reflected from various part of the heart, form an image of the heart.

(b) Ultrasonography is routinely used to show image of patient’s organs such as the liver, gall bladder, uterus etc. to doctors.

(c) A technique called phacoemulsification (phaco stands for the lens of the eye) uses the power of ultrasound to break the hardened gel into tiny pieces.

(d) Ultrasound is also employed to break small ‘stones’ that form in the kidneys into fine grains.

• Sonar : The acronym SONAR stands for SOund Navigation AND Ranging. This is a method for detecting and finding the distance of objects under water by means of reflected ultrasonic waves.
• Note : A sound with a number of frequencies is called a note.
• Tone : A sound with a single frequency is called a tone.

Characteristics of Musical Sound :

• Loudness or intensity : Loudness is the measure of how intense or loud a sound is, loudness of a sound is measured by energy contained per unit volume of the medium through which sound passes.
• Pitch : The pitch is that physical cause by which we can distinguish a shrill sound from a flat one of same intensity. It depends on the frequency of sound.
• Quality or timbre : Quality is that characteristic of a musical sound by which we can distinguish a note sounded on a musical instrument from a note of same pitch and intensity produced.

Structure of human ear : The structure of human ear can be divided into three parts :

(i) Outer ear : The outer ear consists of pina and ear canal. Pina is the part which is visible from the outside. The function of the outer ear is to guide sound waves to the middle ear.

(ii) Middle ear : The middle ear is separated from the ear canal by a tightly stretched membrane called eardrum. Beyond it, three interconnected bones called the hammer, anvil and stirrup are situated.
The function of the middle ear is to pick up, amplify and transmit sound waves to inner ear.

(iii) Inner ear : The inner ear consists of a liquid-filled coiled tube called cochlea, which is shaped like a snail. This tube is connected to the stirrup. The cochlea has special sensory cells called hair cells, which are so called because of the hairlike structures that stick out of them. The hair cells are connected to the auditory nerve, which is connected to the brain.

Sound pollution : Any unpleasant and unwanted sound is noise. Sound pollution means creation of discomfort, disturbance and irritation which result to ill effects to mental and physical health.

Sources of sound pollution :

• The sound vehicle causes sound pollution.
• The machinery used in industry creates sound of intensity more than tolerance limit and thereby causes sound pollution.
• Loud-speakers, amplifiers, air horn played at top volume on the roads, are major sources of sound pollution.
• The sound crackers cracked in different functions, particularly marriages and Puja Festivals causes sound pollution.

Harmful effects of sound pollution :

• Sound pollution may cause hindrance in the growth of nervous system of a child in the womb, That is yet to be born.
• One may become deaf due to sound pollution.
• The sound of high intensity and pitch may cause increase of blood pressure, nervous breakdown, heart disease, loss of memory etc.

Remedial measures of sound pollution :

• Factories, smaller or big, should be prohibited in residential areas.
• Sound resistant measures may be taken in industry.
• Unnecessary use of horn in the vicinity of hospital, school etc. should be restricted.
• Airports should be away from residential area.
• Sound pollution awareness programme should be taken up at regular interval.

WBBSE Class 10 Physical Science Question Answer West Bengal Board

WBBSE Class 10 Physical Science Book Solutions West Bengal Board in English Medium

• Chapter 1 Concerns About Our Environment
• Chapter 2 Behaviour of Gases
• Chapter 3 Chemical Calculations
• Chapter 4 Thermal Phenomena
• Chapter 5 Light
• Chapter 6 Current Electricity
• Chapter 7 Atomic Nucleus
• Chapter 8 Physical and Chemical Properties of Matter
• • Chapter 8.1 Periodic Table and Periodicity of the Properties of Elements
  • Chapter 8.2 Ionic and Covalent Bonding
  • Chapter 8.3 Electricity and Chemical Reactions
  • Chapter 8.4 Inorganic Chemistry in the Laboratory and in Industry
  • Chapter 8.5 Metallurgy
  • Chapter 8.6 Organic Chemistry

WBBSE Class 10 Physical Science Book Solutions West Bengal Board in Hindi Medium

• Chapter 1 अपने पर्यावरण के प्रति अभिरुचि
• Chapter 2 गैसों का व्यवहार
• Chapter 3 रासायनिक परिकलन
• Chapter 4 उष्मीय घटनायें
• Chapter 5 प्रकाश
• Chapter 6 धारा विद्युत
• Chapter 7 परमाणु नाभिक
• Chapter 8 पदार्थ के भौतिक और रासायनिक गुण
  • Chapter 8.1 पदार्थ के भौतिक एवं रासायनिक गुण
  • Chapter 8.2 रासायनिक बन्धन
  • Chapter 8.3 विद्युत एवं रासायनिक अभिक्रियाएँ
  • Chapter 8.4 प्रयोगशाला एवं उद्योग में अकार्बनिक रसायन
  • Chapter 8.5 धातुकर्म
  • Chapter 8.6 कार्बनिक रसायन

WBBSE Class 10 Physical Science Syllabus West Bengal Board 2024

Chapter 1 Concerns About Our Environment

Chapter 1.1 Physical Science and Environment:
How can we interpret our environment in terms of physical science? – Structure of the atmosphere, The ozone layer, Greenhouse effect, and global warming.

Chapter 1.1 Rational Use of Energy:
Should we not conserve energy for future generations? – Harnessing energy resources for sustainable development.

Chapter 2 Behaviour of Gases

Chapter 2.1 Behaviour in the Bulk:
How does a gas behave under different conditions? – Pressure exerted by a gas and its volume, Boyle’s law, Charles’ law, Absolute temperature scale, Corrjbining Boyle’s and Charles’ laws, Ideal gas, Avogadro’s law, Ideal gas equation.

Chapter 2.2 Behaviour at the Molecular Level:
How can we describe behaviour of gases at the molecular level? – Molecular picture of an ideal gas, Deviation from ideal behaviour.

Chapter 3 Chemical Calculations

Chapter 3.1 Stoichiometric Equations:
How to answer ‘how much?’ – Conservation of mass in chemical reactions, Weight versus weight calculations.

Chapter 4 Thermal Phenomena

Chapter 4.1 Thermal Expansion:
How do different materials expand on heating? – Coefficient of expansion for solids, liquids, and gases.

Chapter 4.2 Thermal Conduction:
Does heat flow equally well through all solids? – Thermal conductivity.

Chapter 5 Light

Chapter 5.1 Reflection at Spherical Mirror:
What happens when the reflecting surface of a mirror is convex or concave? – Reflection of light at the spherical surface, Geometry of a spherical mirror, Reflection in the spherical mirror.

Chapter 5.2 Refraction of Light:
What happens when a light ray travels from one medium to another? – Laws of refraction, Familiarity with the structure of a glass slab and a prism, Refraction through a glass slab and a prism. Deviation of light rays due to refraction.

Chapter 5.3 Lenses:
How does a lens differ from a glass slab? – Familiarity with the structure of spherical lenses (concave and convex), Refraction through a lens (convex and concave), Image formation by thin lenses, Simple camera, and the Human eye as application image formation by a lens.

Chapter 5.4 Dispersion of Light:
What happens when white light passes through a prism?
Dispersion of white light as a result of refraction through a glass slab and a prism the idea of monochromatic and polychromatic and polychromatic light.

Chapter 5.5 Light Wave:
What kind of wave is a light wave, and how many different kinds of light waves are there? – Frequency, wavelength, and velocity of a wave, Uses and adverse effects of UV, X-ray, and gamma-ray, and Scattering of light.

Chapter 6 Current Electricity

Chapter 6.1 Electric Current, Potential Difference & EMF:
How does the charge start moving? – Electric charge, Electric potential difference, EMF, and electrical cell as a source of EMF, Electric current.

Chapter 6.2 Ohm’s Law:
What is resistance? – Relation between potential difference and current in a wire: Concept of resistance from Ohm’s law, EMF and internal resistance of a cell, Resistivity and conductivity, Series and parallel combination of resistances, Domestic circuits.

Chapter 6.3 Heating Effect of Electric Current:
How do we measure the consumption of electrical energy? – Joule’s law on the heating effect of current: Concept of electrical energy, Domestic uses of the heating effect of current.

Chapter 6.4 Electrical Power:
Electrical Power, the concept of a kilowatt-hour, B.O.T.

Chapter 6.5 Electromagnetism:
How does a current-carrying wire interact with a magnet? – Action of electric current on magnet: Ampere’s swimming rule, right-hand grasp rule, Action of magnet on current carrying wire: Fleming’s left-hand rule, Working principle of the motor.

Chapter 6.6 Electromagnetic Induction:
How can we get electricity from motion? – Concept of induced EMF and induced current, Basic idea of direct current and alternating current.

Chapter 6.7 Electric Generator:
How does a generator work? – Working principle and functioning of an electric generator.

Chapter 6.8 Domestic Electrical Circuit:
How are electrical wirings done in households? – Components used in domestic electrical circuit, Schematic representation of domestic electrical circuit in simplest form.

Chapter 7 Atomic Nucleus

Chapter 7.1 Radioactivity:
What is radioactivity and how is it related to the atomic nucleus? – Nature of α, β and γ rays, and their origin.

Chapter 7.2 Nuclear energy:
How can we get energy from the nucleus? – Concept of mass defect binding energy and fission, Fusion.

Chapter 8 Physical and Chemical Properties of Matter

Chapter 8.1 Periodic Table and Periodicity of the Properties of Elements:
What is meant by periodicity? – Brief history of the periodic table, Modern periodic table, Periodicity of properties of elements.

Chapter 8.2 Ionic and Covalent Bonding:
What holds the ions in a solid or atoms in a molecule? – Properties of ionic compounds, Ionic bonding, Properties of covalent compounds, Covalent bonding.

Chapter 8.3 Electricity and Chemical Reactions:
How are electricity and chemical changes related? – Electrolytes, Strong and weak electrolytes, Mechanism of electrical conduction in molten/solution states, Electrolysis, Applications of electrolysis.

Chapter 8.4 Inorganic Chemistry in the Laboratory and Industry:
Just how important are the inorganic chemicals? – Laboratory preparation of ammonia, Properties of ammonia, Major industrial uses and industrial manufacture of NH3 and urea, Laboratory preparation of hydrogen sulfide, Properties of H₂S, Laboratory preparation and major uses of N₂, Properties of N₂, Industrial manufacture of HCl, HNO₃, and H₂SO₄.

Chapter 8.5 Metallurgy:
How can we get the metals from their ores and use them? – Uses of Fe, Cu, Zn, and Al and their alloys, Ores, and minerals, Brief introduction to electronic theory of redox processes, Example of Thermite reaction, Metal corrosion: (a) Rusting of iron and its prevention (b) Corrosion of other metals and its health implications.

Chapter 8.6 Organic Chemistry:
How are some carbon compounds distinctly different from inorganic compounds of carbon? – Organic compounds are compounds of carbon, Tetravalency and catenation property of carbon, Structures of C₂H₆, C₂H₄, C₂H₂, Functional groups, Isomerism, Homologous series, IUPAC nomenclature of simple organic compounds, Industrial source and major uses of CH₄, C₂H₄, C₂H₂, LPG & CNG, Reactions of CH₄, C₂H₄, C₂H₂, Some other synthetic organic polymers, Biodegradable polymers, Uses and properties of ethyl alcohol and acetic acid, Harmful effects of methanol and ethanol, Denatured spirit.

WBBSE Class 10 Physical Science Blueprint for 1st Summative Evaluation (Total Marks – 40)

Theme/Sub-theme	MCQ (Group A)	VSA (Group B)	SA (Group C)	LA (Group D)	Total
1. Concerns About Our Environment	1 × 1	1 × 2	2 × 1	–	5
2. Behaviour of Gases	1 × 1	1 × 2	2 × 1	3 × 1	8
3. Light	1 × 3	1 × 3	2 × 1	3 × 2	14
4. Periodic Table and Periodicity of the Properties of Elements	1 × 1	1 × 1	2 × 1	3 × 1	7
5. Ionic and Covalent Bonding	1 × 1	1 × 1	2 × 2	–	6
Total	7	9	12	12	40

WBBSE Class 10 Physical Science Blueprint for 2nd Summative Evaluation (Total Marks – 40)

Theme/Sub-theme	MCQ (Group A)	VSA (Group B)	SA (Group C)	LA (Group D)	Total
1. Chemical Calculations	1 × 1	1 × 1	–	3 × 1	5
2. Thermal Phenomena	1 × 1	1 × 2	–	3 × 1	6
3. Current Electricity	1 × 2	1 × 2	2 × 1	3 × 2	12
4. Electricity and Chemical Reactions	1 × 1	1 × 2	–	3 × 1	6
5. Inorganic Chemistry in the Laboratory and in Industry	1 × 1	1 × 1	2 × 1	3 × 1	7
6. Metallurgy	1 × 1	1 × 1	2 × 1	–	4
Total	7	9	6	18	40

WBBSE Class 10 Physical Science Blueprint for 3rd Summative Evaluation/Selection Test (Total Marks – 90)

Theme/Sub-theme	MCQ (Group A)	VSA (Group B)	SA (Group C)	LA (Group D)	Total
1. Concerns about Our Environment	1 × 1	1 × 2	2 × 1	–	5
2. Behaviour of Gases	1 × 1	1 × 2	2 × 1	3 × 1	8
3. Chemical Calculations	1 × 1	–	–	3 × 1	4
4. Thermal Phenomena	1 × 1	1 × 1	–	3 × 1	5
5. Light	1 × 2	1 × 2	2 × 1	3 × 2	12
6. Current Electricity	1 × 2	1 × 2	2 × 1	3 × 2	12
7. Atomic Nucleus	1 × 1	1 × 1	–	3 × 1	5
8. Periodic Table and Periodicity of the Properties of Elements	1 × 1	1 × 2	–	3 × 1	6
9. Ionic and Covalent Bonding	1 × 1	1 × 1	2 × 2	–	6
10. Electricity and Chemical Reactions	1 × 1	1 × 2	–	3 × 1	6
11. Inorganic Chemistry in the Laboratory and Industry	1 × 1	1 × 2	2 × 1	3 × 1	8
12. Metallurgy	1 × 1	1 × 2	2 × 1	–	5
13. Organic Chemistry	1 × 1	1 × 2	2 × 1	3 × 1	8
Total	15	21	18	36	90

WBBSE Class 10 Solutions

Detailed explanations in West Bengal Board Class 9 Physical Science Book Solutions Chapter 4.2 Mole Concept offer valuable context and analysis.

WBBSE Class 9 Physical Science Chapter 4.2 Question Answer – Mole Concept

Very short answer type questions

Question 1.
What is mole of a substance ?
Answer:
Molecular weight of a substance expressed in gram is called mole of the substance.

Question 2.
What is molar volume of a substance ?
Answer:
Molar volume of a substance is the volume of gram molecular weight of any substance in gaseous state.

Question 3.
What is vapour density of a gas ?
Answer:
Vapour density is the ratio of weight of some volume of a gas or vapour to the weight of same volume of hydrogen at same temperature and pressure.

Question 4.
What is Avogadro’s number ?
Answer:
Avogadro’s number is the number of molecules present in 1 gram molecule of a substance.

Question 5.
What is the value of Avogadro’s number ?
Answer:
The value of Avogadro’s number is 6.023 × 10²³

Question 6.
How many molecules are present in 44g CO₂?
Answer:
6.023 x 10²³

Question 7.
What is the gram molecular weight of nitrogen ?
Answer:
28g

Question 8.
What is the atomicity of helium ?
Answer:
Atomicity of helium is 1.

Question 9.
What will be the volume of 22g of CO₂ at NTP ?
Answer:
\(\frac{22 \cdot 4}{2}\) lit = 11.2 lit

Question 10.
Who postulated the atomic theory ?
Answer:
John Dalton postulated the atomic theory.

Question 11.
What is the molar volume of a gas at NTP ?
Answer:
22.4 lit

Question 12.
Give the name of a gas which has same atom and molecule ?
Answer:
Argon (Ar)

Question 13.
Who gave the first concept of molecule ?
Answer:
idea about the molecules was put forward first by an Italian physicist, Amedeo Avogadro (1811).

Question 14.
Which element has atomicity 4 ?
Answer:
Phosphorus

Question 15.
What is the volume of 1 mole CO₂ at NTP ?
Answer:
22.4 lit

Question 16.
What is the number of 1 mole electron ?
Answer:
6.023 x 10²³

Question 17.
What is the mass of 1 gram-mole of water.
Answer:
18g

Question 18.
If a gas has vapour density 14, what will be its molecular weight ?
Answer:
Molecular weight of the gas = 2 × 14 = 28

Question 19.
How many number of molecules are present in 18g water ?
Answer:
18g water contains 6.023 × 10²³ molecules of water.

Question 20.
What is the atomicity of Argon gas ?
Answer:
The atomicity of Argon is 1.

Question 21.
What will be the volume of gas of 64g SO₂ at NTP ?
Answer:
22.4 lit

Question 22.
Calculate the number of molecules in 4.5 g water.
Answer:
The number of molecules present in 4.5 g water
\(=\frac{6.023 \times 10^{23} \times 4.5}{18}=1.505 \times 10^{23}\)

Question 23.
What will be the weight of 44.8 lit CO₂ at NTP ?
Answer:
The weight of 44.8 lit CO₂ at NTP

Question 24.
What are the number of electrons in 2 moles electron ?
Answer:
The number of electrons present in 2 moles electron
= 2 x 6.023 x 10²³ = 12.046 x 10²³

Question 25.
Which one is heavier : 2 gram-molecule ammonia and 2 gram molecule carbon dioxide ?
Answer:
2 gram-molecule ammonia = (2 x 17) g = 34g NH₃
2 gram-molecule carbon dioxide = (2 x 44)g = 88g CO₂
So, 2 gram-molecule carbon dioxide is heavier than 2 gram-molecule ammonia.

Question 26.
If the vapour density of any gas is 35.5, what will be its molecular weight ?
Answer:
The molecular weight of the gas will be = 2 x 35.5 = 71

Question 27.
One atomic mass unit (amu) is how much grams ?
Answer:
1.6603 x 10^-24g

Question 28.
What is the volume of 7g nitrogen at STP ?
Answer:
5.6L

Short Answer Type Question

Question 1.
State Avogadro’s hypothesis.
Answer:
Avogadro’s hypothesis : Under the same conditions of temperature and pressure, equal volume of all gases (both elementary and compound) contain equal number of molecules.

Question 2.
What is Gay Lussac’s law ?
Answer:
Gay Lussac’s law : If different gases react chemically under the same condition of temperature and pressure and if the product is gaseous then the reactants and the products maintain a simple ratio in their volumes.

Question 3.
What is Berzelius hypothesis ?
Answer:
Berzelius hypothesis : Under the same conditions of temperature and pressure, equal volumes of all gases contain the same number of atoms.

Question 4.
What is vapour density ?
Answer:
Vapour density : Vapour density of a gas is the ratio of the weight of a certain volume of the gas to the weight of the same volume of hydrogen under similar conditions of temperature and pressure.

Question 5.
What is molecule ?
Answer:
Molecule : The smallest particle of an element or compound which can exist in the free state is known as molecule.

Question 6.
What are the types of molecules ?
Answer:
There are two types of molecules :

• Elementary molecule
• Compound molecule

Elementary molecule : The molecule which is formed by one or more atoms of an element is known as elementary molecule, e.g. Na, K, B, C etc.
Compound molecule : The molecule which is formed by atoms of more than one element is known as compound molecule, e.g. H₂O, NH₃, HCl etc

Question 7.
What is molecular weight ?
Answer:
Molecular weight : The molecular weight of a substance is a number which represents how many times a molecule of the substance is heavier than \(\frac{1}{12}\)part of the weight of a C-12 isotope.

Question 8.
What is atomic weight ?
Answer:
Atomic weight : Atomic weight of an element
\(=\frac{\text { Weight of } 1 \text { atom of the element }}{\frac{1}{12} \text { of the weight of } 1 \text { atom of } \mathrm{C}^{12}}\)

Question 9.
What is gram atomic mass ?
Answer:
Gram atomic mass : When the atomic mass of an element is expressed in gram then the amount of element in gram is known as gram atomic mass or gram-atom of the element.

Question 10.
What is gram molecular mass ?
Answer:
Gram molecular mass : When the molecular mass of an element or compound in expressed in gram then that amount of element or the compound in gram is known as gram molecular mass or gram molecule or gram-mole of the element or the compound.

Question 11.
What is Molar volume ?
Answer:
Molar volume or gram-molar volume : The volume of a gaseous substance (element or compound) under a fixed temperature and pressure is known as molar volume or gram molar volume.

Question 12.
What is Avogadro’s number ?
Answer:
Avogadro’s number : One gram-molecule of any element or compound contains equal number of molecules. This number is known as Avogadro’s number.
It is denoted by N and its value is 6.023 × 10²³

Question 13.
What is atomicity ?
Answer:
Atomicity : Atomicity is. the number of atoms by which an elementary molecule is composed of.
e.g. atomicity of He = 1; atomicity of oxygen = 2; atomicity of ozone = 3; atomicity of phosphorous = 4

Question 14.
What do you mean by mole ?
Answer:
Mole : The ‘mole’ is regarded as the amount of the substance in grams which contains Avogadro’s number of elementary entities (such as molecules, atoms, ions) constituting the substance under investigation.

Question 15.
What is atomic mass unit (amu) ?
Answer:
Atomic mass unit (amu) : It is the quantity of mass equal to \(\frac{1}{12}th\) of the mass of carbon atom (c¹²).
1 amu = 1.6606 x 10^-24 g

Question 16.
What is the difference between molecular weight and actual weight of a molecule ?
Answer:
Distinction between molecular weight and actual weight of a molecule : Gram molecular weight is the weight of 6.023 × 10²³ molecules, while the actual weight is the weight of one molecule.
Weight of molecule = \(\frac{\text { Gram-molecular weight }}{6.023 \times 10^{23}}\)

Question 17.
Give statement of two important deductions of Avogadro’s law
Answer:
Two important deductions of Avogadro’s law :

• Molecules of common elementary gases like hydrogen, oxygen, nitrogen etc. are diatomic.
• The gram-molecular (or molar) volume of all gases under the same condition of temperature and pressure is the same and at STP, it is 22.4 litBroad answer type questions

Broad answer type questions 

Question 1.
Explain Berzelius hypothesis on the basis of Dalton’s atomic theory.
Answer:
Explanation of Berzelius hypothesis on the basis of Dalton’s atomic theory : Under the same conditions of temperature and pressure one volume of hydrogen and one volume of chlorine combine to form two volumes of hydrogen chloride gas. According to Berzelius hypothesis, suppose that under the same conditions of temperature and pressure each volume of a gas contains
n-number of atoms.
We can say,
n-number of hydrogen atoms + n-number of chlorine atoms
= 2n number of hydrogen chloride atoms
∴ 1 atom of hydrogen + 1 atom of chlorine
= 2 atoms of hydrogen chloride
∴ \(\frac{1}{2}\) atom of hydrogen + \(\frac{1}{2}\) atom of chlorine
= 1 atom of hydrogen chloride
Now, existence of \(\frac{1}{2}\) atom of hydrogen and chlorine is impossible according to Dalton’s atomic theory, so the hypothesis clashed with the basic concept of atomic theory. Hence, it was rejected by Dalton.

Question 2.
Explain Gay-Lussac’s law with the help of Avogadro’s hypothesis.
Answer:
Explanation of Gay-Lussac’s law with the help of Avogadro’s hypothesis :
From experimental result it has been found that under the same conditions of temperature and pressure one volume of hydrogen combines with one volume of chlorine to produce two volumes of hydrogen chloride gas.

Suppose, under the same conditions of temperature and pressure each volume of gas contains
n-number of molecules.
So,
n-number of hydrogen molecules + n-number of chlorine molecules
= 2n-number of hydrogen chloride molecules.
∴ 1 molecule of hydrogen + 1 molecule of chlorine
= 2 molecules of hydrogen chloride gas.
∴ \(\frac{1}{2}\) molecule of hydrogen + \(\frac{1}{2}\) molecule of chlorine
= 1 molecule of hydrogen chloride gas
This is to say,
1 molecule of hydrogen chloride must contain \(\frac{1}{2}\) molecule of hydrogen and \(\frac{1}{2}\) molecule of chlorine.
This fact does not go against Dalton’s atomic theory. Later on, it is known from Avogadro’s hypothesis that each molecule of hydrogen and chlorine contains two atoms.
So, 1 atom of hydrogen + 1 atom of chlorine
= 1 molecule of hydrogen chloride gas. in this way Dalton’s atomic theory and Gay-Lussac’s law were harmonized with the help of Avogadro’s hypothesis.

Question 3.
What are the application of Avogadro’s hypothesis ?
Answer:
Application of Avogadro’s hypothesis :

• Excepting the inert gases, all other active elementary gases are diatomic i.e. molecule containing two atoms.
• Molecular weight of any gas is twice its vapour density.
• Volume of a gram-mole of all elementary or compound gases is 22.4 lit at NTP.
• Number of atoms or molecules of all elementary or compound gases at their gram-atomic or gram-molecular weight is 6.023 x 10²³ at NTP.
• Molecular formula of a gaseous molecule can be deducted from its volumetric composition.
• Atomic weight of an element can be found out by the application of this hypothesis.

Question 4.
Prove that hydrogen molecule is diatomic in nature.
Answer:
Profit of diatomic nature of hydrogen molecule : From experimental result we know that under the same conditions of temperature and pressure one volume of hydrogen and one volume of chlorine combine chemically to form two volumes of hydrogen chloride gas.

Let us suppose that under the conditions of pressure and temperature at the time of experiment, one volume of hydrogen gas contains re-molecules. So, according to Avogadro’s hypothesis, under the same conditions of temperature and pressure one volume of chlorine contains remolecules and two volumes of hydrogen chloride gas contain 2n molecules.

We can write,
re-molecules of hydrogen + re-molecules of chlorine
= 2re molecules of hydrogen chloride gas.
∴ 1 molecule of hydrogen + 1-molecule of chlorine = 2 molecules of hydrogen chloride gas.
So, 1 molecule of hydrogen chloride
= \(\frac{1}{2}\) molecule of hydrogen + \(\frac{1}{2}\) molecule of chlorine
Now, according to Dalton’s atomic theory, an atom is indivisible.
So, one molecule of hydrogen chloride contains at least one atom of hydrogen and one atom of chlorine. This one atom of hydrogen comes from \(\frac{1}{2}\) molecule of hydrogen.
So one molecule of hydrogen must contain two atoms.

Question 5.
Deduce the relation of density and vapour density of a gas on the basis of hydrogen.
Answer:
Relation of density and vapour density of a gas on the basis of hydrogen
We know,

∴ Density of gas (d) = D x 0.089
(Density of hydrogen gas = 0 089)
∴ d = D x 0.089

Question 6.
Deduce the relation between vapour density and the molecular weight of a gas.
Answer:
Relation between vapour density and the molecular weight of a gas : From the definition of vapour density

Suppose, V volume of the gas contains ‘n’ molecules. Then by Avogadro’s hypotheis, V volume of hydrogen will also contain ‘n’ molecules of hydrogen.

i.e. Molecular weight of a gas 2 x its vapour density.

Question 7.
Gram-molecular volume of all gases at NTP is 224 litres – prove it.
Answer:
If the molecular mass of a gas is M, then we know that the limiting density of the gas 0.0898 × \(\frac{M}{2}\) g/litre
From the definition of limiting density it can be written that mass of 1 litre of a gas at NTP = 00898 x \(\frac{M}{2}\)g
So. at NTP, volume of 0.0898 × g gram of the gas = 1 litre
∴ At NTP volume of 1 gram of the gas = \(\frac{2}{M \times 0.0898}\) litre
∴ At NTP. volume of M g or 1 gram molecule of the gas
\(=\frac{2 \times M}{M \times 0.00898}=\frac{2}{0.0898}\) = 2227 litre
Atomic mass of hydrogen in oxygen scale is 2 016. in that case calculating in a similar way, we can obtain,
Volume of I gram-molecule gas at NTP\(=\frac{2.016}{0.0898}=22.4 \text { litre }\)
So, at NTP, volume of 1 gram-molecule of any gas is 224 litre.

Question 8.
What is the actual weight of a molecule ? What is the actual weight of an atom?
Answer:
Actual weight of a molecule:
Suppose, the molecular mass of any substance M
So, gram-molecular weight of the substance = Mg
We know that there are present 6.023 × 10² molecules in gram-
molecular weight.
∴ Weight of 6.023 x 10²³ molecules Mg
∴ Weight of 1 molecule = \(=\frac{M}{6.023 \times 10^{23}} \mathrm{~g}\)
Actual weight of an atom:
Suppose, the atomic mass of an element = A
So, gram-atomic weight of the element = Ag
We know that there are 6.023 x 10²³ atoms present in gram-atomic weight.
∴ Weight of 6.023 × 10²³ atoms = Ag
∴ Weight of 1 atom \(=\frac{A}{6.023 \times 10^{23}} \mathrm{~g}\)

Question 9.
What are the difference between molecular weight and acutal weight (mass) of a molecule?
Answer:
Difference between molecular weight and actual weight (mass) of a molecule :

Numerical Problem:

Problem 1.
If 200 ml of a gas at NTP weighs 0.27g and the normal density of hydrogen be 0.09 glut, calculate vapour density of the gas.
Answer:
At NTP, vapour density × 0.09 = Normal density
Now, 200 ml of the gas weighs 0.27 g

Problem 2.
Find the actual weight of 1 molecule water. (H = 1, O = 16)
Answer:
Molecular weight of water = (1 × 2 + 16) = 18
gram molecular weight of water = 18 g
So, 18g water contains = 6.023 × 10²³ molecules
Hence, actual weight of 1 water molecule
= \(\frac{18}{6.023 \times 10^{23}}\) = 2.988 × 10^-23g (approx)

Problem 3.
How many gram-moles of chlorine are present in 213 g chlorine and what is the volume of 6.023 x 10³⁰ hydrogen molecules at STP ? Given, atomic weight of chlorine is 35.5.
Answer:
Molecular weight of chlorine 355 × 2 = 71
∴ 71g chlorine = 1 gram-mole chlorine
∴ 213 g chlorine = \(\frac{1 \times 213}{71}\) gram-mole = 3 gram mole
At STP, 6.023 × 10²³ molecules occupy 224 lit
∴ At STP, 6.023 × 10³⁰ molecules occupy
\(=\frac{22.4 \times 6.023 \times 10^{30}}{6.023 \times 10^{23}} \mathrm{lit}\)
= 22.4 ×10⁷ lit

Problem 4.
Find the number of hydrogen and oxygen atoms in 425.6 lit of water vapour at NTP.
Answer:
Number of molecules in 22.4 lit water vapour at NTP
6.023 x 10²³
∴ Number of molecules in 425.6 lit water vapour at NTP
\(=\frac{6.023 \times 10^{23} \times 425 \cdot 6}{22 \cdot 4}=114437 \times 10^{20}\)
Now, 1 molecule of water contains number of hydrogen atom = 2
∴ In 114437 x 10²⁰ molecules of water contains number of hydrogen atom
= 2288874 × 10²⁰
Again, 1 water molecule contains 1 oxygen atom
∴ 114437 x 10²⁰ water molecules contain 114437× 10²⁰ oxygen atom.

Problem 5.
What is the weight of one gram oxygen and how many number of oxygen atoms are there ?
Answer:
One gram-atom of oxygen = 16 g of oxygen
From Avogadro’s hypothesis,
32 g of oxygen contains = 6.023 × 10²³ molecules
∴ 16 g of oxygen contains  \(=\frac{6023 \times 10^{23}}{32} \times 16\) molecules
= 3.0115 × 10²³ molecules
Again oxygen is a diatomic element,
Hence, it contains 2 × 3.0115 x 10²³ atoms = 6.023 × 10²³ atoms.

Problem 6.
Calculate the number of mole in 9 grams of water and calculate the number of hydrogen atoms present in it.
Answer:
18 g water contain = 6.023 × 10²³ molecules
9 g water contain \(\frac{6.023 \times 10^{23}}{2}\) = 3.0115 x 10²³ molecules of water
One molecule of water contains 2 hydrogen atoms. So the number of hydrogen atoms =
(3.0115 × 10²³) x 2 = 6.023 × 10²³

Problem 7.
250 cm³ of a gas at NTP weigh 0.7924. What is the molecular weight of the gas ?
Answer:
250 cm³ of the gas at NTP weighs = 0.7924 g
∴ 22400 of the gas at NTP weighs
= \(\frac{0.7924 \times 22400}{250} \mathrm{~g}\)
= 70 999 g
∴ Gram-molecule of the gas = 70.999 g
∴ Molecular weight of the gas = 70.999

Problem 8.
The molecular weight of a gaseous substance is 200. What is the volume of 5 g of the substance at NTP?
Answer:
The molecular weight of the substance = 200
∴ 1 gram-molecule of the substance = 200 g
Volume of 1 gram-molecule or 200 g of the substance at NTP = 224 lit.
Volume of 5 g of the substance at NTP = \(\frac{22.4 \times 5}{200}=0.56 \mathrm{lit}\)

Problem 9.
Calculate the number of molecules in a drop of water weighing 0.05 g (H = 1, 0 = 16)
Answer:
Molecular mass of water = (2 × 1 + 1 × 16) = 18
∴ Gram-molecular mass of water = 18 g
Now, 18 g of water contains 6.023 × 10²³ molecule
∴ 0.05 g of water contains
\(\frac{6.023 \times 10^{23}}{18} \times 0.05\) molecule
= 1.672 × 10²¹

Problem 10.
Calculate the total number of electrons present in 1.6g of methane (CH₄). [H = 1, C = 12]
Answer:
Molecular mass of methane = 16 g
16 g of methane 1 mole
∴ 16 g of methane = 0.1 mole
Now, 1 mole of methane = 6.023 × 10²³ molecules
∴ 0.1 mole of methane = 6.023 × 10²² molecules
We know that one atom of carbon contains six electrons and one atom of hydrogen contains one electron.
∴ One molecule of CH₄ contains = 6 + 1 × 4 = 10 electrons
∴ 6.023 × 10²² molecules contain = 10 x 6.023 x 10²²
= 6.023 × 10²³ electrons

Problem 11.
A tetra-atomic gas occupies 1.4 lit at 0°C and 76cm pressure. Find the number of atoms in the gas.
Answer:
At 0°C and 760 cm pressure, number of molecules present
= 224 lit of the gas.
= 6.023 × 10²³
= 6.023 × 10²³ x 4 atoms
Number of atoms present in 1.4 lit
= \(\frac{1 \cdot 4}{22 \cdot 4}\) x 6.023× 10²³ × 4 = 1.505 x10²³

Problem 12.
Calclulate the molecular mass of CO₂
Answer:
CO₂ molecule contains one C-atom and two oxygen atoms.
Molecular mass of CO₂= (1 x 12 + 2 x 6) = 44U

Problem 13.
Calculate the formula of unit mass of MgCl₂
Answer:
MgCl₂ molecule contains one Mg-atom and two Cl-atoms.
∴ Formula of unit mass of MgCl₂
= 24+2 x 35.5 = 24+71=95U

Problem 14.
Calculate the following :
(i) 1 milimole of NH₃
(ii) 3.011 x 10²³ number of \({ }_8^{16} \mathrm{O}\)
Answer:
(i) 1 milimole = 10 mole.
Mass = molar mass x no. of moles
= 17 × 10^-3= 0.017g

(ii) Mass = number of moles × atomic mass
\(=\frac{3.011 \times 10^{23}}{6.022 \times 110^{23}} \times 16=8 \mathrm{~g}\)

Example 15 :
Calculate the number of particles in each of the following :
(i) 35.5g of Na atoms.
(ii) 11g of CO₂
(iii) 0.2 mole of 2C atoms.
Answer:
(i) the number of atoms
\(=\frac{\text { given mass }}{\text { molar mass }}\)×Avogadro’s number
\(\frac{35 \cdot 5}{23}\)

(ii) Number of Melecules of CO₂

(iii) No of ¹²C atoms
= No of moles of particle x Avogadro’s number.
= 0.2 × 6.022 × 10²³ = 1.2044 × 10²³

Problem 16.
Calculate the mass of 250 molecules of sodium chloride.
Answer:
Molecular mass of NaCl
= 1 × 23 + 1 × 35.5 = 58.5
1 mole of NaCl = 58.5g NaCl
6.022 × 10²³ molecules of NaCl have mass 58.5g
250 molecules of NaCl have mass = \(\frac{58 \cdot 5 \times 250}{6.022 \times 10^{23}}\) = 2.428 ×10^-23g.

Problem 17.
What is the actual weight of one molecule of H₂ and one atom of hydrogen ? The gram molecule of H₂ is 2.016g.
Answer:
6.022 x 10^-23 molecules of hydrogen weigh 2.016g.
1 molecule of hydrogen weigh \(\frac{2 \cdot 016}{6 \cdot 022 \times 10^{23}} \mathrm{~g}\) = 3.34 × 10^-24g
Hydrogen is diatomic i.e. one molecule of hydrogen consists of two atoms. Therefore, one atom of hydrogen weighs
\(=\frac{3.34 \times 10^{-24}}{2}\) = 1.67 × 10^-24g

Problem 18.
How many moles or milimoles are present in 0.49 g of H₂SO₄? [H = 1, S=32, O=16]
Answer:
Molecular mass of H₂SO₄
=2 x 1+1 x 32+4  x 16=98
1g mole of H₂SO₄ is equal to 1 mole of H₂SO₄
∴ 0.49g of H₂SO₄ is equal to \(\frac{0.49}{98}\) = 0.005 mole
So. no of millimoles
= 0.005 × 103 = 5 millimole
[Since 1 mole 10³ millimole]

Comprehensive WBBSE Class 9 Physical Science Notes Chapter 1 Measurement can help students make connections between concepts.

Measurement Class 9 WBBSE Notes

Science : This word originated from Latin word ‘Scientia’ meaning ‘to know’. Thus knowledge acquired by man through systemetic observations and experiments is called Science.

Physical Quantity : A physical quantity is any measurable quantity of an object or an event. Example : mass, length, time, weight, density etc.

Types of physical quantity :

• Scalar quantity
• Vector quantity

Scalar quantity : The quantity which has only magnitude but no direction is called scalar quantity.
Example : mass, length, time etc.

Vector quantity : The quantity which has both magnitude and direction is called vector quantity.
Example : weight of matter, velocity, acceleration, force etc.

Unit : In measuring any physical quantity, some convenient and definite quantity of it is taken as the standard and in terms of this standard the physical quantity is measured. This standard is called a unit.

Importance of units :

• Measurement of any physical quantity without unit is meaningless. Because we cannot have any idea about a physical quantity with its magnitude only.
• Unit establishes relation between different measures of same quantity.

Characteristics of units :

1. The unit should be well defined.
2. The unit should be of suitable size.
3. The unit should be easily reproducible.
4. The unit should be imperishable.
5. The unit should not change with time or with physical conditions like pressure, temperature etc.

Fundamental unit : This units of physical quantities which are independent of each other and from which other units can be derived are called the fundamental units.
Example : units of length, mass, time etc.

Derived unit : The units of physical quantity which are derived with the help of one or more than one fundamental units are called derived units.

Example : The unit of area is obtained by using the unit of length twice. Similarly units of speed, force, work etc. are all derived units.

Different systems of fundamental units :

1. CGS system : The word ‘C’ is the unit of length-Centimeter, ‘G’ is the unit of mass gram, ‘S’ is the unit of time-Second.
2. EPS system : This is the British system of units in which units of length, mass and time are foot, pound and second respectively.
3. MKS system : This is basically practical systems of units in which units of length, mass and time are metre, kilogram and second respectively.

SI units : In 1960, an international system of units was adopted to have a consistent system of units. This system of units is known as SI units.

Types of SI units :

• Fundamental SI units
• Derived SI units

Fundamental SI Units :

Derived SI units :

Some important rules for writing SI units :

• The symbol of the units are always written in Roman small letter. Example : {kg}, {m}, {s} etc.
• If a unit is named after a person, the symbol is written in capital letter. Example : Newton-N, Ampere-A, Volt-V etc.
• The symbols of the units are always used in singular form. Example : Mass- 5 kg not 5 kgs.
• When temperature is expressed in Kelvin scale, (0°) degree sign is not used i.e. we write 273 K, not 273°k
• Full stop, comma, etc. are not written after the symbols of the units. (e.g. we should write cm and not cm. or cm,)
• The multiplication of two units are written as symbols of unit sequencially, e.g. kgm.
• units like metre per second is written as m/s or ms^-1 or in case of Joule per kelvin per mole is written as JK^-1 mol^-1 or J/K mol but not J/K/mol.

Definitions of units of SI :

Metre : The distance between the two marks made on a platinumiridium bar maintained at 0°C temperature preserved at International Bureau of Weights and Measures (Bureau international des poids et measures) at Sevres near Paris considered as one metre (symbol : m).

Modern definition of metre : The standard metre is exactly equal to 1650763.73 wavelengths in vacuum, of the radiation from Krypton isotope of mass number 86.

Kilogram : Mass of a solid cylinder made of platinum-iridium and preserved at the standard office at Sevres near Paris, is called one kilogram (kg).

Modern definition of second : One second is the duration of 9192631770 periods of radiation corresponding to unperturbed transition between the two hyperfine levels of the ground state of \({ }^{193} {C}_{{s}}\) atom.

Units of volume :

Volume : It is defined to be the space occupied by a substance.
Solids have three dimensions, i.e. length, breadth and height and so the unit of volume is CGS system in cm × cm × cm = cm³.
Unit of volume in SI system is m³.

Volume of some solid figures :

1. Volume of a rectangle solid = length × breadth × height.
2. Volume of a cube = (length)³.
3. Volume of a sphere = \(\frac{4}{3}\) πr³
4. Volume of a cone = \(\frac{1}{3}\) πr² h
5. Volume of a right circular cylinder = πr² h

[r = radius, π = \(\frac{22}{7}\), h = the perpendicular height]

Unit of volume of liquid :

Litre (L)= Volume of 1 kg of pure water at 4°C or 277 K is called a litre. This is not an SI unit. 1 L = 1000 ml = 1000 cm⁹ = 10 dm³

Volume of 1 kg of water is 1000 ml.
Hence, volume of 1000 ml (cc) of pure water at 277 K is called a litre.
It is the unit of liquid in CGS system.
1 ml = 1 cc

Advantage of increased volume of water on solidification : In the winter season in the cold country water solidifies to ice and the volume of water increased on solidification. The density of ice being less than water it floats on the upper surface of water. Again by costing water reaches at the temperature of 4°C and it attains its maximum density and thereby it cannot come on the upper surface of water below the surface of ice and remain at 4°C. So aqua life is possible in winter.

Measurement of length : The measurement of lengths are of two types :

• direct method
• indirect method

Length measuring device in direct method :

• a metre scale : for 10-3 m-102 m length
• a Vernier – Callipers : for distances upto 10-4m
• a screw gauge or a spherometer : for distances upto 10-5 m.

Length measuring device in indirect method for large distance :

Parallax method : The change in the position of an object with respect to the background, when seen from two different positions is known as parallax. The distance between the two positions of observation is called the basis.
Size of an astronomical object : The size of an astronomical object, such as the moon can be measured with the help of an astronomical telescope.

Indirect methods for determination of very small length : The optical microscopes, working on visible light of wavelengths ranging from 4 × 10^-7 m to 8 × 10^-7 m, cannot be used to measure the sizes of molecules (10^-8 m. to .10^-16 m).
An electron microscope is usually used for this purpose.

Wavelength of radiations are expressed in angstrom (A°).
1 A°=10^-8 cm = 10^-10m

Some commonly used prefixes in CGS and SI system :

Measurement of Mass :

The mass of a body is the quantity of matter contained in it. The range of mass varies from that of electron having mass of the order of 10-30 kg to observable universe with masses of the order of 1055 kg}.

Mass measuring device :

1. The mass of ordinary objects-common balance
2. The large masses of planets etc-gravitational methods are adopted.
3. For measuring of very small masses of atomic or sub-atomic particles – mass spectrograph.

Few important points related to common balance :

If the mass of right-hand side pan and left-hand side pan in balance are unequal but same in length then the actual mass of substance

m = \(\frac{m_1+m_2}{2}\)
(where, m₁ and m₂ are mass of substance in two different pans)

If the length of balance is unequal but the pans are same in mass, then the actual mass of substance
m = \(\sqrt{m_1 m_2}\) (where m₁ and m₂ are mass in two pans)

Characteristics of Good balance :

• The pillar should be vertical, i.e. the balance beam horizontal with the help of levelling screws and plumb line.
• The balance should be correct.
• The balance should be rigid.
• The balance should be stable.

Different portion of comman balance :

• balance beam
• stirrups
• pointer
• lever
• scale pans
• plumb line
• glass case

Weight box : A wooden box for holding weights are called weight box is supplied each balance.
It may be noted that in a weight box weights one in the ratio 1 : 2 : 2 : 5.
A common balance can measure a minimum weight of 5 mg accurately. The upper range is about 200 g.

Spring balance: The weight of a body is the force with which is attracted by the earth towards its centre. The weight of a body is measured by spring balance.

Measurement of time : The range of time interval of events varies from as small as of the order 10-24 s} for life span of most unstable particle to as large as of the order of 1017 s} for age of universe.

For measurement of time interval, we need a clock based on any phenomenon that repeats itself regularly.
Commonly used clocks and watches are based on spring, pendulum etc.

Time measuring device :

Stop watch : It is used to measure the interval of incident starting and ending. This watch can measure a time interval of one-tenth of a second accurately. This watch is used in sports, scientific work etc. for measuring time.

Sand watch : This consists of two conical shaped glass vessels joined with each other with a narrow opening in the middle. The upper vessel contains a measured quantity of dry sand and it takes a definite interval of time for the sand to pass from the upper vessel to the lower one. The bottle is inverted for reuse when the upper vessel becomes empty.

Sundial : Sun rays are used in this time-measuring device. It consists of a horizontal circular disc with a thin triangular plate of metal fixed at its centre and pointing in the north-south direction. This triangular plate serves as an obstacle to the rays of the sun and casts shadow on the circular disc on the other side. The circular disc has graduations from 1 to 12 like that on a clock. Any particular time of the day, is indicated by the position of the shadow.

The sundial can be used only on a sunny day and not at night or on a cloudy day.

The pendulum clock : In this clock a simple pendulum is used. At a particular place, a pendulum of a given length takes fixed time to complete one oscillation i.e. starting from one extreme position and coming back to that position. This time is called the time period of the pendulum. A seconds pendulum is that whose time period is 2 seconds. In a pendulum clock such a seconds pendulum is used.

Atomic or caesium clock : These clocks are based on periodic vibrations produced in a caesium atom. These clocks are very accurate and in a year such a clock lose or gain not more than 3s.

Electric oscillators : The A.C. main electric supply has a frequency of 50 Hz. The synchronous rotations of an A.C. motor can be used for having a time scale.

Electronic oscillators : Vacuum tubes and transistors can be used for producing electro-magnetic waves of high frequencies and their small time periods of oscillations are used small time intervals accurately.

Quartz crystal clock : A quartz crystal shows piezo-electric effect. When fluctuating pressure is applied across one pair of faces of a crystal, an oscillating emf is developed across another pair perpendicular faces. These oscillations may be used for measuring time. It has an accuracy of 1 second in every 10^-9second.

Decay of elementary particles : Many unstable elementary particles decay in time interval as short as 10^-10 second to 10^-24 second. Thus studying their decay very small time intervals can be measured.
Radioactive dating : This technique is used for measuring long time interval of the order of 10^-17 second.

Accuracy : The closeness of the measured value to the true value of the physical quantity is known as the accuracy of the measurement.

Precision : It means the extent or limit to which the measurement of a physical quantity is done.
Errors in measurement : The error in the measurement of a physical quantity is defined as the difference between the true value and the measured value of the physical quantity.

Types of errors :

• Systematic errors
• Random errors

Systematic errors : The error which occurs according to a definite pattern is known as systematic error.

Types of systematic errors :

1. Instrumental errors : The errors caused due to defective instrument are called instrumental errors.
2. Personal errors : The errors in the measurement of a physical quantity due to limitations or carelessness of the person experimenting are known as personal errors.
3. Natural errors : The errors due to the change in the conditions of the environment (temperature, pressure etc.) are known as natural errors.

Random error : Random or chance errors are due to unknown causes. These errors cannot be controlled by the observer and are not constant in magnitude. They may be positive or negative.
The errors are expressed in different ways :

Absolute error: The absolute error of a given value of physical quantity is the difference between mean value of the physical quantity and the observed value under consideration.
Absolute error of the i-th observation

the mean value of the measured quantity,
x_i = the value of the measured quantity in the i th observation.]

Mean absolute error : The arithmetic mean or average of all absolute errors of the measurements is called the mean absolute error.
Mean absolute error

Relative error : The ratio of the absolute error to the physical quantity is called the relative error.
Relative error = Sk = \(\frac{\Delta \bar{x}}{x}\)

Percentage error = Relative error × 100%
= \(\frac{\Delta \bar{x}}{x}\) × 100%
Accuracy : The closeness of the measured value to the true value of the physical quantity is known as the accuracy of the measurement.

The degree of accuracy of any measurements depends upon :

the accuracy of the measuring device used.
Precision : It means the extent or limit to which the measurement of a physical quantity is done.

Significant figures : It is in the reported result of a measurement of a quantity is the number of digits that are known with certainty plus one that is uncertain, beginning with the first non-zero digit.

Rounding off : The observed results of various measurements may have different precisions. Thus, the results obtained at various stages of calculation are to be rounded off because the final result cannot be more precised than that of the least precised measurement.

Precautions in the measurement of measuring devices :

Ordinary scale : We often use ordinary scale for measuring length. It is usually a thin rectangular bar of box-wood, metal or plastic.

Precautions in measurement :

• The scale is to be placed on the straight line in such a way that the graduations of the scale be perpendicular to the straight line. By doing this the error in the reading due to thickness of the scale is avoided.
• It is better not to use any end of the scale, for that edge may be broken.
• The length of a straight line should be measured by different parts of the scale and average length should be determined.
• By doing this, the error in the reading due to irregularities in the graduations in the different parts of the scale, if any, is eliminated.
• IN.B. In making an ordinary scale wood or plastic is usually used instead of metal. With the change in atmospheric temperature metal scales changes in length.]

Measuring cylinder : A measuring cylinder is used for measuring the volume of a liquid. It is a glass cylindrical jar graduated in millilitre. The internal volumes of the jar are marked with horizontal marks, the reading starting from the bottom of the jar. The liquid whose volume is to be determined, is poured, into the jar. The volume of the liquid is obtained from the reading corresponding to the level of the liquid in the jar.

Precautions in measurement :

1. During determination of volume of a liquid, it should be noted that free surface of the liquid in the jar is always either concave or convex.
2. In either case (concave or convex), the reading is to be taken along a line tangential to the curved surface of the liquid. The reading is taken avoiding parallax error.

Common balance : In the laboratory we measure the mass of a body usually with a common balance. Actually we find the mass of the body by comparing its mass with some standard weights. Common balance is much more sensitive and masses of even small objects can be determined accurately with its help.

Precautions in measurement :

1. Body whose mass is to be determined must be dry and at room temperature.
2. The weights must be held with a pair of forceps.
3. The balance beam should be on the beam support when weights are being placed on or taken away from the scale pan, otherwise the beam may topple down.

Dimensions : Dimensions of a physical quantity give the relation of its unit with the units from which it is derived.

The dimensions of a physical quantity are expressed as the powers to which the fundamental units of masses, length and time are raised to obtain the derived unit of the quantity.

Fundamental units from which the unit of a physical quantity are derived are each expressed in capital letters.

For example, the length is denoted by [L], unit of mass by [M], unit of time [T] etc. The dimensions are always written within square bracket.
w Dimensonal formula and equation : The dimensional formula of a physical quantity is an expression which gives the fundamental units on which the physical quantity depends and the nature of the dependence.

Thus, the dimensional formula of velocity is [M° L¹ T^-1]. If we represent velocity by [v], then [v]=[ M°L¹T^-1] is called dimensional equation of velocity. So, when a physical quantity is equated to its dimensional formula, we get the dimensional equation of the physical quantity.

Dimensional formulae of some physical quantities :

Different types of variables and constants :

Dimensional constants : The quantities which have dimensions but of constant value are called dimensional constants.
Example : Gravitational constant, Planck’s constant.

Dimensionless constants : The constant quantities having no dimensions are called dimensionless constants.
Example : pure numbers 1,2,3 …., π, e ( = 2 .718).

Dimensional variables : The quantities which have dimensions but do not have any fixed value are called dimensional variables.
Example: volume, velocity, force etc.

Dimensionless variables : The quantities which have neither dimensions nor any constant value are called dimensionless variables.
Example : angle, specific gravity, strain etc.

Use of dimensional analysis :

The dimensional analysis have following applications :

• To check the correctness of a physical equation.
• To derive the relation between different physical quantities involved in a physical phenomena.
• To convert from one system of units to another.

Limitations of dimensional analysis :

1. This method fails to determine the dimensionless constants in the formula.
2. If a physical quantity depends on more than three factors, having dimensions, the formula cannot be determined.
3. This method cannot be used to derive a relation if it involves trigonometric or exponential or log functions.
4. This method fails to derive an exact form of a relation when it consists of more than one part on any one side.
5. It gives no information whether a physical quantity is scalar or vector.
6. Even when dimensions are given, the physical quantity may not be unique, as many physical quantities have the same dimensions.

Some uses of different measuring devices :

1. Determination of areas of an irregularly shaped sheet of metal or paper: This is done with the help of a graph paper.
2. Determination of the length of a curved line: This is done by using a thread and a linear scale.
3. Determination of thickness of a sheet of a thin paper : This is indirectly done with the help of linear scale.
4. Determination of volume of an irregularly shaped solid: This is done by using a measuring cylinder.
5. Determination of rate of fall of water from a tap : This is done by using measuring cylinder and stopwatch.

Comprehensive WBBSE Class 9 Physical Science Notes Chapter 4 Matter: Atomic Structure, Physical and Chemical Properties of Matter can help students make connections between concepts.

Matter: Atomic Structure, Physical and Chemical Properties of Matter Class 9 WBBSE Notes

Indian Philosopher Kanad named the basic smallest particles of matter as Kana.

Greek Philosopher Demokritos called the smallest particle as Adornos’ meaning indivisible.

John Dalton developed the idea of ultimate unit of matter which is known as atom.

Cathode Rays : It consists of negatively charged material particles called electrons. Cathode rays are produced from cathode surface when a gas at low pressure is subjected to electric discharge.

Electrons : Fundamental sub-atomic particles carrying negative charge (1.602 × 10^-19 coulombs) and having mass 91 × 10^-31 kg, discovered by J. J. Thomson.

Proton : A sub-atomic positively charged particle, having charge 1.602 × 10^-19 coulombs and mass 1.67 × 10^-27 kg. Mass of a proton is nearly 1837 times higher than the mass of an electron.

Alpha particle : He²⁺ ions or helium nuclei.

Rutherford’s Experiment : Led to the discovery of nucleus. Radius of nucleus (~ 10^-15 m) is very small as compared with radius of an atom (~ 100^-10m).

Neutrons : Sub-atomic neutral particles having mass 1.675 x 10^-27 kg, discovered by James Chadwick.

Nucleons : The constituent particles of the nucleus i.e., protons and neutrons together are called nucleons.

The mass of an electron is \(\frac{1}{1837}\) times the mass of a proton.

The mass of an electron is \(\frac{1}{1839}\)  times the mass of a neutron.

Rutherford’s model of the atom proposed that a very tiny nucleus is present inside the atom and electrons revolve round this nucleus. The stability of the atom could not be explained by this model.

Neils Bohr proposed that electrons revolve round the nucleus in a limited number of orbits, called permissible orbits. The orbits are also called energy level.

Energy levels are designated by letters K, L, M, N, O, P, Q.

Elements with their outermost shell completely filled are chemically inert.

The outermost shell of an atom is called valence shell and the electrons present in the valence shell are called valence electrons.

The maximum number of electrons that can be accommodated in a shell is given by 2n2, where
n = shell number.

Atomic number (Z)= Number of protons
= Number of electrons
– Mass number (A) = Number of protons + Number of neutrons

Isotopes : Atoms of same element having same atomic number but different mass number.

Isobars : Atoms of different elements having same mass number but different atomic number.

Isotones: Atoms of different elements having same neutron number but different atomic and mass number.

Nuclear force : In the nucleus, continuous transformation of the closely packed neutrons and protons take place i.e. protons transform into neutrons and vice versa. As a result of this transformation a strong attractive force acts within the nucleus. This force is known as nuclear force.

Ion : A charged atom or radical is called an ion.

Cation : Positively charged atom or radical is called a cation.

Anion : Negatively charged atom or radical is called an anion.

Frequency (v) of the radiation emitted when an electron jumps from an orbit having energy E₁ to an orbit having energy E₁ is given as \(v=\frac{E_2-E_1}{h}\)

Electronic arrangement of elements:

Each element has its own characteristic emission spectrum.

Series of lines in the line spectrum of hydrogen atom :

Orbital: Region of space around the nucleus where probability of finding an electron is maximum.

Quantum Mechanics: It is a theoretical science that takes into account the dual nature of electron.

Degenerate orbitals: The orbitals having equal energy.

Ground state: The lowest energy state of an atom.

Excited state: An atom is said to be in excited state if some lower energy orbital is vacant and the electron is present in higher energy orbital.

Electronic configuration : Distribution of electrons among various orbitals in an atom.

Quantum numbers : A set of four number used to specify energy, size, shape, orientation of the electron orbital and spin of the electron.

Principal Quantum number (n) : Tells us about energy of the electron and size of the orbital.

Azimuthai Quantum number (1): Tells us about the shape of the orbital and orbital angular momentum.

Magnetic Quantum number (m): Tells us about orientations of the electron cloud in a sub-shell.

Spin Quantum number (s) : Tells us about spin of the electron.

s-orbitais are spherically symmetrical.

Shape of s-orbitais:

p-orbitals are dumb-bell shaped:

The five d-orbitais are : d_xy d_yz d_xz dx² – y² and dz²

Angular momentum \(m v r=\frac{n h}{2 \pi}\)
[m mass of the electron, u = velocity of the electron, r = radius of the orbit,
h = Planck’s constant, n = Principal Quantum number]

de-Brogue equation:
\(\lambda=\frac{h}{m v}\) A = [λ = wavelength, m = mass of the electron, υ = velocity of the electron, h = Planck’s constant]

Heisenberg’s uncertainty principle:
\(\Delta x \cdot \Delta p \geq \frac{h}{4 \pi}\)
[Δx = uncertainty in position,
Δp = uncertainty in momentum
h = Plancks constant]

Schrodinger’s wave equation:

[m = mass of an electron, V potential energy, E = total energy, Ψ = wave function of electron (eigen function

Aufbau Prinuiple : The orbitais are filled in the order of increasing energy,starting with the orbital of lowest energy.

Pauli’s Exclusion Principle: No two electrons in an atom can have same set of ail the four quantum numbers.

Hund’s Rule of Maximum Spin Multiplicity: The pairing of electrons in orbitais of a particular sub-shell cannot take place until all the orbitais of the sub-shell are singly occupied.

Nuclear isomer : Atoms (nuclides) which have the same atomic number and mass number but have different radioactive properties are called nuclear isomers.
Example:

Isodiapher : Atoms in which the difference between the number of neutrons and the number of protons is saine are termed as isodiapher.

Example:

Isoster: Molecules or ions with same number of atoms and also the same number of electrons are said to form isosteric group or more simply isosters.
Example : in N₂O and CO₂, number of atoms 3 and numbers of electrons = 20. So, these are isosters.

Gay Lussac’. Law : Under the same conditions of temperature and pressure, when two or more gases combine together. they do so in simple ratio by volumes, and the volumes of the products if gaseous also bear a simple ratio to the volume of the reacting gases.

Berzelius hypothesis: Under the same conditions of temperature and pressure, equal volume of all gases contain the same number of atoms.

Avogadros hypothesis : Under the same conditions of temperature and pressure equal volume of all gases (both elementary and compound) contain equal number of molecules.

Molecule : The smallest particle of an element or compound which can exist in the free state is known as molecule.

Elementary molecule : The molecule which is formed by one or more atoms of an element is known as elementary molecule. Example : Hydrogen (H₂), Oxygen (O₂) etc.

Compound molecule : The molecule which is formed by atoms of more than one element is known as compound molecule. Example: Water (H₂O), Carbon dioxide (CO₂) etc.

Vapour density: Vapour density of a gas is the ratio of the weight of a certain volume of the gas to the weight of same volume of hydrogen under similar conditions of temperature and pressure.

Molecular mass =

Atomic mass = \(\frac{\text { mass of } 1 \text { atom of an element }}{\frac{1}{12} \times \text { mass of } 1 \text { carbon }\left({ }^{12} \mathrm{C}\right) \text { atom }}\)

Gram atomic mass : When the atomic mass of an element is expressed in gram then the amount of element in gram is known as gram atomic mass of the element.

Gram molecular mass : When the molecular mass of an element or compound is expressed in gram then that amount of element or the compound in gram is known as gram molecular mass of the element or the compound.

Molar volume : The volume of a gaseous substance (element or compound) under a fixed temperature and pressure is known as molar volume of the gas.

Avogadro’s number: One gram-molecule of any element or compound containing equal number of molecules (N_A 6.022 × 10²³) is known as Avogadro’s number.

R. A. Millikan determined the value of Avogadro’s number by ‘oil drop experiment’ in 1913.

Modern definition of Avogadro’s number : The number of atoms present in exactly 12g of carbon (12C) is designated as Avogadro’s number.

Avogadro’s constant : ‘Avogadro’s number/mole’ is called Avogadro’s constant
i.e. 6.022 × 10²³ Avogadro’s number has no unit. But Avogadro’s constant has the unit mol’. It is a universal constant.

Atomicity : It is the number of atoms by which an elementary molecule is composed of.

Atomic mass unit (amu) : It is an unit used for measusing the atomic mass of an atom.
1 amu = 16606 ×10^-24g

Average atomic mass
\(\frac{\sum \text { natural abundance of isotope }(\%) \times \text { its atomic mass }}{100}\)

Mole (SI System) : A mole is the amount of a substance which contains as many entities (atoms, molecules, ion or any other particle) as there are atoms in exactly 0012 kg (or 12g) of carbon-12 isotope.

Solution : A solution (or true solution) is a uniformly homogeneous mixture of two or more substances whose relative proportions may be varied upto a certain limit.

Solute : The substances which are present in smaller quantities and gets dissolved are called solutes.

Solvent : The medium in which the solutes are uniformly dispersed through dissolution is known as the solvent.  solute + solvent = solution

Homogeneous mixture : A homogeneous mixture is one in which the constituent substances are so intimately mixed that even a very close examination cannot distinguish any surface of separation between them.

Heterogeneous mixture : A heterogenous mixture on the other hand is one in which the particles of the constituent substances of a mixture are distinctly distinguishable.

Some useful solvents : Water is a versatile liquid solvent of organic and inorganic substances, carbon disuiphide is the solvent of sulphur, oxalic acid is the solvent of rust, benzene is the solvent of rubber etc.

Types of solution:

(a) Solution of solid in solids : Solution of one metal in another gives the alloys like brass, bronze etc.

(b) Solution of liquids in solids : Mercury (liquid metal) can remain dissolved in solid metal like gold, silver, sodium etc. to form amalgams.

(c) Solution of gases in solids : Hydrogen gets dissolved in spongy palladium and in some similar metals. This phenomenon is called occlusion.

(d) Solution of solids less in liquids : Aqueous solution of common salt (NaCl), sugar, copper sulphate are examples.

(e) Solution of liquids in liquids : A mixture of water and alcohol, a mixture of benzene and xylene are examples.

(f) Solution of gases in liquids: Ordinary water contains air dissolved in it. Soda water contains carbon dioxide under pressure. Dissolved oxygen in water allows aquatic life to survive.

(g) Solution of solids in gases : e.g. smoke, where solvent is air and solute is carbon and dust particles dissolved in air.

(h) Solution of liquids in gases: e.g. fog.

(j) Solution of gases in gases : e.g. air, mainly a mixture of nitrogen and oxygen.

Classification based on concentration:

• saturated solution
• unsaturated solution
• supersaturated solution

Suspension : The particle in suspension state is 10^-4cm in diameter.

True solution : The particle in true solution is 10-8 cm in diameter.

Sol : It is a colloidal solution in which a solid is dispersed in a liquid (e.g.paints).

Emulsion : It is a colloidal solution in which a liquid is dispersed in another liquid (e.g. milk).

Gel : It is a colloidal solution in which a liquid is dispersed in a solid (e.g.fruit jellies, cheese etc.)

Aerosol : It is a colloidal solution in which a solid or a liquid is dispersed in a gas (e.g. fog. smoke, clouds etc.)

Lyophilic colloids : These are the substances which pass readily into the colloidal state whenever mixed with a suitable solvent (e.g. protein, starch etc.)

Lyophobic colloids : These are the substances which do not yield colloidal solutions on mere shaking with a liquid (e.g. gold, silver. Fe(OH)⁴, As).

Positive sols : These are the sols which carry positive charge on the dispersed phase particles. (e.g. sols of Fe(OH), Al(OH)₃ etc.)

Negative sols : These are the sols which carry negative charge on their particles (e.g. sols of Cu, Ag, Au, As₂S₃ etc.)

Multi-molecular colloids :These are the colloids in which the individual particles consist of aggregates of atoms of small molecules having molecular sizeless than 10^-7cm in diameter. (e.g. sols. of gold atoms, platinum sol.)

Macro-molecular colloids : These are the colloids in which the size of the particles of the dispersed phase are of the order of colloidal dimensions (e.g. sol of starch, cellulose)

Peptisation : It is the process of converting precipitates into colloidal state by adding small amount of a suitable electrolyte.

Dialysis : It is the process of separating substances in colloidal state from those present in ionic states with the help of a semipermeable membrane

Tyndal I effect : It is the process of scattering of light from the surface of colloidal particles and in this process a beam of light passed through a colloidalsolution becomes visible as a bright streak.

Brownian motion : This is ceaaeless erratic irregular and random motion of colloidal particles suspended in a dispersion medium.

Elcctrophoresis : It is the movement of charged colloidal particles, under the influence of an electric field, towards the oppositely charged electrodes.

Coagulation : It is the phenomenon of change of colloidal state to suspended state. It can be brought about by adding an electrolyte to a colloidal solution.

Hardy schuize rule : It states that greater the valency of the coagulating ion added, the greater is its power to coagulate.

Gold number : ¡t is the weight in milligrams of a protective colloid which prevents the coagulation of 10 mL of a gven gold sol on adding 1 mL of 10% solution of sodium chloride.

Smaller the value of gold number, of a lyophilic colloid, the greater is its protective power.

Colloids in everyday life :
(i) Foods, medicines and pharmaceutical preparation.
(ii) Soaps, synthetic detergents,
(iii) Paints, varnishes, enamels, resins, gums, glues, rubber.
(iv) Industrial processes such as tanning, dyeing, rubber plating and removal of smoke from air are based on the colloidal nature of particles.

Suspended particulate matter (SPM) : Many partículate matter, such as fine particles of dust, sand, coal dust, pollen grains, cement dust, metal dust, fly-ash etc. remain suspended in normal air; these are termed as pollutants which cause air pollution.

Solubility : Solubility of a given solute in a solvent is defined as the weight in grams of the solute dissolved in 100 grams of the solvent so as to saturate the solution at a given temperature.

Solubility is just a number, it has no unit,

Effect of temperature and pressure on soluhility of gas in liquid:
(a) Effect of tempe rature : The solubility of gases in liquids decreases with the increase of temperature and at lower temperature, solubility increases.

(b) Effect of pressure : At a definite temperature the solubility of a gas in a particular solvent is proportional to the pressure applied on the gas. This is known as Henry’s law.

Effect of temperature Of solubility of solid in liquid :

• The solubility of KNO₃ in water, increases tremendously with the increase of temperature.
• In case of NaCl, the increase of solubility in water is much less, almost same.
• The solubility of CaSO₄in water decreases with the increase of temperature.
• The solubility of Glauber’s salt (Na₂SO₄.10₄O) in water first increases upto a certain temperature (324°C) and then sharply decreases with temperature.

Solubility curve :

Crystals : Crystals are homogeneous solid particles (big or small) with definite geometric shape, and are bounded by plane surfaces which meet at sharp edges.

Crystallisation : Crystallisation is a process by which crystals of a substance are obtained from its solution.

Water of crystallisation : When crystal is formed from the aqueous solution of it then one or more fixed number of water molecules associated with each molecules of that substance by chemical bonding are called water of crystallisation.

Hydrated crystals: Crystals having the water molecules associated with them are called hydrated crystals.

Example:
MgSO₄. 7H₂O (Epsome salt)
FeSO₄. 7H₂O (Green vitriol)
CuSO₄. 5H₂O(Blue vitriol)
H₂C2O₄. 2H<₂O (Oxalic acid)

Efflorescent substance and Efflorescence: Certain hydrated crystals when exposed to air at ordinary temperature lose their water of crystallisation partially or fully and are transformed into amorphous varieties. Such a substance is called efflorescent substance and the phenomenon is known as efflorescence.

Example:

Cause of efflorescence: Efflorescence is happened only when the vapour pressure within the hydrated crystal at ordinary temperature is greater than the vapour pressure of the atmosphere.

Deliquescent substance and Deliquescence: Certain water soluble substances when exposed to the atmosphere at ordinary temperature absorb water vapour from the air and are soluble in it to make saturated solution, such substances are called deliquescent substances and the phenomenon is known as deliquescence.

Example:
(i) MgCl₂ (Magnesium chloride)
(ii) CaCl₂(Calcium chloride)
(iii) FeCl₃ (Ferric chloride)
(iv) NaOH (Caustic soda)
N.B. : Edible Salt (NaCl) becomes moist in the rainy season due to the presence of deliquescent MgCl₂. 6H₂O as impurity in it.

Cause of deliquescence: Deliquescence occurs when the vapour pressure of water in the delequescent substance is less than the vapour pressure of the atirosphere of ordinary temperature.

Hygroscopic substances: There are certain substances which when exposed to air absorb moisture from air but are not dissolved in the absorbed water. These substances are called hygroscopic substances. Example: Conc. HSO₄, CaO etc.

Drying agent: There are certain substances which absorb water or water vapour molecule from other compounds without forming any chemical bond. These substances are called drying agent.
Example: Anhydrous CaCl₂, Anhydrous MgSO₄, Conc. HSO₄ etc.

The concentration of a solution can be expressed in several different ways:
(a)Percentage by weight (% w/w): It is the amount of solute in grams present in 100 grams of the solution.

(b) Weightivolume percentage (% w/V): It is the amount of solute in grams present in 100 mL of the solution.

(c) Volume/volume percentage (% V/V): It is the volume in mL of the solute present in 100 mL of the solution.

(d) Strength: It is the number of grams of the solute dissolved per litre of the solution.

(e) Normality (N): It is the number of gram equivalent of solute dissolved per litre of a solution.

(f) Molarity (M): It is the number of moles of solute dissolved per litre of the solution.

(g) Molality (m): It is the number of moles of solute dissloved in 1000 gram of the solvent.

(h) Formality (F): Formality of a solution is the number of gram formula weight of the ionic solute (e.g. NaCl) dissolved per litre of the solution.

(j) Parts per million (ppm) : It is defined as the number of parts of a component per million parts of the solution.

Vapour pressure of liquid : The pressure developed above the liquid at a given temperature at the equilibrium point.

Lowering of vapour pressure : It is the difference in the vapour pressure of the pure solvent and that of the solution.

Raoult’s Law : The vapour pressure of a solution is equal to the product of the mole fraction of the solvent and its vapour pressure in pure state.

Ideal solution: The solution which obey Raoult’s law at all concentration and follow the conditions
\(\Delta \mathrm{H}_{\text {mix }}=0 ; \Delta \mathrm{V}_{\text {mix }}=0 \)

Non-ideal solution : The solution which show positive or negative deviations from Raoults law.

Azeotrope : The mixture of liquids which boils at constant temperature like pure liquid and has same composition of components in liquid as well as vapour phase.

Colligative properties : The properties of the solution which are independent of nature of solute but depend upon the concentration of solute particles.

Osmosis : The passage of solvent from pure solvent or solution of low concentration to the solution of higher concentration through the semi permeable membrane.

Isotonic solutions : The solutions of same molar concentration and same osmotic pressure at given temperature.

Acids:
(j) Arrhenius concept : A substance which furnishes H ions in aqueous solution.

(ii) Bronsted concept : A substance which is proton donor.

(iii) Lewis concept : A substance which is acceptor of electron pair.

Types of acids : According to their structure:

(i) Hydracids : Those acids which have hydrogen atoms and other non-metal or radicals but no oxygen atom are called hydracids. Example : Hydrochloric acid (HCl), Hydrocyanic acid (HCN).

(ii) Oxyacids : Those acids which have one or more oxygen atoms with hydrogen and other non-metal atoms are called oxyacids. Example : Nitric acid (HNO₃), Sulphuric acid (H₂SO₄).

Types of oxyacids:
(a) Ous-acids : Those oxiacids which have lesser number of oxygen atom and the valency of the main element naming the acid is less are called otis-acids.

Example: HNO₂ (Nitrous acid)
[valency of N = 3]
H₂SO₃ (Sulphurous acid)
[valency of S = 4]

(b) Ic-acids : Those oxyacids which have greater number of oxygen atom and the valency of the main element naming the acid is high. are called ic-acids.

Example: HNO₃ (Nitric acid)
[valency of N= 5]
H₂SO₄ (Sulphuric acid)
[valency of S = 6]

(c) Hypoacids : Those acids which have less oxygen atom than ous-acids are called hypoacids.
Example : HOCl (Hypochiorous acid)

(d) Per-acids : Those acids which have greater number of oxygen atoms than ic-acids are called per-acids. Example : HClO₄ Perchloric acid)

Basicity of acid : The number of H^⊕ ions produced in aqueous solution of a molecule of acid, is called basicity of acid.

Types of acids according to basicity of acid:

(a) Monobasic acids : Those acids which ionise in aqueous solution to produce one He ion are called monobasic acids. Example : HCl, HNO₃, HBr etc.

(b) Diba sic acids : Those acids which ionise in aqueous solution to produce inherit; two He ions are called dihasic acids. Example : H₂SO₄, H₂CO₃, H₂SO₃ etc.

(c) Tribasic acids : Those acids which ionise in aqueous solution to produce three He ions are called trihasic acids. Example : H₃PO₄

• Types of acids according to their ionisation :

Strong acids : The acids which ionise mostly in aqueous solution and small number of molecules exist in molecular state, are called strong acids.
Example : HCl, HNO₃, H₂SO₄ etc.

Weak acids : The acids whose most of the molecules exist in molecular state and ionise a little number of molecules, are called weak acids.
Example : CH₃COOH, H₂CO₃, H₂S etc.

Types of acids according to their source :
Mineral acids : The acids which are produced from minerals are called mineral acids.

Example :
(i) Hydrochloric acid (HCl)
[Prepared from NaCl]

(ii) Nitric acid (HNO₃)
[Prepared from KNO₃]

Organic acids : The acids containing carbon atom and prepared from animals or plants are known as organic acids.
Example :
(i) Formic acid (HCOOH) [Source : Ant]
(ii) Lactic acid [CH₃CH(OH)COOH] [Source : milk]

Properties of acids :
(i) An acid produces H ^⊕ ion in aqueous solution.
(ii) Acids are generally soluble in water and the solution is acidic or sour in taste.
(iii) The aqueous solution of acids turns blue litmus into red, orange coloured methyl orange into pinkish red. On the other hand, phenol phthalein remains colourless in aqueous solution of an acid.
(iv) Electropositive metals like Zn, Mg, Fe (which belong above hydrogen in electrochemical series) react with dilute acid to produce hydrogen.
Example :
Zn + H₂SO₄ = ZnSO₄ + H₂↑
2Al + 6HCl – 2AlCl₃ + 3H₂↑

Acids form salt and water reacting with metallic oxides and hydroxides.
Example :
2NaOH + H₂SO₄ = Ha₂SO₄ + 2H₂O
CaO + 2HCl = CaCl₂ + H₂O

Acids form carbon dioxide reacting with metallic carbonate and bicarbonate.
Example :
Na₂CO₃ + 2HCl = 2NaCl + CO₂ + H₂O
NaHCO₃ + HCl = NaCl + CO₂ + H₂O

Fuming nitric acid : It is a product produced by dissolving excess amount of nitrogen dioxide in nitric acid.

Acid rain : Dissolved NO₂ and SO₂ of atmosphere produce nitric acid (HNO₃) and sulphuric acid
(H₂SO₄)which come down along with rain. This phenomenon is known as acid rain.

Stone cancer : Corrosion of marble or stone due to acid rain is called stone cancer.

Aqua regia : It is a mixture of 3 volumes concentrated hydrochloric acid and 1 volume concentrated nitric acid which can dissolve metals like gold and silver.

Royal water : Aqua regia is called royal water.

Fuming sulphuric acid : Fuming sulphuric acid or oleum (H₂S₂O₇) is obtained when sulphur trioxide is passed over 98% Sulphuric acid.

Muriatic acid : Hydrochloric acid.

Aqua fortis : Nitric acid.

King of chemicals : Sulphuric acid.

Laughing gas : Nitrous oxide (N₂O).

Preparation of hydrogen chloride in laboratory :

Preparation of nitric acid in laboratory :

Manufacture of sulphuric acid by contact process :

Identification of acids :

(a) Identification of Hydrochloric acid :

(b) Identification of Sulphuric acid :

(c) Identification of Nitric acid (Ring test) :

Formation of Hydrogen from nitric acid : Magnesium liberates hydrogen from cold and very dilute nitric acid.
Mg + 2HNO₃ = Mg(NO₃)₂+ H₂

Passive iron : Cold and concentrated nitric acid or fuming nitric acid when comes in contact with iron produces passive iron. Passive iron is chemically inactive.

Uses of hydrochloric acid :

• It is used in dyeing and calicoprinting.
• It is used as a cleaning agent in galvanising and tin plating.
• It is used in the manufacture of glucose, glue and many useful metallic chlorides.

Uses of nitric acid :

• It is used in the production of explosives such as dynamite, gun cotton, nitro-glycerine, trinitro toluene (TNT).
• It is used to prepare rayon, artificial silk and dyes.

Uses of sulphuric acid :

• It is used for the preparation of different chemical compounds like HCl, HNO₃, ether, alcohol etc.
• It is used as an important raw material for dyes, medicine, plastic etc.
• Oxides : Binary compound of oxygen with any element (metallic and non-metallic) are called oxides.

Types of Oxides :

(a) Acidic oxide : An acidic oxide is an oxide which reacts with a base to form salt and water.

• Non-metallic oxide : CO₂, SO₂, P₂O₅, SiO₂
• Metallic oxide : CrO₃ (Chromium trioxide), Mn₂O₇ (Manganese heptoxide) etc.

(b) Basic oxide: A basic oxide is an oxide which reacts with an acid to form salt and water. These are generally the oxides of metals.
Example :
CaO (Calcium oxide)
MgO (Magnesium oxide)
FeO, CuO etc.

(c) Neutral oxide : The oxides which neither react with acid nor with base are called neutral oxide.
Example : CO (Carbon monoxide)
H₂O (Water), N₂O (Nitrous oxide)
NO (Nitric oxide)

(d) Amphoteric oxide : The oxides which react with both acids and bases to produce salt and water are called amphoteric oxide.
Example :
Al₂O₃ (Aluminium oxide)
ZnO (Zinc oxide)
SnO (Stanous oxide)
PbO (Lead monoxide) etc.

Peroxide : The oxides which react with cold and dilute mineral acid to produce hydrogen peroxide and have peroxylinkage (— O — O —) are called peroxide.
Example :
Na₂O₂ (Sodium peroxide)
BaO₂ (Barium peroxide)

Mixed oxide : An oxide which is formed by oxides of more than one oxide of a metal having variable valency is called mixed oxide.
Example :
Fe₃O₄ [FeO + Fe₂O₃] ; Mn₃O₄ [2MnO + MnO];
Pb₃O₄ [PbO₂ + 2PbO] etc.

Poly-oxide : The oxides which have oxygen atoms more than an ordinary oxide and also do not react with acid to produce hydrogen peroxide are called poly-oxides.
Example :
Mn₂O₇ (Manganese heptoxide)
PbO₂ (Lead dioxide)

Sub-oxide : An oxide which has less oxygen atoms rather than required for oxidation state of the element is called sub-oxide. Example : Carbon sub-oxide (C₃O₂)

Super-oxide : Metals like Na, Li, Ca form super-oxide and have negative ion (O-O) in it.
Example : LiO₂, KO₂ etc.

Alkali: Those metallic oxides and hydroxides react with acids to produce salts and water are known as alkali. Example : CaO, Na₂O, NaOH, Al(OH)_a etc.

Base (According to Arrhenius concept): The metallic oxides or hydroxides after dissolving in water produce negatively charged hydroxyl ions are known as base.

Properties of bases :

• The aqueous solution of bases turn red litmus into blue, orange coloured methyl orange into yellow.
• Bases produce salts and water when react with acids.
• Concentrated solutions of some strong bases are slippery to touch.
• The aqueous solution of bases can conduct electricity.
• Strong bases when heated with electro-positive elements (like Al, Zn etc.) produce hydrogen.

Example : 2NaOH + Al + 2H₂O = 2NaAlO₂ + 3H₂↑
Strong bases absorb carbon dioxide from atmosphere to form carbonate salt and water.
Example : 2NaOH + CO₂ = Na₂CO₃ + H₂O

Difference between alkali and bases : The oxides and hydroxides of metal and metal like radicals are called alkalis. But those alkalis which dissolve in water to form OH are called bases.

Example :

• Acidity of bases : The number of OH ions produced in aqueous solution of one molecule of bases are called the acidity of bases.

Types of bases according to acidity of bases :

(a) Monoacidic bases : The bases which produce one OK^– ion in aqueous solution from one molecule are called monoacidic bases.
Example:

(b) Diacidic bases : The bases which produce two OH” ion in aqueous solution from one molecule are called diacidic bases.
Example :

Types of bases according to their ionisation :

Strong bases : The bases which ionise mostly in aqueous solution and small number of molelcules exist in molecular state are called strong bases.
Example : NaOH, KOH, Ca(OH)₂ etc.

Weak bases : The bases whose most of the molecules exist in molecular state and ionise a little number of molecules are called weak bases.
Example : NH₄OH, Cu(OH)₂ etc.

Types of bases according to their source :

Mineral bases : The bases which are produced from minerals are called mineral bases.
Example : NaOH, KOH etc.

Organic bases : The bases cortaining nitrogen atom and whose sources are animals and plants are called Organic bases.
Example:

Salts : The replaceable hydrogen atoms in an acid when replaced by metal or basic radical partially or fully then the compound so produced is called salt.

Definition of salts according to Arrhenius concept: The compound formed by cation other than hydrogen ion (H)⁺ and anion other than hydroxyl ion (OH~) is known as salt.

Types of salt :

• Normal salts : Normal salts are formed by the complete replacement of the hydrogen atoms in an acid by metals or basic redicals.

Example :
(i) NaCl [Sodium chloride] [NaOH + HCl = NaCl + H₂O]
(ii) Na₂SO₄ [Sodium Sulphate] [2NaOH + H₂SO₄ = NA,SO₄ + H₂O]

Acid salt: An acid salt is one which is formed only by partial replacement of hydrogen atom in a molecule of an acid (di-basic, tri-basic) by metal or basic radical.
Example :
NaHSO₄ (Sodium bisulphate)
NaHCO₃ (Sodium bicarbonate)
Na₂HPO₄ (Di-sodium hydrogen phosphate)

Basic Salt: When the hydroxyl (OH) or, oxide (O) radicals are partially neutralised by acid then basic salts are produced.
Example :
(i) Pb(OH)Cl [Basic lead chloride]
(ii) Pb(OH)NO₃ [Basic lead nitrate]

Double salts : When two normal salts are mixed with each other with their molecular weight ratio to form a solution and cooled the solution to make a combined crystal which is stable in solid state but ionises into different constituent ions in solution are called double salts.
Example :
K₂SO₄ Al₂(SO₄)₃ . 24H₂O
(NH₄)₂SO₄FeSO₄ . 6H₂O

Complex salts : When the solution of two mixed salts is concentrated and cooled then a crystal of new salt is produced. The constituent salts lose their identity in new salt. In aqueous solution they produce new complex ions. This type of salts are called complex salts.
Example : K₄[Fe(CN)₆] ; [Cu(NH₃)] SO₄

Neutralisation : The reaction in which equivalent amount of an acid reacts with equivalent amount of a base and thus the properties of acid and bases are completely lost forming salt and water, is called neutralisation reaction.
Example :

Titration : The process by which bases are neutralised with the help of acids or vice-versa is called titration.

Indicators : The chemicals which are able to determine the end point of neutralisation reaction by changing their colours are called indicators.

Indicators used in neutralisation reaction :

Name of the indicators and their actual colour	In acid solution	In alkali solution	In neutral solution
1. Litmus (violet)	red	blue	violet
2. Phenolphthalein	colourless	pink	colourless
3. Methyl-orange (orange)	red	yellow	orange

Choice of indicators :
(a) Strong acid and weak base — Methyl orange
(b) Weak acid and strong base — Phenolphthalein
(c) Strong acid and strong base — Any indicator
(d) Weak acid and weak base — No such indicator.

Vanishing colour : Vanishing colour is a coloured solution which becomes colourless on exposure to air after a lapse of time.
Example :
The aqueous solution of ammonium hydroxide (NH₄OH) mixed with phenolphthalein turns pink colour. Now if this solution is allowed to spray on cloth, at once a pink colour is developed upon cloth. But since ammonia in NH₄OH is highly volatile, after lapse of time NH₃ disappears from the cloth and the solution becomes neutral. As a result pink colour of phenolphthalein vanishes.

Vanishing colour cannot be prepared with dilute NaOH solution, due to the fact that NaOH does not evaporate.

pH : The symbol pH is derived from ‘Poteuz,’ the Danish word for power, pH refers to potency of hydronium ion.  pH = -log [H] = – log[OH^–]

pH range :

• The pH of acidic solution is less than 7. [pH < 7, the solution will be acidic]
• The pH of basic solution is greater than 7. [ pH > 7, the solution will be basic]
• The pH of neutral solution is 7. [pH = 7, the solution will be neutral] So, the range of pH is from 0 to 14.

Colour changes for universal indicator at different pH values.

• The approximate pH of a solution can be determined with the help of pH papers.
• pH -papers have different colours in solutions of different pH.
  A pH paper can determine pH of a solution with an accuracy of about 0 -5.

pH-metres : For accurate measurement of pH (upto accuracy 0.001 units), pH metres are used.

Universal indicator : A mixture of organic dyes which gives different colours with solutions of different

pH values is called a universal indicator.

	Name of the Fluid	pH
1.	0 1(M) NaOH solution	13
2.	Milk of magnesia	10
3.	Egg white, sea water	7-8
4.	Human blood	7-4
5.	Tears	7-4
6.	Milk	6-8
7.	Human Saliva	64
8.	Tomato juice	4-2
9.	Soft drinks and Vinegar	30
10.	Lemon juice	2-2
11.	Gastric juice	1.2
12.	1(M) HCl solution	0

Importance of pH :

Field	Utility
1. Biotechnology	Biochemical and organic reactions at controlled pH value give best results.
2. Agriculture	Some crops such as citrus fruits grow better in alkaline soils. Sugarcane grows better in neutral soil, whereas rice grow better in slightly acidic soil. The pH of soil is tested before growing particular crop.
3. Medicine	Determination of pH value of blood and urine helps to diagnose certain disease.
4. Cosmetics	Soaps, shampoos, face creams are prepared for the consumers having different pH value of skin secretions.
5. Milk plants	pH of milk is rigorously controlled at pH 6 8 as otherwise it turns sour.
6. Breweries	Controlling pH value of wine to obtain a desired flavour.

Antacids : Medicines which can remove the excess acid from the stomach and raise the pH to appropriate level are called antacids.

Types of antacids:

Cimetidine (Tegamet): It binds to the receptors that trigger the release of hydrochloric acid into the stomach. This results in release of lesser amount of acid.

Cimetidine remained the largest selling drug in the world for a long time.

(ii) Ranitidine (Zantac): After cimetidine, ranitidine (Zantac) was introduced for treatment of hyperacidity.

Separation of mixtures:

Fractional distillation : The process by which two or more miscible liquids are separated by distillation using the difference in their boiling point is called fractional distillation.

When the boiling points of two or more miscible liquids differ by about 15°C — 20°C these are separated by simple fraction distillation.

fractionating column : If the boiling points of two liquids are very close i.e. difference is small in that case a fractionating column is used to separate the liquid mixtures.

Separation of immiscible liquids using separating funnel : Immiscible liquids are easily separated by using a separating funnel. Each liquid forms its own liquid layer due to difference in density.

Chromatography : It is a method by which mixture of different substances are separated. In this method due to difference in adsorption of different substances in solid phase (adsorbent) and difference in migration of the substances in the mobile phase (liquid or gas), various substances are separated.

Paper Chromatography : It is a very easy technique to separate the various organic dyes present in the writing ink or printing ink.

Advantages of chromatography :

• A small amount of the compound present in the mixture can be separated.
• The properties of the individuals present in the mixture do not alter.
• It is a very easier method to separate the different components present in the mixture.

Sublimation : There are some solids which on heating directly transform into vapour phase without transforming through the intermediate liquid state. This phenomenon is called sublimation.
Example : NH₄Cl (ammonium chloride) and SiO₂ (sand) are separated by this method.

Natural water : Natural water, except rain water, is generally impure and contains dissolved metals, apart from other impurities.

Sources of natural water :

• Rain water
• River water
• Spring and well-water
• Mineral water
• Sea water

Soft water : Water in which lather is easily formed with soap is called soft water.

Hard water : The class of water that does not easily form lather with soap, is called hard water.

Hardness of water : The property of natural water which does not allow it to form lather is known as hardness of water.

Cause of hardness of water : The hardness of water is caused by the presence of bicarbonate, chloride and sulphate of calcium, magnesium and iron in water.

Types of hardness :

• Temporary hardness : It is due to the presence of soluble bicarbonate of calcium, magnesium and iron in water.
• Permanent hardness : This is caused by the presence of chloride and sulphate of calcium, magnesium and iron in water.

Removal of temporary hardness :

Removal of permanent hardness :
Lime-soda process :
CaSO₄ + Na₂CO₃ = CaCO₃↓ + Na₂SO₄
CaCl₂ + Na₂CO₃ = CaCO₃↓ + 2NaCl

Permutit process :
2Na (Permutit) + Ca(HCO₃)₂ = Ca(Permutit)₂ ↓+ 2NaHCO₃

[Zeolite, a natural mineral is a giant molecule of sodium aluminium silicate (NaAlSO₄ 3H₂O). Artificially prepared mineral called permutit is used to remove permanent hardness]

Calgon process :

Ion exchange resin process :

De-ionised water : Very soft and pure water which is obtained by treating hard water through the ion-exchange process, is free from all types of ions i.e., cations and anions. Such type of water is known as de-ionised water.

Water is an universal and versatile solvent :

• Water is a polar compound. Most of the ionic compounds are soluble in polar solution and thus water is, therefore, a good ionising solvent for acids, bases and salts.
• Water is an amphiprotic solvent i.e. it can act both as an acid (a proton donar) and a base (a proton acceptor) in the reaction.
• Many covalent compounds like sugar, urea, alcohol, organic acid, organic bases are soluble in water through hydrogen bonding. Hence, we may conclude that water is an universal as well as versatile solvent.

Water pollution : Some soluble and insoluble matters in water which produce some harmful effect to men and aquatic animals are known as water pollution.

Pollution due to natural source :

Pollution due to arsenic : Ground water is sometime contaminated with arsenic acid and arsenius acid as a deprotonated form. Consumption of these arsenic contaminated water for a long period can cause dangerous poisonous effects in the form of black spots on palms and feet and skin become rough. This is known as black foot disease. From scientific researches it can be assumed that arsenic poison may be a cause of skin cancer.

Pollution due to fluoride : Sea-water, river-water contain fluoride from natural source. Excess fluoride in water is responsible for fluorosis disease which is a cause of dental decay and bone decay.

Pollution due to artificial or human activity : Fertilizers, insectisides, pestisides, are regularly used for increasing the production of food grains. These substances when washed with water are mixed in ponds, river etc. As a result of this, water is polluted.

Removal of arsenic from water :

• Arsenic is removed from water if water is allowed to pass through a bag made of cloth containing alumina or ferric hydroxide replacing the candle used in the filter. Arsenic or arsenius ions are absorbed in the alumina bed or ferric hydroxide bed and thus are removed.
• Take 20 litre of water, \(\frac{1}{4}\)th amount of tea spoon of bleaching powder and the same amount of ferrous sulphate in a bucket. After stirring well, arsenic will be removed.