08. Short and Long Answer Questions: Journey Inside the Atom

Short Answer Type Questions

Ques 1: What were the contributions of Acharya Kanada and the Greek philosophers Leucippus and Democritus to the concept of the atom?

Ans: Acharya Kanada, an ancient Indian thinker, proposed that if matter (dravya) is divided repeatedly, a stage is reached where the smallest indivisible particles remain. He called these particles parmanus, recording his ideas in the Sanskrit text Vaisesika Sutras. According to him, a parmanu is infinitely small and cannot be perceived by the senses. Combinations of parmanus in pairs (dyads) and triplets (triads) form all material substances.

Similarly, the Greek philosophers Leucippus and Democritus proposed that matter consists of tiny indivisible particles, which they named atomos – a Greek word meaning indivisible. Both traditions arrived at remarkably parallel ideas about the smallest units of matter, although neither was based on experimental observations.

Ques 2: How did J. J. Thomson discover the electron? What conclusion did he draw about the nature of cathode rays?

Ans: In 1897, J. J. Thomson studied the conduction of electric current through gases at very low pressure. He used a glass tube fitted with two electrodes and applied a high voltage. He observed rays travelling from the cathode (negative electrode) to the anode (positive electrode). These were called cathode rays.

By studying these rays in electric and magnetic fields, Thomson concluded that they consist of streams of negatively charged particles with a much smaller mass than atoms. He named these particles electrons. He also found that the nature of cathode rays was independent of both the cathode material and the gas in the tube, showing that electrons are a fundamental component of all atoms. The charge of an electron is taken as $- 1$ (i.e., $- 1.602 \times 1 0^{- 19}$ C) as a matter of convention.

Ques 3: Describe Thomson’s model of the atom. Why is it also called the plum pudding model?

Ans: After discovering electrons, Thomson faced a puzzle: since electrons are negatively charged and atoms are electrically neutral, there must be a positive charge present somewhere. To explain this, he proposed that an atom is a sphere of uniform positive charge with electrons distributed throughout it.

This model was compared to a pudding with plums embedded in it – the positively charged sphere acts like the pudding and the electrons act like the plums scattered inside it. Hence, it is popularly called the plum pudding model. A simpler comparison is a watermelon, where the red pulp represents the positive charge and the seeds represent the electrons.

This model was the first genuine attempt to describe how an atom’s positive and negative charges stay balanced, though it was later disproved by Rutherford’s gold foil experiment.

Ques4: What is radioactivity? How did its discovery challenge Dalton’s atomic theory?

Ans: Scientists discovered in the late 19th century that certain elements emit invisible energy and particles spontaneously. This phenomenon is known as radioactivity.

Until that time, atoms were believed to be the smallest, indivisible units of matter – as Dalton had proposed. However, the discovery of radioactivity showed that atoms must be composed of even smaller subatomic particles, since energy and particles were being emitted from within the atom itself.

This directly challenged Dalton’s atomic theory, which stated that atoms are indivisible and cannot be broken down into smaller parts. Radioactivity proved that atoms have an internal structure and are made up of subatomic components, thus laying the foundation for further investigation into the structure of the atom by scientists such as Thomson and Rutherford.

Ques 5: How did Rutherford discover the proton? How does the presence of protons explain the electrical neutrality of atoms?

Ans: Rutherford showed through the gold foil experiment that the nucleus of an atom carries positive charge. He identified the particles responsible for this positive charge and named them protons. Protons are much heavier than electrons and possess a charge that is equal in magnitude but opposite in sign to that of an electron. The relative charge of a proton is taken as $+ 1$ .

For an atom to be electrically neutral, the total positive charge from protons must exactly balance the total negative charge from electrons. This means:

Number of protons = Number of electrons

For example, a helium atom has 2 protons and 2 electrons, while a sodium atom has 11 protons and 11 electrons. In both cases the charges cancel out completely, making the atom electrically neutral. Similarly, all atoms in nature are electrically neutral.

Ques 6: According to Bohr’s model, how does an electron move from one shell to another? How does the energy of shells change as we move away from the nucleus?

Ans: According to Bohr’s model, electrons revolve in fixed circular paths called stationary states or shells (energy levels), designated as K, L, M, N, … corresponding to $n = 1, 2, 3, 4, \dots$ An electron in a stationary state does not lose energy while revolving.

Movement between shells: An electron can move from one shell to another by absorbing or releasing a fixed amount of energy equal to the difference in energy between the two shells. If energy is absorbed, the electron jumps to a higher shell (farther from the nucleus); if energy is released, it falls back to a lower shell.

Energy of shells with distance from nucleus: The energy of shells increases as we move away from the nucleus. The K-shell $(n = 1)$ , being closest to the nucleus, has the least energy. The L-shell $(n = 2)$ has more energy than the K-shell, and so on. Thus, the farther a shell is from the nucleus, the higher its energy level.

Ques 7: How was the neutron discovered? What role does it play in the nucleus?

Ans: The neutron was discovered in 1932 by James Chadwick, a student of Ernest Rutherford. Scientists had noticed a puzzle: a helium atom has two protons yet its mass is about four times that of a hydrogen atom (not double). This suggested there must be another particle in the nucleus that adds mass without adding charge.

Chadwick discovered a new subatomic particle with a mass nearly equal to a proton but with no electrical charge. This neutral particle was named the neutron, represented by the symbol $n^{0}$ .

Neutrons are found in the nucleus of all atoms except hydrogen. Along with protons, they contribute to the mass of the atom. Neutrons also help reduce the repulsion between positively charged protons in the nucleus by increasing the spacing between them and strengthening the nuclear force that holds the nucleus together.

Ques 8: What are atomic number and mass number? Write the standard notation used to represent an atom, with an example.

Ans: The atomic number of an element is the number of protons present in the nucleus of its atom. It is denoted by the symbol $Z$ . Since atoms are electrically neutral, the number of electrons equals the number of protons, so $Z$ also gives the number of electrons.

The mass number is the total number of protons and neutrons (collectively called nucleons) in the nucleus. It is denoted by $A$ . $A = Number of protons + Number of neutrons$

The standard notation for an atom places the mass number at the top-left and the atomic number at the bottom-left of the element’s symbol: $_{Z}^{A} X$

For example, carbon has atomic number $Z = 6$ and mass number $A = 12$ , so it is written as $_{6}^{12} C$ .

Ques9: What are the rules followed by Bohr and Bury for distributing electrons in the shells of an atom?

Ans: Bohr and Bury suggested the following rules for the distribution of electrons in the energy shells of an atom:

The maximum number of electrons that can be present in any shell is given by the formula $2 n^{2}$ , where $n$ is the shell number. Thus: K-shell $(n = 1)$ holds up to $2$ electrons, L-shell $(n = 2)$ holds up to $8$ , and M-shell $(n = 3)$ holds up to $18$ .
The maximum number of electrons in the outermost shell is 8 (except the first shell, which holds a maximum of 2).
Electrons are filled in shells in a stepwise manner, starting from the shell closest to the nucleus (K), then L, M, N, and so on. The next shell begins filling only after the inner shell is complete.

Ques10: What is valency? Explain with the examples of sodium and oxygen how valency is determined from the electronic configuration.

Ans: The valency of an element is the number of electrons an atom gains, loses, or shares to complete its outermost shell (octet) and attain a stable configuration. It represents the combining capacity of an atom.

Sodium (Na): Its electronic configuration is $2, 8, 1$ . The outermost shell has only 1 electron. Since fewer than 4 electrons are present in the valence shell, sodium tends to lose 1 electron to achieve an octet. Therefore, the valency of sodium is 1.

Oxygen (O): Its electronic configuration is $2, 6$ . The outermost shell has 6 electrons. Since more than 4 valence electrons are present, oxygen tends to gain 2 electrons to complete its octet. Therefore, the valency of oxygen is 2.

Ques11: What is average atomic mass? Chlorine exists as two isotopes 1735Cl (75%) and 1737Cl (25%). Calculate the weighted average atomic mass of chlorine.

Ans: The average atomic mass of a natural element is the weighted average of the masses of all its naturally occurring isotopes, calculated by considering their relative (percent) abundances in nature. It is more accurate than a simple arithmetic mean because isotopes do not occur in equal proportions.

The weighted average atomic mass is calculated by multiplying the mass of each isotope by its fractional abundance and then adding the results.

Calculation for Chlorine:

Chlorine has two isotopes: $^{35} Cl$ (75% abundance) and $^{37} Cl$ (25% abundance). $Weighted average atomic mass = (35 \times \frac{75}{100}) + (37 \times \frac{25}{100})$ $= \frac{2625}{100} + \frac{925}{100} = \frac{3550}{100} = 35.5 u$

Therefore, the weighted average atomic mass of chlorine is 35.5 u. Note that no single chlorine atom has a mass of 35.5 u – this value reflects the average across a large number of atoms as they occur in nature.

Ques 12: What is Dalton’s atomic theory? How did it serve as the starting point for understanding atomic structure?

Ans: In 1808, John Dalton proposed his atomic theory based on scientific experiments of that time. The key ideas of his theory are:

All matter is composed of extremely small particles called atoms.
Atoms are indivisible – they cannot be broken down into smaller parts.
Atoms are the fundamental building blocks of matter.
Atoms of the same element are identical to each other, while atoms of different elements are different.

Dalton’s atomic theory was the first scientific description of how matter is made. It replaced the earlier imaginary concepts (like the parmanu of Kanada and the atomos of Greek philosophers) with ideas grounded in experimental evidence.

It became the starting point for the current understanding of atomic structure because it prompted scientists to ask further questions – What are atoms made of? What do they look like? What makes atoms of one element different from another? – questions whose pursuit eventually led to the discovery of subatomic particles and modern atomic models.

Long Answer Type Questions

Ques 1: Describe Rutherford’s model of the atom. What were the conclusions he drew from the gold foil experiment? Also mention any two limitations of this model.

Ans: From the results of the gold foil experiment, Rutherford proposed the nuclear model (also called the planetary model) of the atom. His key conclusions and proposals were:

Most of the atom is empty space, since most $α$ -particles passed through the gold foil without deflection.
All the positive charge and most of the mass of the atom are concentrated in an extremely small, dense central region called the nucleus. The large deflection and bouncing back of a few $α$ -particles indicated this concentrated positive charge.
Electrons revolve around the nucleus in circular orbits, somewhat like planets revolving around the Sun – hence it is called the planetary model.
The nucleus is about $1 0^{5}$ times smaller than the atom. The diameter of an atom is approximately $1 0^{- 10}$ m, while the diameter of the nucleus is approximately $1 0^{- 15}$ m.

Rutherford also discovered the proton – he showed that the nucleus carries positive charge due to particles called protons. Protons are much heavier than electrons and possess an equal but opposite charge. For an atom to be neutral, the number of protons must equal the number of electrons.

Limitations of Rutherford’s Model:

Cannot explain atomic stability: A revolving electron is always accelerating. An accelerating charged particle must continuously lose energy (emit radiation). As a result, the electron should spiral inward and collapse into the nucleus. But atoms are stable, which Rutherford’s model could not explain.
Does not describe electron arrangement: Rutherford’s model did not specify how electrons are arranged or distributed around the nucleus, nor did it account for the specific energies that electrons possess.

Ques 2: Compare the atomic models proposed by Thomson, Rutherford, and Bohr. What was the major drawback of each model and how did the next model overcome it?

Ans: Atomic models evolved as new experimental evidence emerged. The comparison is given below:

Feature	Thomson’s Model	Rutherford’s Model	Bohr’s Model
Structure proposed	Atom as a sphere of uniform positive charge with electrons embedded in it (plum pudding)	Dense central nucleus with electrons revolving around it (planetary model)	Electrons revolve in fixed energy shells (orbits) around the nucleus
Proposed by / Year	J. J. Thomson (after 1897)	Ernest Rutherford (1911)	Niels Bohr (1913)
Major drawback	Could not explain deflection of $α$ -particles in the gold foil experiment	Could not explain why electrons do not lose energy and fall into the nucleus (stability of atom)	Was found to have limitations; a new quantum mechanical model was later proposed (to be learnt in higher grades)
How next model overcame it	Rutherford’s experiment showed positive charge is concentrated in a nucleus, not spread evenly	Bohr introduced the concept of stationary states where electrons do not lose energy	Quantum mechanical model (higher grades) treats electrons as electron clouds

Each model was a step forward driven by curiosity, questioning, and experimentation. Although the early models were not fully correct, they represent how science progresses – one step at a time.

Ques3: Explain the rules for writing symbols of elements as per IUPAC norms. Why are standard symbols necessary? Give five examples of elements whose symbols are derived from languages other than English.

Ans: John Dalton was the first to introduce symbols for elements in 1803, using pictorial representations. In 1813, Berzelius suggested that symbols should be derived from the Latin names of elements, giving rise to alphabetic chemical symbols. Today, the International Union of Pure and Applied Chemistry (IUPAC) approves all names and symbols.

Rules for writing element symbols (IUPAC norms):

Many symbols are the first letter or the first two letters of the element’s name (e.g., Hydrogen → H; Aluminium → Al).
The first letter is always a capital (uppercase) letter, while the second letter (if any) is written in lowercase. Example: Cobalt → Co (not CO); Aluminium → Al (not AL).
Some symbols are formed from the first letter and a letter other than the second letter of the name. Example: Chlorine → Cl; Zinc → Zn.
Symbols of some elements come from their Latin, Greek, or German names rather than their English names.

Five examples of symbols derived from non-English names:

Element	Symbol	Source language & name
Iron	Fe	Latin – Ferrum
Copper	Cu	Latin – Cuprum
Gold	Au	Latin – Aurum
Mercury	Hg	Greek – Hydrargyros
Tungsten	W	German – Wolfram

Why standard symbols are necessary: Scientists use standard symbols because they are internationally recognised and allow scientists worldwide to communicate clearly, regardless of language barriers. If every scientist used different symbols for the same element, it would cause enormous confusion in research and industry.

Ques 4: Define isotopes and isobars. Write the electronic configuration of the three isotopes of hydrogen. State two similarities and two differences between isotopes and isobars.

Ans: Isotopes: Atoms of the same element having the same atomic number ( $Z$ ) but different mass numbers ( $A$ ) due to a different number of neutrons are called isotopes.

Isobars: Atoms of different elements that have the same mass number ( $A$ ) but different atomic numbers ( $Z$ ) are called isobars.

Three isotopes of hydrogen and their electronic configurations:

Isotope	Symbol	Protons	Neutrons	Electrons	Electronic Config.
Protium	$_{1}^{1} H$	1	0	1	K: 1
Deuterium	$_{1}^{2} H$	1	1	1	K: 1
Tritium	$_{1}^{3} H$	1	2	1	K: 1

All three isotopes have the same electronic configuration (1 electron in the K-shell) and hence the same chemical properties.

Two similarities between isotopes and isobars:

Both isotopes and isobars are atoms that differ in their nuclear composition (number of neutrons or protons).
Both are subclassifications of atoms used to explain why atoms of the same or different elements can have varying masses.

Two differences between isotopes and isobars:

Property	Isotopes	Isobars
Atomic number (Z)	Same (same element)	Different (different elements)
Mass number (A)	Different	Same

Ques 5: An atom of an element X has a mass number of 31 and 16 neutrons in its nucleus.
(i) Find the atomic number of X.
(ii) Identify the element X.
(iii) Write its electronic configuration.
(iv) Find the number of valence electrons and determine its valency.
(v) Write the standard notation for the atom of element X.

Ans: (i) Finding the atomic number:

Given: Mass number $A = 31$ , Number of neutrons $= 16$ $A = Number of protons + Number of neutrons$ $31 = Z + 16$ $Z = 31 - 16 = 15$

Therefore, the atomic number of element X is 15.

(ii) Identification of element X:

The element with atomic number 15 is Phosphorus (P).

(iii) Electronic configuration of Phosphorus:

Total electrons = 15 (equal to atomic number)

K-shell: 2 electrons
L-shell: 8 electrons
M-shell: 5 electrons

Electronic configuration: 2, 8, 5

(iv) Valence electrons and valency:

The outermost (M) shell contains 5 valence electrons. Since the number of valence electrons is more than 4, phosphorus tends to gain 3 electrons to complete its octet of 8 electrons in the outermost shell. The number of electrons gained to attain a stable configuration gives the valency.

Therefore, the valency of X (Phosphorus) is 3.

(v) Standard notation for atom of element X: $_{15}^{31} P$