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I. CONCEPTUAL UNDERSTANDING
1. What information does the electronic configuration of an atom provide? (AS - 1) (2 Marks)
A: The distribution of electrons in shells, sub-shells and orbitals of an atom is known as electronic configuration.
The short hand notation consists of the principal energy level (n value) the letter representing the sub-level (l value) and the number of electrons (x) in the sub-shell is written as
→ for Hydrogen atom (H) having nl^{x} atomic number (Z) = 1, the number of electrons is one, then the electronic configuration is 1s^{1}.
2. a) How many maximum number of electrons that can be accommodated in a principal energy shell? (AS - 1) (2 Marks)
A: The maximum number of electrons than can be accommodated in a shell is 2n^{2} where n is the principal quantum number.
b) How many maximum number of electrons that can be accommodated in a sub - shell? (AS - 1) (2 Marks)
A: * Each sub - shell holds a maximum of twice as many electrons as the number of orbitals in the sub-shell.
* The maximum number of electrons that can occupy various sub - shells is given in the following table.
c) How many maximum number of electrons can be accommodated in an orbital? (AS - 1) (1 Mark)
A: The maximum number of electrons that can be accommodated in an orbital is 2.
d) How many sub-shells present in a principal energy shell? (AS - 1) (2 Marks)
A: * The number of sub - shells present in a principal energy shell is equal to the principal quantum number.
* The following table gives the details.
e) How many spin orientations are possible for an electron in an orbital? (AS - 1) (1 Mark)
A: * Only two orientations of the spin of an electron in an orbital are possible.
* One clock - wise (+ ) and the other anticlock - wise ( - )
* Their spin quantum numbers are + and - respectively.
3. In an atom the number of electrons in M - shell is equal to the number of electrons in the K and L shell. Answer the following questions. (AS - 1)(4 Marks)
a) Which is the outermost shell? (1 Mark)
A: N - shell
b) How many electrons are there in the outermost shell? (1 Mark)
A: 2 electrons
c) What is the atomic number of the element? (1 Mark)
A: The atomic number of the element is 22.
d) Write the electronic configuration of the element.
A: The electronic configuration is 1s^{2} 2s^{2} 2p^{6} 3s^{2} 3p^{6} 4s^{2} 3d^{2}.
4. Rainbow is an example for continuous spectrum. Explain. (AS - 1) (4 Marks)
A: For continuous spectrum formation of Rainbow in nature is a familiar example.
There are seven colours namely violet, indigo, blue, green, yellow, orange and red (VIBGYOR) in a rainbow.
* These colours are spread continuously and the intensity of each colour changes from one point to another.
* Each colour in a rainbow is characterised by a specified wavelength from red (higher wavelength) to violet (shorter wavelength).
* Due to higher intensity of one particular emitted colour we cannot observe all the seven colours clearly.
5. How many elliptical orbits are added by Sommerfeld in the thirds Bohr's orbit? What was the purpose of adding these elliptical orbits? (AS - 1) (2 Marks)
A: * Sommerfeld added two elliptical orbits to Bohr's third orbit.
* The purpose of adding these is to account for the splitting of line spectra which Bohr's model failed to explain.
6. What is an absorption spectrum? (AS - 1) (2 Marks)
A: Absorption spectrum: An absorption spectrum is like the photographic negative of an emission spectrum.
* A continuum of radiation is passed through a sample which absorbs radiation of certain wavelengths.
* The missing wavelengths which correspond to the radiation absorbed by the matter, leave dark spaces in the bright continuous spectrum.
7. What is an orbital? How is it different from Bohr's orbit? (AS - 1) (2 Marks)
A: Orbital: The region of space around the nucleus where the probability of finding the electron is maximum is called an orbital.
* Boh'rs orbits are stationary orbits of fixed energy at different distances from the nucleus.
* Bohr's orbits circular orbits and can accommodate a maximum of 2n2 electrons in it.
* An orbital has no boundary and it can accommodate only two electrons.
8. Explain the significance of three quantum numbers in predicting the positions of an electron in an atom. (AS - 1) (4 Marks)
A: * Each electron in an atom is described by a set of three numbers n, l and ml. These are called quantum numbers. These indicate the probability of finding the electron in space around the nucleus.
* The principal quantum number (n) is related to the size and energy of the main shell n has positive integer values of 1, 2, 3, ........
Shell | K | L | M | N |
n | 1 | 2 | 3 | 4 |
* As 'n' increases, the shells become larger and the electrons in those shells are farther from the nucleus.
The angular momentum quantum number (l): 'l' has integer values from 0 to (n - 1) for each value of n. Each 'l' value represents one sub - shell.
* The value of 'l' for particular sub - shell is generally designated by the letter s, p, d, f....... as follows.
l | 0 | 1 | 2 | 3 |
Name of sub - shell | s | p | d | f |
The magnetic quantum number (ml): It has integer values between -l and +l including zero.
* If l = 0 then ml = 1 because (2l + 1) = 1
When l = 1, then ml = 3 because (2l + 1) = 3
This means ml has three values namely -1, 0 and 1. These are labelled as px, py and pz
* The position of an electron can be predicted as follows.
For example: If the values of n, l and ml are 2, 1 and -1 respectively then the electron is present in 2 p_{x} orbital.
9. What is nlx method? How is it useful? (1 Mark)
A: * The short hand notation of electronic configuration is nl^{x}.
* This method is useful to get the following information.
10. Following orbital diagram shows the electronic configuration of nitrogen.
Which rule does not support this? (2 Marks) (AS - 1)
A: * In the electronic configuration of Nitrogen atom given the electronic configuration does not get the support of Hund's rule.
* According to Hund's rule 'the orbitals of equal energy (degenerate) are occupied with one electron each before pairing of electrons starts.
* So the correct electronic configuration for Nitrogen atom is N (Z = 7)
11. Which rule is violated in the electronic configuration 1s^{0} 2s^{2} 2p^{4}? (AS - 1) (1 Mark)
A: * Aufbau rule is violated in the electronic configuration 1s^{0} 2s^{2} 2p^{4}.
* According to Aufbau principle, the lowest - energy orbitals are filled first.
* As the lowest-energy orbital 1s^{0} is not filled up, this electronic configuration violated Aufbau principle.
* 1s orbital must be first filled up before the electron enters 2s orbital.
12. Write the four quantum numbers for the differentiating electron of Sodium atom (Na) (AS - 1) (1 Mark)
A: * The electronic configuration of Sodium (Na) (Z = 11) atom is 1s^{2} 2s^{2} 2p^{6}3s^{1}.
* The differentiating electron is in 3s orbital. Its four quantum numbers are
13. What is emission spectra? (AS - 1) (1 Mark)
A: The spectrum of radiation emitted by a substance that has absorbed energy is called an emission spectrum.
II. Asking questions and making Hypothesis
14. i) An electron is an atom has the following set of four quantum numbers to which orbital it belong to (AS - 2) (2 Marks)
ii) Write the four quantum numbers for 1s^{1} electron.
A: i) The electron in an atom belong to 2s orbital. Its spin is in clockwise direction.
ii) The four quantum numbers for 1s^{1} electron are
15. Which electronic shell is at a higher energy level K or L? (AS -2) (1 Mark)
A: * L shell is at a higher energy level than K shell.
* L shell is away from the nucleus compared to the distance of K - shell from the nucleus.
IV. Information Skills and Projects
16. Collect the information regarding the wavelengths and corresponding frequencies of three primary colours red, blue and green. (AS - 4) (4 Marks)
A:
S.No. | Primary Colours | Wavelength (nm) (1 nm = 10^{-9}m) |
Frequency (Hz) |
1. | Red | 740 to 625 | (405 to 480) × 10^{14} |
2. | Blue | 500 to 430 | (600 to 700) × 10^{14} |
3. | Green | 565 to 520 | (530 to 590) × 10^{14} |
VII. Application to Daily Life, Concern to Biodiversity
17. The wavelength of a radiowave is 1.0 m. Find its frequency (AS - 7) (1 Mark)
A: Given: λ = 1.0 m, υ = ?
We know c = 3 × 10^{8} m.s^{-1}
Questions and Answers given in the Lesson
1. What happens when you heat an iron rod on a flame? Do you find any change in colour on heating an iron rod? (AS - 1) (2 Marks)
A: * When we heat an iron rod some of the heat energy is emitted as light.
* First it turns red (lower energy corresponding to higher wavelength) and as the temperature rises it glows orange, yellow, blue (higher energy and of lower wavelength) or even white (all visible wavelengths) if the temperature is high enough.
2. Do you observe any other colour at the same time when one colour is emitted? (AS - 1) (1 Mark)
A: When the temperature is high enough, other colours will also be emitted, but due to higher intensity of one particular emitted colour (e.g.: red), others cannot be observed.
3. Variety of colours is seen from fire works. How do these colours came from fire works? (AS - 1)(2 Marks)
A: * The atoms of elements present in the fire works absorb heat energy and the electrons in the respective atoms move to excited states from their ground state.
* When these electrons return to their ground states the absorbed energy is emitted as light energy in the visible spectrum.
* So we observe variety of colours from fire works as different elements present in them emit different colours of light.
4. Do you observe yellow light in street lamps? How is this produced? (AS - 1) (1 Mark)
A: Sodium vapours produce yellow light in street lamps.
5. Why do different elements emit different flame colours when heated by the same non-luminous flame? (AS - 1) (4 Marks)
A: * All elements are made up of atoms and molecules. These atoms and molecules posses fixed energies.
* Generally these atoms will be in their ground state as it gives stability to them.
* The atoms or molecules are stable only when they are in their lowest energy state.
* When they are heated on a non-luminous flame, the electrons in the respective atoms gain energy and move to higher energy states.
* These excited atoms (electrons in higher energy state) emit light energy to lower their energies in order to come back to lowest energy state.
* The light emitted by the atoms in such process have certain fixed wavelength for each kind of atom.
* When different elements are heated naturally the light emitted by them will have different flame colours as the light comes from different atoms.
6. What happens when an electron gains energy? (AS - 1) (1 Mark)
A: When an electron gains energy it move to a higher energy level, the excited state.
7. Does the electron retain the energy forever? (AS - 1) (1 Mark)
A: * The electron loses the energy and comes back to its ground state.
* The energy emitted by the electron is seen in the form of electromagnetic energy.
8. Did Bohr's model account for splitting of line spectra of a hydrogen atom into finer lines? (AS - 1) (1 Mark)
A: Bohr's model failed to account for splitting of line spectra.
9. Why is the electron in an atom restricted to revolve around the nucleus at certain fixed distances? (AS - 1) (1 Mark)
A: To explain the atomic spectra, Bohr - Sommerfeld model proposed that the electrons are restricted to revolve around the nucleus at certain fixed distances.
10. Do the electrons follow defined paths around the nucleus? (AS - 1) (1 Mark)
A: * The electrons do not follow definite paths around the nucleus.
* They revolve around the nucleus in a region called orbital.
11. What is the velocity of the electron? (AS - 1) (1 Mark)
A: The velocity of the electron is very nearer to the velocity of light.
12. Is it possible to find the exact position of the electron? (AS - 1) (2 Marks)
A: * Electrons are invisible to the nacked eye. As the electrons are very small, light of very short wavelength is required to find the position of the electron.
* This short wavelength light interacts with the electron and disturbs the motion of the electron.
* Hence, simultaneously the position and velocity of electron cannot be measured accurately.
* We can note the region of space around the nucleus where the probability of finding the electron is maximum.
13. Do atoms have a definite boundary, as suggested by Bohr's model? (AS - 1)(1 Mark)
A: Atoms do not have a definite boundary as suggested by Bohr's model.
14. What do we call the region of space where the electron might be, at a given time? (AS - 1) (1 Mark)
A: The region of space around the nucleus where the probability of finding the electron is maximum is called an orbital.
15. What information do the quantum numbers provide?
A: * The quantum numbers describe the space around the nucleus where the electrons are found and also their energies.
* These are called atomic orbitals.
16. What does each quantum number signify? (AS - 1) (2 Marks)
A: * The principal quantum number provides information about the size and energy of the main shell.
* The angular-momentum quantum number provides information related to the shape of the sub-shell.
* The magnetic quantum number describes the orientation of the orbital in space relative to the other orbitals in the atom.
17. What is the maximum value of 'l' for n = 3? (AS - 1) (1 Mark)
A: The maximum value of l is 3 for n = 3
18. How many values can 'l' have for n = 4? (AS - 1) (1 Mark)
A: * l can take the values of 0 to (n - 1).
* So for n = 4, l can take values of 0 to 3. It means 0, 1, 2, 3. (4 values)
19. Do all the p - orbitals have the same energy? (AS - 1) (1 Mark)
A: * All the p - orbitals have the same energy.
* They differ in their orientation only.
20. How are the two electrons in Helium atom arranged? (AS - 1) (1 Mark)
A: * These two electrons are arranged in pair in 1s orbital.
* So the electronic configuration for the atom is 1s2.
21. What are the spins of two electrons in a He atom? (AS - 1) (1 Mark)
A: * The two electrons in a He atom have anti-parallel spins.
* Electrons with paired spins are denoted by .
22. How many electrons can occupy an orbital? (AS - 1) (1 Mark)
A: An orbital can hold only two electrons and they must have opposite spins.
23. For Carbon atom (C) (Z = 6) where does the 6^{th} electron go? (AS - 1) (1 Mark)
A: * The configuration of Carbon atom (C) (Z = 6) is 1s^{2} 2s^{2} 2p^{2}.
* The 5^{th} and 6^{th} electrons into separate 2p orbitals, with both electrons having the same spin.
ACTIVITIES
Activity: 1
1. Describe an activity to explain the wave nature of light. (AS - 3) (4 marks)
A: * When we throw a stone into a still pond, we observe ripples, which are transmitting the disturbance in the form of waves on the surface of water.
* Sound waves are produced when something vibrates.
* In the same way electromagnetic waves are produced when an electric charge vibrates.
* A vibrating electric charge creates a change in electric field. The changing electric field creates a changing magnetic field.
* The process continues with both the created fields being perpendicular to each other and at right angles to the direction of propagation of the wave.
Activity: 2
2. Describe an activity to establish the fact that metals produce colour inflame. (AS - 3) (4 Marks)
A: * Take a pinch of cupric chloride in a watch glass and make a paste with concentrated hydrochloric acid.
* Take this paste on a platinum loop and introduce it into a non-luminous flame.
* Cupric chloride produces a green colour flame.
* Take a pinch of strontium chloride in a watch glass and make a paste with concentrated hydrochloric acid.
* Take this paste on a platinum loop and introduce it into a non-luminous flame.
* Strontium chloride produces a crimson red flame.
* Thus the above activities establish the fact that metals produce colour in flame.
Activity: 3
3. Complete the electronic configuration of the following elements. (AS - 3)(4 Marks)
Element | Atomic Number Z |
Electronic configuration |
C | 6 | |
N | 7 | |
O | 8 | |
F | 9 | |
Ne | 10 | |
Na | 11 | |
Mg | 12 | |
Al | 13 | |
Si | 14 | |
P | 15 | |
S | 16 | |
Cl | 17 | |
Ar | 18 | |
K | 19 | |
Ca | 20 |
A:
S. No. | Element | Atomic number (Z) | Electronic configuration |
1. | Carbon (C) | 6 | 1s^{2 }2s^{2} 2p^{2} |
2. | Nitrogen (N) | 7 | 1s^{2} 2s^{2} 2p^{3} |
3. | Oxygen (O) | 8 | 1s^{2} 2s^{2} 2p^{4} |
4. | Fluorine (F) | 9 | 1s^{2} 2s^{2} 2p^{5} |
5. | Neon (Ne) | 10 | 1s^{2} 2s^{2} 2p^{6} |
6. | Sodium (Na) | 11 | 1s^{2 }2s^{2} 2p^{6} 3s^{1} |
7. | Magnesium (Mg) | 12 | 1s^{2} 2s^{2} 2p^{6} 3s^{2} |
8. | Aluminium (Al) | 13 | 1s^{2} 2s^{2} 2p^{6} 3s^{2} 3p^{1} |
9. | Silicon (Si) | 14 | 1s^{2} 2s^{2} 2p^{6} 3s^{2} 3p^{2} |
10. | Phosphorous (P) | 15 | 1s^{2} 2s^{2} 2p^{6} 3s^{2} 3p^{3} |
11. | Sulphur (S) | 16 | 1s^{2} 2s^{2} 2p^{6} 3s^{2} 3p4 |
12. | Chlorine (Cl) | 17 | 1s^{2} 2s^{2} 2p^{6} 3s^{2} 3p^{5} |
13. | Argon (Ar) | 18 | 1s^{2} 2s^{2} 2p^{6} 3s^{2} 3p^{6} |
14. | Potassium (K) | 19 | 1s^{2} 2s^{2} 2p^{6} 3s^{2} 3p^{6} 4s^{1} |
15. | Calcium (Ca) | 20 | 1s^{2} 2s^{2} 2p^{6} 3s^{2} 3p^{6} 4s^{2} |
Additional Questions and Answers
I. Conceptual Understanding
1. How many colours are there in a rainbow? (1 Mark)
A: There are seven colours namely Violet, Indigo, Blue, Green, Yellow, Orange and Red (VIBGYOR) in a rainbow.
2. When are electromagnetic waves produced? (1 Mark)
A: Electromagnetic waves are produced when an electronic charge vibrates.
3. What is the speed of an electromagnetic wave? (1 Mark)
A: Visible light is an electromagnetic wave and the speed of light is 3 × 10^{8} ms^{-1}.
4. What are the characteristics of electromagnetic waves? (4 Marks)
A: * Electromagnetic energy travelling through a vacuum behaves in some way like ocean waves travelling through water.
* Like ocean waves, electromagnetic energy is characterized by wavelength (λ) and frequency.
* The wavelength (λ) of the wave is the distance from one wave peak to the next.
* The frequency (υ) of a wave is simply the number of wave peaks that pass by a given point per unit time, expressed in units of reciprocal seconds (1/s or s^{-1}).
* The relation between these quantities is given by λα or c = υλ
5. What is an electromagnetic spectrum? Give a familiar example for it. (2 Marks)
A: * Electromagnetic waves can have a wide variety of frequencies.
* The entire range of electromagnetic wave frequencies is known as the electromagnetic spectrum.
* The familiar example of the visible spectrum in nature is the formation of a rainbow.
6. Explain the term 'visible spectrum'. (2 Marks)
A: * Each colour in a rainbow is characterized by a specified wavelength from red (higher wavelength) to violet (shorter wavelength).
* These colours (wavelengths), that the human naked eye is sensitive to, are called visible light.
* The range of wavelengths covering red colour to violet colour is called the visible spectrum.
7. Describe an electromagnetic spectrum. (2 Marks)
A: * Electromagnetic waves can have a wide variety of wavelengths. The entire range of wavelengths is known as the electromagnetic spectrum.
* The electromagnetic spectrum consists of a continuous range of wavelengths of gamma rays at the shorter wavelength to radiowaves at the longer wavelength.
* But our eyes are sensitive only to visible light.
8. Explain Max Planck's theory. (4 Marks)
A: * Max Planck broke with the 'continuous energy' tradition of electromagnetic energy by assuming that the energy is always emitted in multiples of hυ.
For example: hυ, 2hυ, 3hυ, ......, nhυ.
* That is, the energy for a certain frequency E can be represented by the equation E = hυ, where 'h' is Planck's constant which has the value 6.626 × 10^{-34} Js and 'υ' is the frequency of the radiation absorbed or emitted.
* The energy (E) for the red colour (higher wavelength or lower frequency) is lower compared to the energy of blue colour (lower wavelength or higher frequency).
* The energy emitted from a material body increases with increase in heat energy.
* The significance of Planck's proposal is that, electromagnetic energy can be gained or lost in discrete values and not in a continuous manner.
* Hence, emission or absorption of light spectrum is a collection of a group of wavelengths.
9. Explain Bohr's model of Hydrogen atom and its limitations. (4 Marks)
A: Bohr's model of Hydrogen atom:
* Electrons in an atom occupy stationary orbits (states) of fixed energy at different distances from the nucleus.
* When an electron 'jumps' from a lower energy state (ground state) to higher energy states (excited state) it absorbs energy or emits energy when such a jump occurs from a higher energy state to a lower energy state.
* The energies of an electron in an atom can have only certain values E_{1}, E_{2}, E_{3}, ......... that is, the energy is quantized.
* The states corresponding to these energies are called stationary states and the possible values of the energy are called energy levels.
Limitations of Bohr's model:
a) Bohr's model explains all the line spectra deserved in the case of hydrogen atom. It is a successful model as for as line spectra of hydrogen atom is concerned.
b) Bohr's model failed to account for splitting of line spectra.
c) This model failed to account for the atomic spectra of atoms of more than one electron.
10. Explain Bohr-Sommerfeld model of an atom? What are its limitations? (4 Marks)
A: Bohr-Sommerfeld model of an atom
The allowed electronic orbits for the main quantum numbers by Bohr - Sommerfeld model
* In an attempt to account for the structure (splitting) of line spectra known as fine spectra, Sommerfeld modified Bohr's atomic model by adding elliptical orbits.
* While retaining the first of Bohr's circular orbit as such, he added one elliptical orbit to Bohr's second orbit, two elliptical orbits to Bohr's third orbit, etc. such that the nucleus of the atom is one of the principal foci of these elliptical orbits.
* He was guided by the fact that, in general, periodic motion under the influence of a central force will lead to elliptical orbits with the force situated at one of the foci.
Limitations:
* Bohr-Sommerfeld model, though successful in accounting for the fine line structure of hydrogen atomic spectra, does not provide a satisfactory picture of the structure of atom in general.
* This model failed to account for the atomic spectra of atoms of more than one electron.
11. Do the electrons follow definite paths around the nucleus? (2 Marks)
A: * If the electron revolves around the nucleus in definite paths or orbits, the exact position of the electron at various times will be known.
* Electrons are invisible to the naked eye. As the electrons are very small, light of very short wavelength is required to know the position and velocity of electron.
* This short wavelength of light interacts with the electron and disturbs the motion of the electron.
* Hence simultaneously the position and velocity of electron cannot be measured accurately.
* So the electrons do not follow definite paths around the nucleus.
12. Distinguish between orbit and orbital. (4 Marks)
A:
Orbit | Orbital |
* The path of an electron around the nucleus is called orbit. | * The region in space where there is finite probability of finding electron is called atomic orbital. |
* It is two dimensional. | * It is three dimensional. |
* Its shape is circular. | * It has spherical or dumb bell or double dumb bell shape. |
* The maximum number of electrons in any orbit is 2n. | * The maximum number of electrons in orbital is '2'. |
13. What does principal quantum number signify? (4 Marks)
A: Principal quantum number (n)
* The principal quantum number is related to the size and energy of the main shell.
* 'n' has positive integer values of 1, 2, 3, ........
* As 'n' increases, the shells become larger and the electrons in those shells are farther from the nucleus.
* An increase in 'n' also means higher energy. n = 1, 2, 3, ....... are often represented by the letters K, L, M, ....... For each 'n' value there is one main shell.
Shell | K | L | M | N |
n | 1 | 2 | 3 | 4 |
14. Write a note on the angular momentum quantum number. (4 Marks)
A: The angular - momentum quantum number (l)
* 'l has integer values from 0 to n - 1 for each value of 'n'. Each 'l' value represents one sub - shell.
* Each value of 'l' is related to the shape of a particular sub - shell in the space around the nucleus.
* The value of 'l' for a particular sub - shell is generally designated by the letters s, p, d, ......... as follows:
l | 0 | 1 | 2 | 3 |
Name of the sub-shell | s | p | d | f |
* When n = 1, there is only one sub - shell with l = 0. This is designated as 1s orbital.
* When n = 2, there are two sub - shells, with l = 0, the '2s' sub - shell and with l = 1, the '2p' sub - shell.
15. Explain the significance of the magnetic quantum number. (4 Marks)
A: The magnetic quantum number (ml):
* The magnetic quantum number (ml) has integer values between -l and +l, including zero. Thus for a certain value of l, there are (2l + 1) integer values of ml as follows:
-l, (-l + 1), ........, -1, 0, 1, .........., (+l - 1), + l
* These values describe the orientation of the orbital in space relative to the other orbitals in the atom.
* When l = 0, (2l + 1) = 1 and there is only one value of ml, thus we have only one orbital i.e. 1s.
* When l = 1, (2l + 1) = 3, that means ml has three values, namely, -1, 0 and 1 or three p orbitals, with different orientations along x, y, z axes.
These are labelled as p_{x}, p_{y} and p_{z}.
16. Do all the p - orbitals have the same energy? (2 Marks)
A: * The number of 'ml' values indicates the number of orbitals in a sub - shell with a particular l value.
* Orbitals in the sub - shell belonging to the same shell possess same energy.
17. Give in a tabular form the number of orbitals present in each sub shell. (2 Marks)
A:
18. Show in a tabular form the maximum number of electrons than can occupy various sub -shells. (4 Marks)
A:
19. Why is spin quantum is introduced? (2 Marks)
A: * The three quantum numbers n, l, and ml describe the size (energy), shape, and orientation respectively of an atomic orbital in space.
* As you have observed in the case of street lights (sodium vapour lamp), yellow lights emitted.
* This yellow light is comprised of a very closely spaced doublet when analyzed using high resolution spectroscope.
* Alkali and alkaline earth metals show such type of lines.
* To account for such a behaviour of electron an additional quantum number is introduced. This is spin quantum number.
* This represents the property of the electron. It is denoted by 'm_{s}'.
20. What does the spin quantum number refer? Explain its importance. (2 Marks)
A: * This quantum number refers to the two possible orientations of the spin of an electron, one clockwise and the other anticlockwise spin.
* These are represented by + and - .
* If both are positive values, then the spins are parallel otherwise the spins are anti - parallel.
* The importance of the spin quantum number is seen when electrons occupy specific orbitals in multi - electron atoms.
21. What is electronic configuration. How do you represent the electronic configuration of Hydrogen atom? (4 Marks)
A: Electronic configuration: The distribution of electrons in shells, sub-shells and orbital in an atom is known as electronic configuration.
Electronic configuration of hydrogen atom:
* The shorthand notation consists of the principal energy level (n value), the letter representing sub-level (l value), and the number of electrons (x) in the sub-shell is written as a superscript as shown below:
nl^{x}
* For the hydrogen (H) atom having atomic number (Z) = 1, the number of electrons is one, then the electronic configuration is 1s^{1}.
* The electron configuration can also be represented by showing the spin of the electron.
* For the electron in H, as we have seen, the set of quantum numbers is:
n = 1, l = 0, ml = 0, m_{s} = or - .
22. State and explain Pauli exclusion principle. According to this principle how many electrons can occupy an orbital? (4 Marks)
A: Pauli principle: No two electrons of the same atom can have all four quantum numbers the same.
Explanation:
* Helium atom has two electrons. The first electron occupies '1s' orbital. The second electron joins the first in the 1s - orbital, so the electron configuration of the ground state of 'He' is 1s^{2}.
* According to Pauli exclusion principle no two electrons of the same atom can have all four quantum numbers the same.
* If n, l and ml are same for two electrons then ms must be different. In the helium atom the spins must be paired.
* Electrons with paired spins are denoted by . One electron has ms = + , the other has ms = - . They have anti - parallel spins.
* The major consequence of the exclusion principle involves orbital occupancy. Since only two values of ms are allowed, an orbital can hold only two electrons and they must have opposite spins.
Hence, the electronic configuration of Helium atom is : 1s^{2}
23. How is Aufbau principle helpful in filling up the orbitals to build up the electronic configuration? (4 Marks)
A: Aufbau principle
* As we pass from one element to another one of next higher atomic number, one electron is added every time to the atom.
* The maximum number of electrons in any shell is '2n^{2}', where 'n' is the principal quantum number.
* The maximum number of electrons in a sub - shell (s, p, d or f) is equal to 2 (2l + 1 ) where l = 0, 1, 2, 3... Thus these sub-shells can have a maximum of 2, 6, 10, and 14 electrons respectively.
* In the ground state the electronic configuration can be built by placing electrons in the lowest available orbitals until the total number of electrons added is equal to the atomic number. This is called the Aufbau principle The German word Aufbau means building up.
* Thus orbitals are filled in the order of increasing energy.
* Two general rules help us to predict electronic configurations.
a) Electrons are assigned to orbitals in order of increasing value of (n + l)
b) For sub-shells with the same value of (n + l), electrons are assigned first to the sub-shell with lower 'n'.
24. State and explain Hund's rule taking the configuration of Carbon as an example. (2 Marks)
A: Hund's rule
* According to this rule electron pairing in orbitals starts only when all available empty orbitals of the same energy (degenerate orbitals) are singly occupied.
* The configuration of Carbon (C) atom (Z = 6) is 1s^{2} 2s^{2} 2p^{2}. The first four electrons go into the 1s and 2s orbitals. The next two electrons go into separate 2p orbitals, with both electrons having the same spin.
* Note that the unpaired electronics in the 2p orbitals are shown with parallel spins
V. Communication through Drawing, Model Making
25. Draw the diagram of an electromagnetic wave and label its parts. (4 Marks)
A:
26. Draw the diagram of an electromagnetic spectrum and note the wavelengths of different radiations. (4 Marks)
A:
27. Draw the shapes of orbitals in s and p sub - shells. Mention their shapes. (4 Marks)
A: s - orbital is spherical in shape, p-orbital is dumbell-shaped as shown below.
28. Draw the shapes of orbitals in d sub - shell. Mention their shape. (4 marks)
A: d - orbitals is double dumbell - shaped as shown below.
29. Show the electronic configuration of some elements in the increasing atomic number (Z) value. (4 Marks)
A: The electronic configuration of some elements in the increasing atomic number
(Z) value is given below.
30. Draw the Moeller chart diagram. Mention the ascending order of energies of various atomic orbitals. (4 Marks)
A: The following diagram shows the increasing value of (n + l). Ascending order of energies of various atomic orbitals is given below.
1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s < 4f < 5d < 6p < 7s < 5f < 6d < 7p < 8s ......
31. Write the table which represents the shells, sub - shells and the number of orbitals in the sub -shells. (4 Marks)
A:
Writer: C.V. Sarveswara Sarma