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Everything in the universe is made up of atoms, from the tiniest drop of water to the vast universe. But can you imagine what an atom looks like? Can you see beyond that tiny structure? Well, some scientists did. This chapter Structure of Atom takes us on a trip to see the structure of that tiny atom which some brilliant scientists have already discovered, and some have even looked beyond the atoms, into the electrons, protons and neutrons.
This chapter contains concepts like atomic theories, quantum numbers, photoelectric effect etc., that are too crucial to understand. The NCERT notes for Class 11 Chemistry chapter 2 will act as a key guide to grasp these concepts effectively. These NCERT notes offer a structured approach to help you excel in your exams. The atomic models like the Thomson model of atom, the Rutherford model, and the Bohr model of atom are well explained in the notes. We have also included formulas and diagrams to provide you clear picture. Follow these notes for feasible learning!
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NCERT Solution for class 11 chemistry chapter 2 |
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The notes of this chapter contain all the concepts in detail that will be enough for you to attempt the questions. The diagrams discussed here will help you visualise the concepts and the formulas will help you tackle the numericals.
The structure of the atom was obtained from the experiments on electrical discharge through gases.
The discovery of the electron began with Michael Faraday's work in the 1830s, where he found that electricity passing through solutions caused chemical changes, suggesting that electricity has tiny particles.
Later, in the 1850s, scientists studied electrical discharges in partially evacuated cathode ray tubes. These glass tubes had metal plates (electrodes) inside, and when high voltage was applied at low pressure, a stream of particles moved from the negative electrode (cathode) to the positive electrode (anode) - these were called cathode rays. When these rays passed through a hole in the anode and hit a zinc sulphide screen, they created a glowing spot, showing the presence of particles, which were later identified as electrons.
Properties of Cathode Rays
They travel in a straight line.
Cathode rays begin from the cathode and then move towards the anode.
They are invisible but can be made visible with the help of materials like fluorescent or Phosphorescent.
Since under the supply of electric charge, they are moved to the positive charge, which indicates that cathode rays consist of negatively charged particles.
The properties of cathode rays do not depend upon the material of the electrode in which it is used and the nature of gases that are present in the cathode ray tube.
Experimentally, J.J. Thomson determined the charge-to-mass ratio of the electrons. According to Thomson’s experiment, the amount of deviation for a particle from its path under the presence of an electric and a magnetic field depends upon
Deflection tends to be higher when the magnitude of the charge is higher for the particles.
When the mass of a particle is lighter, the deflection will be greater.
With the increase of voltage across the electrodes for with the increase of the strength of the magnetic field the deflection of electrons will also rise.
Thomson determined the value,
Where the me is the mass of the electron and e is the magnitude of charge.
The oil drop experiment conducted by R.A. Millikan found the charge on the electron.
Charge of an electron=-1.6022×10-19C
Mass of electron = 9.1094×10-31kg
Goldstein conducted another experiment with the help of a perforated cathode ray tube. A new type of eraser passes through the perforation of the card order by reducing the pressure, and it moves in just the opposite direction to that of the cathode rays. These are named canal rays or anode rays.
Properties of anode rays
The magnitude of the positive charge of anode rays depends upon the nature of the gas that is present in the tube.
It also depends on the gas for the value of the charge-to-mass ratio.
The behavior of anode rays is just the opposite of cathode rays and the magnetic and electric fields.
The smallest and lightest positive ion was obtained from hydrogen and was called a proton. This positively charged particle was characterised in 1919.
Later, the neutral particles were discovered by Chadwick (1932) by bombarding a thin sheet of beryllium with α-particles. When electrically neutral particles having a mass slightly greater than that of protons were emitted. He named these particles as neutrons.
A series of atomic models were introduced that led to the discovery of atoms and their subatomic particles.
Name |
Symbol |
Absolute Charge (C) |
Relative Charge |
Mass (kg) |
Mass (u) |
Approx. Mass (u) |
Electron |
e |
–1.602176 × 10⁻¹⁹ |
–1 |
9.109382 × 10⁻³¹ |
0.00054 |
0 |
Proton |
p |
+1.602176 × 10⁻¹⁹ |
+1 |
1.6726216 × 10⁻²⁷ |
1.00727 |
1 |
Neutron |
n |
0 |
0 |
1.674927 × 10⁻²⁷ |
1.00867 |
1 |
J.J. Thomson proposed a structure for an atom that can be regarded as a sphere of an approximate radius carrying a positive charge due to protons, and in which the negatively charged electrons are embedded. Thereby, the atom can be visualized as a pudding for a cake of positively charged protons with electrons in it. And the mass of atoms is evenly spread over the atom.
Rutherford's
A stream of high energy α–particles from a radioactive source was directed at a thin foil (thickness ∼ 100 nm) of gold metal. The thin gold foil had a circular fluorescent zinc sulphide screen around it. Whenever α–particles struck the screen, a tiny flash of light was produced at that point
Observations of Rutherford's experiment:
The observations of the Rutherford Alpha scattering experiment are
Most of the part of the atom is empty and atom is spherical in shape.
Each atom consists of a small, heavy, positively charged portion located at the centre, known as nucleus.
All positive charges of an atom (i.e protons) are present in the nucleus and electrons move around the nucleus in circular orbits.
Electrons and the nucleus are held together by electrostatic forces of attraction.
Atomic Number
The number of protons present in the nucleus is referred to as the atomic number. For example, the number of protons in a sodium atom is 11 and also the atomic number of sodium is 11. For maintaining electrical neutrality the number of electrons in an atom is equal to the number of protons.
Mass Number
The sum of the number of protons and neutrons present in a nucleus together is called the mass number.
The element that has the same atomic number but a different mass number is an isotope. 1H1, 1H2, 1H3.
And the element that is processing the same mass number but a different atomic number is an isobar. 6C14, 7N14.
Historically, results obtained from studies of radiation interactions with matter have yielded enormous information about the structure of atoms and molecules. Niels Bohr used these findings to refine the model presented by Rutherford. Two developments played a major role in the formulation of Bohr’s model of atom.
Radiation causes wavelike and particle-like properties, which means that it has a dual character
The atomic spectra can be explained only by assuming the quantum state electronic energy levels in atoms.
Electromagnetic radiation -
According to electromagnetic wave theory, energy is emitted continuously from a source in the form of radiation (or waves), known as electromagnetic radiation. Electromagnetic radiations have both magnetic field as well as electric field components which oscillate in the phase perpendicular to each other as well as perpendicular to the direction of wave propagation. These waves do not require any medium for propagation and can propagate through a vacuum. There are many types of electromagnetic radiations which constitute what is known as the electromagnetic spectrum. There are several parameters required to characterise or define a wave. These parameters are defined below:
1. Wavelength (
The maxima are called as Crests and the minima are called as Troughs. Alternatively, the distance between two consecutive crests or two consecutive troughs is also called as the wavelength.
2. Time Period (T): It is the time required for one complete oscillation or one complete cycle by a wave.
3. Frequency (
4. Speed (c): It is the distance travelled by the wave in one second. In one time period, the wave travels a distance equal to its wavelength.
The speed of all the different components of light is the same i.e. they travel with the speed of 3
5. Wave number (
The rays present on the left extreme of the spectrum have the greatest frequency, the least wavelength, and the greatest energy. As the frequency increases, wavelength decreases and the energy increases.
Max Planck put forward a theory for explaining the phenomenon of blackbody radiation and the photoelectric effect. The theory focuses on
The radiant energy absorbed or emitted is in the form of small pockets of energy and these small pockets of energy are quantum.
The energy of each Quantum is directly proportional to the frequency of the radiation emitted.
The ideal body that has the ability to emit and absorb all frequencies is a black body. And the corresponding radiation emitted by a black body is black body radiation.
where h is Planck's constant and it has a value equal to 6.63
Photoelectric Effect
When a beam of light passes on the surface of some metals, electrons are emitted from the metal surface. This phenomenon is the photoelectric effect. And it has been observed that only photons of light of a particular frequency, that is the threshold frequency, can cause the photoelectric effect. The kinetic energy of the electron emitted from the surface of the metal is directly proportional to the frequency of the striking photons. And also when the intensity of the photon of the light is increasing, more electrons are ejected.
There were certain observations in the photoelectric effect experiment.
(1) There was a requirement of a minimum energy for each metal for the photoelectric effect to occur. This minimum energy is known as work function (W0) and it can be closely associated with the ionisation energy of the metal.
Mathematically, the work function, threshold frequency and threshold wavelength can be associated as
Note: hc is approximately equal to 2 x 10-25 J-m or 12400 eV-nm. (eV is the energy in electron volts)
(2) The number of electrons ejected is proportional to the intensity (brightness) of light striking the metal, but does not depend upon the frequency of light.
(3) There was almost no time lag between the striking of light and ejection of photo electrons
(4) The kinetic energy of the ejected electrons (photoelectrons) depends upon the frequency of the light used.
Einstein's photoelectric equation
From conservation of energy
where m is the mass of the electron and v is the velocity associated with the ejected electron. Also, h is Planck’s constant and v0 is threshold frequency of metal.
(5) The Kinetic energy of ejected photoelectron is also sometimes associated with Stopping Potential. It is defined as the minimum opposing potential applied due to which the kinetic energy of electron becomes zero.
where,
Vs = Stopping potential
e = Charge on electron
Dual behavior of Electromagnetic Radiation
Light and other electromagnetic radiation consist of a dual nature. That is, they have particle and wave-like properties. Radiations emitted by the matter show the properties of a particle by exhibiting propagation. Electrons also exhibit wave-particle duality.
When a light ray passes through a prism the wave which is having a short wavelength forms a colored band more than the one that has a longer wavelength. These bands then spread to form a series of colored bands and are called the spectrum. The one that is deviated least is the one which is having the longest wavelength, that is, their red color.
Continuous spectrum
When white light is passed through a prism, it will split into seven different colored bands, just like a rainbow and these colors are continuous and it is called the continuous spectrum.
Emission spectra
When radiation is emitted from a source and is then passed to a prism, which is then received on a photographic plate it is called the emission spectrum. The emission spectrum can be observed by heating a substance to a high temperature.
Line spectra
The emission spectra of atoms in the gas phase do not show a continuous spread of wavelength from red to violet, rather they emit light only at specific wavelengths with dark spaces between them. Such spectra are called line spectra or atomic spectra because the emitted radiation is identified by the appearance of bright lines in the spectra. Sodium emits yellow light while potassium emits violet light.
Absorption spectra
A ray of light, when allowed to pass through the vapors of a substance and the transmitted light is then allowed to strike a prism, dark lines appear. The dark line represents that the radiations corresponding to them are absorbed by the substance. The spectrum is called the absorption spectrum.
The figure below is showing absorption and emission spectra.
Line Spectrum of Hydrogen
When an electric discharge is passed through gaseous hydrogen, the H2 molecules dissociate, and the energetically excited hydrogen atoms produced emit electromagnetic radiation of discrete frequencies. The spectrum, consisting of a large number of lines, is obtained and the spectrum is called the hydrogen spectrum. The series of lines is named as the Lyman series, Balmer series, Paschen series, Brackett series, and Pfund series.
1. The electron in the hydrogen atom can move around the nucleus in a circular path of fixed radius and energy. These paths are called orbits, stationary states, or allowed energy states and are arranged concentrically around the nucleus. The force of attraction between the nucleus and an electron provides the centripetal force required by the electron to carry out the circular motion.
2. The energy of an electron in the orbit does not change with time. However, the electron will move from a lower stationary state to a higher stationary state when a required amount of energy is absorbed by the electron or energy is emitted when the electron moves from a higher stationary state to a lower stationary state
3. Energy can be absorbed or emitted when electron transitions between two different orbits and the frequency of the photon involved can be calculated using the formula:
4. The angular momentum of an electron is quantised. In a given stationary state, it can be expressed as
L= mvr= nh/2
So, only those energy states (or orbits) are allowed in which the above equation holds true for the angular momentum.
Note: Bohr's model is only valid for hydrogen-like species or unielectronic species that contain only a single electron
The line spectrum of the hydrogen atom can be explained using Bohr's model, which says that electrons absorb or emit energy when they jump between orbits (energy levels). When an electron moves from a lower orbit (ni) to a higher one (nf), energy is absorbed. When it moves from a higher orbit to a lower one, energy is emitted as light. The energy difference between two levels is given by:
Where
And the wavenumber (lines in the spectrum) is
This formula matches the one given by Rydberg earlier using experimental data. Each spectral line in hydrogen corresponds to a specific electron transition. In a large sample of hydrogen atoms, many such transitions occur, producing multiple lines. The brightness of a line depends on how many photons of the same energy are emitted or absorbed.
It fails to account for the finer details (doublet, that is, two closely spaced lines) of the hydrogen atom spectrum observed by using spectroscopic techniques.
This model is also unable to explain the spectrum of atoms other than hydrogen, for example, He.
Further, Bohr’s theory was also unable to explain the splitting of spectral lines in the presence of a magnetic field (Zeeman effect) or an electric field (Stark effect). It could not explain the ability of atoms to form molecules by chemical bonds.
Two important developments that contributed significantly to the formulation of such a model are
The scientist de-Broglie proposed that just like light, matter should exhibit both particle and wave-like properties. This means that just as the photon has momentum as well as wavelength, electrons should also have momentum as well as wavelength, and he proposed the following mathematical relationship
where m is the mass of the particle, v is its velocity, p is the momentum, and K.E. is the Kinetic Energy of the particle.
V is the voltage across which the charged particle having charge q is accelerated.
(2) de Broglie’s prediction was confirmed experimentally when it was found that an electron beam undergoes diffraction, a phenomenon characteristic of waves.
(3) It needs to be noted that, according to de Broglie, every object in motion has a wave character. This wavelength is quite significant for the subatomic particles, which have very small masses. The wavelengths associated with ordinary objects are however so short that their wave properties cannot be detected as they have large masses.
It states that it is impossible to determine simultaneously the exact position and exact momentum (or velocity) of an electron.
If an attempt is made to ensure that any one of these two quantities is measured with a higher accuracy, then the other quantity becomes less accurate.
Mathematically, the product of uncertainty in position(Δx) and uncertainty in momentum(ΔP) is equal to or greater than h/4π
It can be proved mathematically that the uncertainty principle is only significant for subatomic particles but not for large-sized objects.
Theoretical science that deals with the study of the motion of microscopic objects that have both particle-like and wave-like properties.
The energy of an electron in an atom is in a quantized state.
The existence of quantised electronic energy levels is a direct result of the wave-like properties of electrons and are allowed solutions of the Schrödinger wave equation.
The exact position and exact velocity of an electron in an atom cannot be found out simultaneously.
An atomic orbital is the wave function ψ for an electron in an atom.
The square of the orbital wave function |ψ|2 is the probability of finding an electron.
For a system (such as an atom or a molecule whose energy does not change with time), the Schrödinger equation is written as
As the atomic energy levels or the orbits are quantized, they can be expressed in terms of quantum numbers. Quantum numbers are;
Principal quantum number(n)
It represents the principal shell of an atom. It can have integral values except zero, like 1,2,3,.... Also denoted as K, L, M, etc.
The maximum number of electrons in a principal shell can be 2n2, where n is the principal quantum number.
This quantum number gives information about :
Where Z is the atomic number and n is the principal quantum number.
Azimuthal quantum number(l):
The azimuthal quantum number represents the subshell or subenergy shell in an atom.
l has values from 0 to (n-1).
For eg: for n = 2 ; l = 0, 1
Subshell notation for 0, 1 is s and p.
No. of electrons [2(2l+1)]: for s subshell l = 2; for p subshell, l = 6.
Magnetic quantum number(m):
It represents the number of orbitals present in a subshell.
m has values ranging from -l to +l, including zero.
For eg, for ‘s’ subshell :
The value of l is 0
m has value = 0
For ‘p’ subshell :
The value of l is 1
m has value -1, 0, +1
Spin quantum number(s):
Electrons in an orbital can spin either clockwise or anticlockwise.
Thus, an electron can have only two possible values of this quantum number, either
According to the German physicist, Max Born, the square of the wave function (i.e.,ψ 2) at a point gives the probability density of the electron at that point.
The region where this probability density function reduces to zero is called nodal surfaces or simply nodes.
Shape of s orbital: spherical
In general, it has been found that the ns-orbital has (n – 1) nodes, that is, the number of nodes increases with the increase of principal quantum number n.
Shape of p orbital: Dumb-bell
Shape of d orbital: Double Dumb-bell
The total number of nodes are given by (n–1), i.e., sum of l angular nodes and (n – l – 1) radial nodes.
The orbitals having the same energy are called degenerate. The 1s orbital in a hydrogen atom, as said earlier, corresponds to the most stable condition and is called the ground state and an electron residing in this orbital is most strongly held by the nucleus.
An electron in the 2s, 2p or higher orbitals in a hydrogen atom is in an excited state.
Within a given principal quantum number, the energy of orbitals increases in the order
Aufbau Principle -
According to this rule, “orbitals are filled in the increasing order of their energies starting with the orbital of lowest energy”. Energy of various orbitals are compared with (n+l) rule.
The orbitals having a lower value of (n+l), has lower energy.
If the value of (n+l) is the same for two orbitals, then the orbital with the lower value of ‘n’ would have lower energy and be filled first.
For eg: 3p and 3d.
For 3p : n=3, l=1 so n+l= 4
For 3d : n=3, l=2 so n+l=5. Thus energy of 3p is lower than that of 3d.
Pauli Exclusion Principle -
“No two electrons in an atom can have the same set of four quantum numbers".
Hund’s Rule of Maximum Multiplicity -
The pairing of electrons in the orbitals belonging to the same subshell ( p,d, or f ) does not take place until each orbital belonging to that subshell has got one electron each i.e. it is singly occupied.
Writing of electronic configuration of any element is based on three rules. They are:
Aufbau principle
Pauli’s exclusion principle
Hund’s rule of maximum multiplicity
The distribution of electrons into orbitals of an atom is called its electronic configuration.
The electronic configuration of the different elements can be represented in two ways:
Subshell notation
Orbital diagram
Let’s understand with the help of examples:
The hydrogen atom has only one electron, which goes in the orbital with lowest energy, namely 1s. The E.C. of hydrogen is 1s1 (subshell notation)
The electronic configuration of lithium is 1s2 2s1. Lithium has 3 electrons as its atomic number is 3. 2 electrons are filled in the 1s orbital, and 1 electron is filled in the 2s orbital.
let us consider fluorine (Z = 9) :
F(Z = 9) = 1s2, 2s2, 2px2, 2py2, 2pz1 or
The importance of knowing the exact electronic configuration of an element lies in the fact that the chemical properties of an element are dependent on the behavior and relative arrangement of its electrons.
The completely filled or half-filled subshells have a symmetrical distribution of electrons in them and are therefore more stable.
Symmetrical distribution of electrons: It is well known that symmetry leads to stability. The completely filled or half-filled subshells have a symmetrical distribution of electrons in them and are therefore more stable.
Exchange Energy: The stabilizing effect arises whenever two or more electrons with the same spin are present in the degenerate orbitals of a subshell. These electrons tend to exchange their positions and the energy released due to this exchange is called exchange energy. The number of exchanges that can take place is maximum when the subshell is either half-filled or completely filled. As a result, the exchange energy is maximum and so is the stability.
For example, the valence electronic configurations of Cr and Cu are 3d54s1 and 3d104s1, respectively and not 3d44s2 and 3d94s2.
Question: The extra stability of a half-filled subshell is due to
(A) Symmetrical distribution of electrons
(B) Smaller coulombic repulsion energy
(C) The presence of electrons with the same spin in non-degenerate orbitals
(D) Larger exchange energy
(E) Relatively smaller shielding of electrons by one another
Identify the correct statements
(1) (B), (D) and (E) only
(2) (A), (B), (D) and (E) only
(3) (B), (C) and (D) only
(4) (A), (B) and (D) only
Answer: The extra stability of a half-filled subshell is due to the symmetrical distribution of electrons, smaller coulombic repulsion, larger exchange energy, and smaller shielding of electrons.
Hence, the correct answer is option (2).
Question: Which one of the following about an electron occupying the 1s orbital in a hydrogen atom is incorrect? (Bohr's radius is represented by
(1). The probability density of finding the electron is maximum at the nucleus
(2) The electron can be found at a distance
(3) The 1s orbital is spherically symmetrical
(4) The total energy of the electron is maximum when it is at a distance
Answer:
1.
2. Electrons can exist up to infinity from the nucleus. Hence, statement 2 is true.
3. The 1s orbital is spherically symmetrical- it is true
4. The energy of an electron is maximum at an infinite distance from the nucleus.
More further the electron is from the nucleus, the more is its energy. The total energy of an electron is given by
So as
Hence, statement 4 is false.
Hence, the correct answer is option (4).
Question: Which of the following sets of quantum numbers are correct?
n l m1
(i) 1 1 +2
(ii) 2 1 +1
(iii) 3 2 -2
(iv) 3 4 -2
Choose from the options given below
1) (i) and (ii)
2) (ii) and (iii)
3) (ii) and (iii)
4) (i) and (iv)
Answer:
As in the case n = 1, then I ≠ 1, therefore option (i) is incorrect.
In case n = 3, I ≠ 4, therefore option (iv) is also incorrect.
In case n = 2, I = 0, 1. When l = 1, then m = -1, 0, +1. Therefore, option (ii) is correct.
In case n = 3, l = 0, 1, 2. When l = 2, m = -2, -1, 0, +1, +2. Therefore, option (iii) is also correct.
Hence, the correct answer is option (2).
To solve the questions effectively, students have to focus on understanding the fundamental concepts, along with practicing questions. Students should utilise the sources and content provided online to ensure a complete understanding of the subject.
Here are a few tips that help students to solve the questions with a good approach:
1. Understand the concepts
Firstly, start with a thorough reading of the NCERT textbook.
2. Keep revising the important concepts
The atomic models are one of most important topics of this chapter. Learn their findings and limitations
3. Focus on the important topics like
Concepts like the line spectrum of hydrogen and quantum numbers are frequently asked in exams. Prepare them well.
4. Summarise the concepts
Make a summary sheet with formulas and quantum numbers. Also, practice electron configurations for elements up to atomic number 30.
5. Solve the questions
Try to solve the in-text and end-of-exercise problems provided in the NCERT textbook. Questions from the NCERT books are asked directly in the NCERT boards and other competitive exams. Do previous year questions from NEET and JEE to get used to question patterns.
NCERT notes for each chapter of class 11 are given below. Follow them for deep learning.
Follow the links below to get chapter-wise NCERT solutions and make your learning better.
Students can refer to the links given below for the NCERT exemplar subject-wise solutions
Students can also follow NCERT solutions for other subjects to ace their exam preparations. Click on the link below
NCERT Solutions for Class 11 Maths |
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NCERT Solutions for Class 11 Chemistry |
NCERT Solutions for Class 11 Biology |
Also read
Heisenberg's Uncertainty Principle states that it is fundamentally impossible to determine simultaneously, and with absolute precision, both the position and the momentum (or velocity) of a subatomic particle like an electron. If you know the position very accurately, your knowledge of its momentum becomes uncertain, and vice versa. Mathematically: Δx ⋅ Δp ≥ h/4π.
Hund's Rule states that for a given subshell (orbitals of equal energy or degenerate orbitals, like the three p-orbitals or five d-orbitals), electrons will first occupy all the orbitals individually with parallel spins (same spin direction) before any orbital is doubly occupied (paired up). This maximizes the total spin multiplicity and leads to a more stable configuration.
The principal quantum number (n) describes the energy level or shell of an electron in an atom, with higher values of 'n' indicating higher energy levels and larger distances from the nucleus.
The principal quantum number (n) describes the energy level or shell of an electron in an atom, with higher values of 'n' indicating higher energy levels and larger distances from the nucleus.
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