Class 11th Structure of Atom Notes

We will learn about atoms and their structure in the Class 11th notes in chapter 2 “Structure of Atom”. Chemistry is the study and comprehension of three particles, regardless of how they react with other elements or how they combine.   

Our Structure of Atoms class 11 pdf elaborately explains the 3 fundamental particles, i.e., electron, neutron, and protons. These are the basic components of an atom. The word ‘atom’ has been derived from the Greek word ‘atoms’, which means ‘individual’. Let us have a look at the structure of atoms in class 11 chapter 2 notes for self-study.  

Sub-atomic Particles  

The atom, according to Dalton’s Atomic Theory, is the ultimate particle of matter. It correctly described several chemical laws, such as the law of mass conservation, the law of constant composition, and the law of multiple proportions. However, it could not explain the existence of subatomic particles such as electrons and protons, which were eventually discovered.  

Discovery of Electrons  

William Crooks investigated the conduction of electricity through low-pressure gases in 1879. He experimented with a discharge tube, which is a 60-cm-long cylindrical hard glass tube. It has two metal electrodes on both ends and is sealed on both ends.  

It was noticed that only at very low pressures and very high voltages could the electrical discharge through the gases.   

Evacuation could adjust the pressure of various gases. Current flows through a stream of particles traveling in the tube from the negative electrode (cathode) to the positive electrode (anode) when a sufficiently high voltage is put across the electrodes (anode). They were called Cathode rays, or cathode ray particles. 

Properties of Cathode Rays

1. The cathode rays follow a straight path.  

2. Cathode rays travel from the cathode to the anode.  

3.  Although these rays are invisible, their behavior can be detected through the use of particular materials (fluorescent or phosphorescent) that illuminate when they are struck.  

4.  Negatively charged particles make up cathode rays. When a pair of metal plates are used to apply an electric field to cathode rays, they are deflected towards the positive plate, showing the existence of a negative charge.  

5.  The material of the electrodes and the type of gas present in the cathode ray tube have no bearing on the characteristics of cathode rays.  

Charge to Mass Ratio  

The ratio of electrical charge to the mass of the electron was measured by JJ Thomson by using a cathode ray tube and applying magnetic and electrical fields perpendicular to each other, as well as the path of the electrons. According to Thomson, the amount of deviation of the particles from their path depends upon a certain condition.   

1. Magnitude: The magnitude of the charge of the particle is proportional to the interaction with electric and magnetic fields, which shows greater deflection.   
2. Mass: Lesser the less mass of the particle, the greater the deflection.  
3. Strength of electric or magnetic field: The deflection of electrons is directly proportional to the voltage across electrons or the strength of the electric field.  

When only an electric field is applied, the electrons hit the cathode ray tube at point A. Whereas if only a magnetic field is applied, the electrons hit the cathode ray tube at point C. In the case of the absence of both, the electron hits at point B.  

On accurate measurement, the amount of deflection observed by electrons, e/me = 1.758820 x 1011 C kg-1 where me = Mass of the electron in kg. and e = magnitude of the charge on the electron in coulomb (C)  

Discovery of Anode Rays

In 1886, Goldstein added a perforated cathode to the discharge tube. When the pressure was reduced, he noticed a new form of luminous rays passing through the cathode’s holes or perforations and moving in the opposite direction of the cathode rays. These rays were named positive rays, anode rays, and canal rays. Anode rays are emitted from a space between the anode and the cathode, not from the anode itself.  

Properties of Anode rays  

1. The composition of the gas in the discharge tube determines the value of positive charge (e) on the particles that make up anode rays.
2.  It is discovered that the gas from which they emerge affects the charge-to-mass ratio of the particles.  
3. A multiple of the fundamental unit of electrical charge is carried by some positively charged particles.   
4.  In a magnetic or electric field, the behavior of these particles is the polar opposite of that of electrons or cathode rays.  

Proton  

The proton is the smallest and lightest positive ion created from hydrogen.  

Mass of proton = 1.676 x 10-27 kg  

Charge on a proton = (+) 1.602 x 10-19 C  

Neutron  

A neutron is a neutral particle. Chadwick was the one who discovered it (1932).  

He discovered that highly penetrating rays comprise neutral particles called neutrons by bombarding tiny sheets of beryllium with rapidly moving a-particles.  

Thomson Model of Atom 

1. J. J. Thomson recommended that an atom be thought of as a sphere with a radius of about 1CT8 cm with a positive charge because of protons and negatively charged electrons buried within it.
2. In this model, the atom is depicted as a positive-charged pudding or cake with electrons contained within it. In this model, the atom is depicted as a positive-charged pudding or cake with electrons contained within it.  
3. According to this model, the atom’s mass is uniformly distributed across the atom.

Failure Thomson Model of Atom

Although this concept could explain the atom’s general neutrality, it could not fully explain Rutherford’s scattering tests from 1911.  

Rutherford’s a-particle Scattering Experiment  

n 1911, Rutherford experimented with scattering by bombarding thin foils of metals such as gold, silver, platinum, or copper with a beam of fast-moving a-particles. A circular fluorescent zinc sulphide film around the thin gold foil. A brief flash of light was produced whenever a-particles collided with the screen.  

He made the following observations because of his experiments:  

1. Most of the a-particles passed through the foil without deflection  
2.  A few a-particles were deflected at modest angles.  
3. Only a handful simply deflected backward, approximately 180 degrees.  

Rutherford drew the following conclusions based on his observations:  

1. There must be enough space within the atom because the majority of the a-particles passed through the foil without being deflected.  
2.  Small angles deflected a small fraction of a-particles. The positive charge had to be concentrated in a very tiny region so that a few positively charged a-particles were resisted and deflected. The nucleus was the name given to this relatively small part of the atom.  
3.  Compared to the entire volume of an atom, the nucleus has a very modest volume.  

Rutherford’s Nuclear Model of an Atom 

1. The positive charge and most of the atom’s mass were squeezed into a very tiny area. Rutherford used the term “nucleus” to describe this tiny part of the atom.  
2.  The nucleus is surrounded by electrons that circulate about it in circular patterns called orbits at a top speed.  
3. Electrostatic forces of attraction hold electrons and nuclei together. 

Atomic Number  

The atomic number is equal to the number of protons in the nucleus (z). The number of protons in the hydrogen nucleus is 1, whereas the number of protons in the sodium atom is 11, hence their atomic numbers are 1 and 11. The number of electrons in an atom must be equal to the number of protons to maintain electrical neutrality (atomic number, z). One is the number of electrons in a hydrogen atom, while eleven is the number of electrons in a sodium atom.  

the number of protons in an atom’s nucleus. = Number of electrons in a neutral atom. Atomic Number (z) = the number of protons in an atom’s nucleus.  

Mass Number  

Nucleons are the number of protons and neutrons in the nucleus as a whole. The total number of nucleons in an atom is known as the mass number (A).  

Number of protons (p) + Number of neutrons (n) = Mass Number (A) (n).

Isotopes  

Atoms of the same element having the same atomic number but different mass numbers are called isotopes

Isobars:  

Atoms of distinct elements having the same mass number but different atomic numbers are known as isobars.  

Developments Leading to Bohr’s Model of Atom  

The formulation of Bohr’s atom model was influenced by two major advances. These were:  

1. Electromagnetic radiation has a dual character, meaning it has both wave and particle features.  
2.  Atomic spectrum experimental results can only be explained by assuming quantized electronic energy levels in atoms.  

Electromagnetic Wave Theory  

In 1864, James Clerk Maxwell proposed this hypothesis. The following are the main points of this theory:  

1. Radiant energy is energy that is continuously emitted as radiation from any source (such as a heated rod or the filament of a bulb through which electric current is conducted).  
2. The radiation comprises perpendicular oscillating electric and magnetic fields that are both perpendicular to the radiation’s propagation direction.  
3. The radiation has wave characteristics and travels at a speed of 3 x 108 m/sec.  
4. There is no need for a material medium for the propagation of these waves. The sun’s beams, for example, travel through space, which is a non-material medium.  

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Characteristics of Wave  

Wavelength: The distance between any two consecutive crests or troughs is known as the wavelength. It is denoted by X, and its SI unit is the meter.  

Frequency of wave: The number of waves traveling through a place in one second is known as the frequency of a wave. It is denoted by the v (nu) and is measured in Hertz (Hz).  

1 Hz equals 1 cycle per second.  

Velocity: The linear distance traversed by a wave in one second is referred to as its velocity. It is denoted by c and is measured in cm/sec or m/sec.  

Aptitude: The height of the crest or the depth of the through determines the amplitude of a wave. It is denoted by V and is measured in length units.  

Wave Number: The number of waves present at a 1-meter length is referred to as the wave number. It will be equal to the wavelength’s reciprocal. It is denoted by bar v (read as a nu bar).  

Black Body Radiation

A black body is an ideal body that emits and absorbs all frequencies, and the radiation emitted by such a body is known as black body radiation. The frequency distribution of a black body’s emitted radiation is solely determined by its temperature.  

At a given temperature, the intensity of radiation emitted increases as the wavelength decreases, achieves a maximum value at a specific wavelength and then begins to drop as the wavelength decreases further.  

Planck’s Quantum Theory  

Max Planck proposed Planck’s Quantum Theory in 1900 to explain the occurrence of ‘Black Body Radiation’ and the photoelectric effect.  

In 1905, Einstein pushed this theory even further. The following were the primary points of this theory:  

1.  Radiant energy is emitted or absorbed as little packets of energy. Each of these energy packets is referred to as a quantum.  
2. The frequency of the radiation is precisely proportional to the energy of each quantum.  

E = hv = hc / λ  

Where h is a proportionality constant, it is called Planck’s constant. Its value is equal to 6.626 x 10-34 J sec.

Photoelectric Effect 

n 1887, Hertz found that when a beam of light of a specific frequency affects the surface of certain metals, electrons are released or ejected. This phenomenon is known as the photoelectric effect.  

Observations in Photoelectric effect  

1. The photoelectric effect can only be caused by photons of light with a specified minimum frequency, known as the threshold frequency (v0). The value of v0 varies depending on the metal.  
2. The kinetic energy of the released electrons is directly proportional to the affecting photons’ frequency and is completely independent of their intensity.  
3. The number of electrons ejected each second from the metal surface is determined by the intensity, not the frequency, of the affecting photons or radiation. 

Explanation of Photoelectric Effect 

Using Planck’s quantum theory, Einstein could explain the various aspects of the photoelectric effect in 1905:  

1. Photoelectrons are only ejected if the incident light has a specific frequency (threshold frequency v0)  

2. If the incident light frequency (v) exceeds the threshold frequency (v0), the excess energy (hv − hv0) is transferred to the electron as kinetic energy. K.E. is the expelled electron energy of the emitted electron.

3. More electrons are expelled as the intensity of light increases, but their energies remain the same.

Dual Behavior of Electromagnetic Radiation  

Scientists discovered that light and other electromagnetic radiation had a dual nature after studying their behaviour. There are two types of nature: waves and particles. When radiation interacts with matter, it takes on particle-like characteristics, as opposed to wave-like characteristics (interference and diffraction) when it propagates. This wave-particle duality can also be found in small particles like electrons.  

Spectrum  

When passing a ray of white light through a prism, the wave with the shorter wavelength bends more than the wave with the longer wavelength. Because typical white light comprises waves with all the visible wavelengths, an array of white light is divided into a spectrum of colored bands. The light of the red color has the longest wavelength and is deviated the least, while the light of the violet color has the shortest wavelength and is deviated the most.  

Continuous Spectrum  

When a ray of white light is examined via a prism, it is discovered that it breaks into seven distinct wide bands of colors ranging from violet to red (like a rainbow). These colors are so similar that they blend into one another. As a result, the spectrum is referred to as a continuous spectrum.  

Spectra of Emission  

When the radiation emitted from a source is passed via a prism and then received on the photographic plate, they are known as Emission Spectra. Radiation can be produced in a variety of ways, including:  

1. from the sun or a brightly lit electric bulb  
2.  by sending a low-pressure electric discharge through a gas.  
3.  by raising the temperature of a substance.  

Line Spectra  

When the vapors of a volatile material are allowed to fall on the flame of a Bunsen burner, they are then analyzed using a spectroscope. On the photographic plate, several distinct colored lines appear, which vary depending on the substance. Sodium and its salts, for example, emit yellow light, but potassium and its salts emit violet light.  

Absorption Spectra  

When white light passes through a substance’s vapors and the transmitted light strikes a prism, black lines form in the otherwise continuous spectrum. The black lines show that the white light’s equivalent radiation was absorbed by the material. The absorption spectrum is the name for this type of spectrum.  

The dark lines appear at the same locations as the lines in the emission spectra.

Line Spectrum of Hydrogen  

The spectrum comprises a huge number of lines that are grouped into various series when an electric discharge is transmitted through hydrogen gas confined in a discharge tube under low pressure and the produced light is studied by a spectroscope. The hydrogen spectrum refers to the entire spectrum. Johannes Rydberg discovered that every series of lines in the hydrogen spectrum could be explained by the following expression based on experimental data.  

Bohr’s Model of Atom 

In 1913, Niels Bohr presented a new atom model based on Planck’s Quantum Theory. The following are the primary features of this model:  

i) The electrons in an atom orbit around the nucleus in precise circular trajectories, termed orbits.  

(ii) Because each orbit is related to a specific amount of energy, these are referred to as energy orbits.  

energy shells or levels These are numbered 1, 2, 3, 4……….. or K, L, M, N………..  

(iii) For the electron, only those energy orbits are permitted, in which the electron’s angular momentum is a whole integer multiple of h/2 π  

Angular momentum of electron (mvr) = nh/2π (n = 1, 2, 3, 4 etc.).  

m = mass of the electron.  

v = tangential velocity of the revolving electron.  

r = radius of the orbit.  

h = Planck’s constant.  

n is an integer.  

(iv) As long as an electron is present in a specific orbit, it does not absorb or lose energy, and hence, its energy remains constant.  

(v) An electron absorbs energy only in fixed amounts as quanta when energy is supplied to it and then leaps to a higher energy level away from the nucleus known as the excited state. Because the excited state is unstable, the electron may return to a lower energy state, emitting the same amount of energy. (∆E = E2–E1).  

Achievements of Bohr’s Theory 

1. The stability of an atom has been explained by Bohr’s theory.  

2. The energy of one electron in a hydrogen atom and one electron species has been calculated using Bohr’s theory.  

3. The atomic spectrum of the hydrogen atom has been explained by Bohr’s hypothesis.  

Limitations of Bohr’s Model  

1. The theory could not explain the atomic spectra of multi-electron atoms or atoms with over one electron. 

2. The fine structure of the spectral lines was not explained by Bohr’s hypothesis.

3. Bohr’s theory could not convincingly explain the Zeeman effect and Stark effects. 

4. Bohr’s theory could not explain how atoms can form molecules through chemical bonds. 

 5. It violated Heisenberg’s principle of uncertainty. 

Dual Behavior of Matter  

De Broglie postulated in 1924 that matter, like radiation, should have dual behavior, i.e., particle and wave qualities. This means that electrons, like photons, have both velocity and wavelength. 

De Broglie derived the following relationship between the wavelength (λ) and momentum (p) of a material particle from this example. 

ƛ = h/mc=h/p 

Where, m= mass of particle 

               v= velocity of the particle 

                p= momentum of the particle 

This equation is called the de Broglie formula. 

Heisenberg’s Uncertainty Principle 

It states that “It is impossible to determine simultaneously, the exact position and exact momentum (or velocity) of an electron”. 

If uncertainty in position= Δ x 

And uncertainty in momentum= Δp 

When both are measured simultaneously, according to this formula, ∆ x.∆ p≥h/4π? 

Quantum Mechanical Model of Atom

Quantum mechanics is a theoretical science that studies the motions of small objects with both observable wave and particle features. 

Quantum Numbers 

Atomic orbitals can be described by quantizing their related energy and angular momentum (i.e., they have specific values). Quantized values are quantized values that can be stated as a quantum number. These are used to obtain all the information about an electron, including its position, energy, and spin. 

Principal Quantum Number

It is the most essential quantum number since it identifies the electron’s primary energy level or shell. The letter v represents it and can have any integral value other than zero, e.g., n = 1, 2, 3, 4, …………etc.   

The letters K, L, M, N, O, P………. Beginning with the nucleus

Azimuthal or Subsidiary or Orbital Angular Quantum Number

It is discovered that the spectra of the elements contain not only the major lines but also many fine lines. This number aids in the spectrum’s understanding of fine lines.  

The information provided by the azimuthal quantum number is:  

1. In the main shell, the number of subshells.  
2. The angular momentum of each electron is present in any subshell.  
3. The relative energies of several subshells   
4. The forms of the many subshells that make up a single main shell.  

The letter T stands for this quantum number. It can be any value between 0 and n –1 for a given value of n.  

Magnetic Orbital Quantum Number

The number of preferred electron orientations in a subshell is determined by the magnetic orbital quantum number. The magnetic orbital quantum number determines the number of orbitals present in any subshell since each orientation corresponds to one orbital.  

The letter m or ml represented the magnetic quantum number, and it can take any value of l from–l to + l, including zero.  

As a result, m has 2l + 1 value for the energy value of l.  

Spin Quantum Number

The magnetic characteristics of the substances can be explained using this quantum number. The behavior of a spinning electron is like that of a micro magnet with a defined magnetic moment. When two electrons share an orbital, their magnetic moments oppose and cancel each other.  

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Shape of Atomic Orbitals  

Shapes of s-orbitals

The s-subshell contains the s-orbital. l = 0 and ml = 0 for this subshell. As a result, a spherical s-orbital with only one orientation has uniform electron density along all three axes.  

The chance of an Is electron is determined to be highest near the nucleus and diminishes as the distance from the nucleus increases. The chance of a 2s electron is also highest near the nucleus and declines to zero near the nucleus. Nodal surface, or simply node, is the spherical empty shell for 2s electrons.  

Shape of p orbitals  

p-orbitals are present in the p-subshell with l = 1 and m1 can have three different orientations – 1, 0, + 1. The p-subshell has three orbitals, which are labeled as px, py, and pz orbitals depending on which axis they are pointed along. A p-orbital’s shape is that of a dumbbell, with two lobes. 

Shape of d-orbitals

In a d-subshell with l = 2 and m [= -2, -1, 0, +1, and +2], d-orbitals are present. This shows that there are five separate orbitals for each of the five orientations. d orbitals are of five types: dxy, d_xy, d_yz, d_x²_-y² 

Electronic Configuration

The electronic configuration is the distribution of electrons in distinct orbitals. The electrons in orbitals must follow the following guidelines:
• Aufbau’s Principal
• Pauli’s Exclusion Principal
• Hund’s Rule of Maximum Multiplicity

Aufbau’s Principle  

According to this theory, the lowest-energy orbitals are filled first, followed by the high-energy orbitals.  

1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s < 4f < 5d < 6p < 7s < 5f < 6d < 7p  

(n + l) rule (Bohr Bury’s Rule)  

According to this, the orbital which has a lower value of (n + l) is lower in energy.  

Pauli’s Exclusion Principle

According to this concept, no two electrons in an atom have the identical value of all four quantum numbers, except a maximum of two electrons with the opposite spin.  

Hund’s Rule of Maximum Multiplicity 

This rule states that electrons should be distributed throughout the orbitals of a subshell in such a way that the largest number of unpaired electrons with a parallel spin is got.  

Structure of Atom Class 11 Questions and Answers