Structure Of Atom

Question 1

State Dalton’s Atomic Theory.

Answer 1

The atomic theory of matter was first proposed by John Dalton and he called it Dalton’s atomic theory.

1.Matter consists of indivisible atoms.

2.All atoms of a given element have identical properties, including identical mass. Atoms of different elements differ in mass.

3. Compounds are formed when atoms of different elements combined in a fixed ratio.
4. Chemical reactions involve reorganisation of atoms. They are neither created nor destroyed in a chemical reaction.

Question 2

Drawbacks of Dalton’s Atomic Theory?

Answer 2

1. Could not explain the law of gaseous volumes.
2. Could not provide a reason for combining of atoms.
3. Did not explain the existence of isomers and isobars.
4. Contradicts the existence of nuclear reactions which leads to destruction of atoms.
5. Does not explain existence of complex organic compounds (do not form by combining in whole no. ratios).
6. Does not explain existence of sub-atomic particles.

Question 3

Explain discovery of electron and Faraday’s cathode ray experiment.

Answer 3

Michael Faraday showed that if electricity is passed through a solution of electrolyte, chemical reactions occurred at the electrodes which resulted in the liberation and deposition of matter at the electrodes.

Faraday began to study electrical discharge in partially evacuated tubes known as cathode ray discharge tubes. The electrical discharge through the gases could be “*observed only at very low pressure and at very high voltage**. It was noted that **current starts flowing through a stream of particles moving in the tube from the negative electrode (cathode) to the positive electrode (anode, these were called cathode rays or cathode ray particles.

Results:
1. The cathode rays start from the cathode and move towards the anode.
2.These rays themselves are not visible but their behaviour can be observed with the help of certain kinds of materials such as fluorescent or phosphorescent and with glow when hit by them.
3. In the absence of electrical or magnetic field these rays travel in straight lines.
4. In the presence of electrical or magnetic field the behaviour of cathode rays are similar to that expected from negatively charged particles, suggesting that; cathode ray particles are actually negatively charged particles called electrons.
5.The characteristics of cathode rays (electrons) do not depend upon the material of electrodes and the nature of gas present in the cathode. Which means; electrons are basic constituents of atoms.

5. & 6. lead to discovery of electron.

Question 4

Explain and JJ Thomson’s experiment and how he measured charge to mass ratio of electron.

Answer 4

JJ Thomson measured the ratio of electrical charge to the mass of electron by using cathode ray tube and applying electrical and magnetic field perpendicular to each other as well as to the path of electrons.

The results were:
1. The magnitude of the negative charge on the particle: greater the magnitude of the charge on the particle,greater is the interaction with the electrical or magnetic field and greater is the deflection.
2. The mass of the particle: lighter the particle, greater the deflection.
3. The strength of the electrical or magnetic field: the deflection of electrons from its original path increases with the increase in the voltage across the electrodes, or the strength of the magnetic field.

He determined the ratio e/m(e) as:
e/m = 1.758820 × 10¹¹ C/kg.

Remember, electrons are negatively charged and therefore their charge is also negative.

Question 5

How was charge of an electron determined?

Answer 5

The charge of an electron was determined by the oil drop experiment performed by R.A. Milikan. He found the charge on the electron to be:
** -1.6 × 10-¹⁹ C**

Question 6

How was the mass of the electron determined?

Answer 6

The mass of the electron was determined after the discovery of charge to mass ratio of electron and the charge of the electron.

m(e) = e/(e/m(e)) = 9.1 × 10-³¹ kg

Question 7

Explain discovery of Protons and Neutrons.

Answer 7

The electrical discharge carried out in the modified cathode tube led to the discovery of canal rays carrying positively charged particles.

Characteristics:
1.Unlike cathode rays, mass of positively charged particles depend upon the nature of gas present in the cathode rays, these are simply the positively charged gaseous ions.
2.The charge to mass ratio of the particles depends upon the gas from which these originate.
3. Some of the positively charged particles carry a multiple of the fundamental unit of electrical charge.
4. The behaviour of these particles in the magnetic or electrical field is opposite to the observed for electron or cathode rays.

The smallest and lightest positive ion was obtained from hydrogen and was called proton.

Later, the need was felt for the presence of electrically neutral particles. These particles were discovered by Chadwick by bombarding a thin sheet of beryllium by Alpha particles, when electrically neutral particles having a mass slightly greater than that of protons were emitted, he named these particles as neutrons.

Question 8

Explain the discovery of neutrons.

Answer 8

After the discovery of protons and electrons, the need was felt for the presence of electrically neutral particles. These particles were discovered by Chadwick by bombarding a thin sheet of beryllium by Alpha particles, when electrically neutral particles having a mass slightly greater than that of protons were emitted, he named these particles as neutrons.

Question 9

Which is greater in mass: neutrons or protons?

Answer 9

Neutrons are very slightly greater in mass than protons. During calculations we take them approximately the same in mass.

Question 10

Explain JJ Thomson’s model of atom.

Answer 10

JJ Thomson’s model of atom is also known as the plum pudding or raisin pudding or watermelon model.

He proposed that an atom possesses a spherical shape (radius of approximately 10-¹⁰) in which the positive charge is uniformly distributed throughout the atom. he electrons are embedded into it such a manner as to give the most stable electrostatic arrangement.
An important feature of this model is that the mass of the atom is assumed to be uniformly distributed over the atom.
This model was able to explain the overall neutrality of the atom temporarily.

Question 11

Explain Rutherford’s model of atom and how he achieved it.

Answer 11

In Rutherford’s experiment, he bombarded very thin gold foil with Alpha particles, this is known as the famous Rutherford’s Alpha Particle Scattering Experiment. A stream of high energy Alpha particles from a radioactive source was directed at a thin foil of gold metal, the thin gold fall had a circular fluorescent zinc sulphide screen around it. When Alpha particles from the screen struck at any point, flash of light was produced at that point

Expectations:
According to Thomson model of atom, the mass of each gold atom in the foil should have been spread evenly over the entire atom. Alpha particles had enough energy to pass directly through such a uniform distribution of mass. It was expected that the particles would slow down and change directions only by small angles as they passed through the foil.

Observations:
1.Most of the Alpha particles passed through the gold foil undeflected.
2. A small fraction of the Alpha particles was deflected by small angles.
3. A very few Alpha particles (1 in 20,000) bounced back ,that is, were deflected by nearly 180°.

Conclusions:
1. Most of the space in the atoms is empty as most of the Alpha particles pass through the foil undeflected.
2. A few positively charge Alpha particles were deflected, the deflection must be due to enormous repulsive forces showing that the positive charge of the atom is not spread through out the atom as Thomson has presumed. The positive charge was to be concentrated in a very small volume that repelled and reflected the positively charged Alpha particles.
3. Calculations by Rutherford showed that the volume occupied by the nucleus is negligibly small compared to the total volume of the atom.

Rutherford’s Model:
1. The positive charge and most of the mass of the atom was densely concentrated in extremely strong region. This very small portion of the atom was called the nucleus.
2. The nucleus is surrounded by electrons that move around the nucleus with a very high speed in circular paths called orbits.
3. Electrons and the nucleus are held together by electrostatic forces of attraction.

Question 12

The charge on the proton is equal but opposite to that of?

Answer 12

Electron.

Question 13

The numbers of protons present in the nucleus is equal to?

Answer 13

Atomic no. (z) or no. of electrons in neutral atom.

Question 14

The number of electrons in an atom (neutral) is equal to?

Answer 14

No. of protons or atomic no. (z)

Question 15

Describe mass number.

Answer 15

Mass number is determined by no. of protons and neutrons, or by amount of nucleons.

mass no. = no. of protons (z) + no. of neutrons (n)

Question 16

What are nucleons?

Answer 16

Sub-atomic Particles present in nucleus (protons and neutrons) are known as nucleons.

Question 17

Define Isobars and Isotopes. Give examples.

Answer 17

**Isobars **are the atoms with the same mass number but different atomic number.
Ex: ¹⁴C6 and ¹⁴N7

Atoms with identical atomic number but different atomic mass number are known in Isotopes.
Ex: Protium (¹H1), Deuterium (²T1), Tritium (³T1)

Imp: chemical properties of atoms are controlled by the number of electrons which are determined by the number of protons in the nucleus number of neutrons present in the nucleus have very little affect on the chemical properties of an element therefore all the isotopes of a given element show same chemical behaviour.

Question 18

Do all isotopes of a given element show same chemical behaviour? Why?

Answer 18

Yes. Chemical properties of atoms are controlled by the number of electrons which are determined by the number of protons in the nucleus number of neutrons present in the nucleus have very little affect on the chemical properties of an element. Therefore, all the isotopes of a given element show same chemical behaviour.

Question 19

Drawbacks of Rutherford’s Model?

Answer 19

1. Fails to explain the stability of atom: When a body is moving in an orbit it undergoes acceleration even if it is moving with a constant speed because of changing direction. According to the Electromagnetic Theory of Maxwell, charged particles when accelerated should emit electromagnetic radiation. Therefore, an electron in a orbit will emit radiation, the energy carried by radiation comes from electronic motion. The orbit will the continue to shrink, calculations show that an electron should take only 10-⁸ secs to spiral into the nucleus but this does not happen. So the model should be wrong.
2. It says nothing about distribution of the electrons around the nucleus and the energy of these electrons.

Question 20

Which factors lead to development of Bohr’s Model?

Answer 20

1. Dual character of the electromagnetic radiation which means that the radiations processes both wave like and particle like properties.
2. Experimentql results regarding atomic spectra.

Question 21

Describe Maxwell’s discovery of Wave Nature of Electromagnetic Radiation.
Or, describe Wave Nature of Radiation.

Answer 21

James Maxwell suggested that when electrically charged particle moves under acceleration , alternating electrical and magnetic fields are produced and transmitted. These fields are transmitted in the forms of waves called electromagnetic waves or electromagnetic radiation.

He found that light waves are associated with oscillating electric and magnetic character.

MF is perpendicular to EF. This components have the same wavelength, frequency, speed and amplitude but they vibrate in two mutually perpendicular planes.

Characteristics:
1. The oscillating electrical and magnetic field produced by oscillating charged particles are perpendicular to each other and both are perpendicular to the direction of propagation of the wave.
2. Unlike sound waves or waves produced in water, electromagnetic waves do not require medium and can move in vacuum.
3. There are many types of electromagnetic radiations which differ from each another in wavelength or frequency, these constitute what is called electromagnetic spectrum.
4. Different kinds of units are used to represent electromagnetic radiations.
For example: SI unit of freq. (v) is Hz/s.
SI unit of wavenumber (V) is m-¹.
SI unit of wavelength (λ) is m.
5. In vacuum all types of electromagnetic radiation, regardless of wavelength, travel at the same speed (c): the speed of light.
c = 3 × 10⁸ m/s.

The frequency, wavelength and velocity of light are related by the equation: c = v λ

Question 22

Explain the Electromagnetic Spectrum.

Answer 22

There are many types of electromagnetic radiations which differ from each another in wavelength or frequency these constitute what is called electromagnetic spectrum.

The small portion around is 10¹⁵ Hz, is what is ordinarily called visible light. It is the only part of ES which our eyes can see or detect. Special instruments are required to detect non-visible radiation.

The Spectrum:
Radio - Micro - Infrared - Visible - UV - XRays - Gamma - Cosmic

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Question 23

What is the wave number?

Answer 23

The wave number is a commonly used quantity in spectroscopy (V). It is defined as the number of wavelengths per unit length, or the reciprocal of wavelength unit.
Unit: m-¹ or 1/m

Question 24

Give examples of some phenomena that were explained and not explained by the wave nature of electromagnetic radiation.

Answer 24

Explained:
1. Diffraction
2. Interference

Not Explained:
1. Nature of emission of radiation from hot bodies (black-body radiation).
2. Ejection of electrons from metal surface when radiation strikes it (**photoelectrical effect*”).
3. Variation of heat capacity of solids as a function of temperature.
4.Line spectra of Hydrogen with special reference to Hydrogen

Question 25

Hot objects emit electromagnetic radiations over a wide range of wavelengths.

Answer 25

At high temperatures, an appreciable proportion of radiation is in the visible region of the spectrum. **As the temperature is raised, the higher proportion of short wavelength (blue light) is generated**. This means that **red radiation is the most intense at a particular temperature and the blue radiation is the more intense at another temperature.** **This means intensity of radiations of different wavelength emitted by hard body depend upon its temperature. Objects made of different material and capital different temperatures different amount of radiation** **When the surface of an object is radiadated with light (electromagnetic radiation), a part of radiant energy is generally reflected as such, a part is observed and a part of it is transmitted.**

Question 27

True or False? When the surface of an object is radiadated with light (electromagnetic radiation), a part of radiant energy is generally reflected as such, a part is observed and a part of it is transmitted. State Reason.

Answer 27

True. The reason for incomplete absorption is that ordinary objects are as a rule imperfect absorbers of radiation.

Question 28

What is Black Body Radiation?

Answer 28

An ideal body which emits and absorbs radiation of all frequencies **uniformly** is called a **Back Body.** The radiation emitted by such a body is called **Black Body Radiation.** Black body is in thermal equilibrium with its surroundings. It radiates same amount of energy per unit area as it absorbs from its surroundings in any given time. The amount of light emitted from a black body and its spectral distribution depends only on its temperature. At a given temperature intensity of radiation emitted increases with the increase of wavelength, reaches a maximum value at a given wavelength and then starts decreasing with further increase of wavelength. **In practicality, No such body exists. Carbon Black from approximates fairly close to black body**.

Question 29

What is Planck's constant? How did he achieve it?

Answer 29

Max playing assumed at the absorption and emission of radiation arises from oscillator. i.e; atoms in the wall of black body. **Planck assumed that radiation could be subdivided into discrete chunks of energy. He suggested that atoms and molecules could emit or absorb energy only in discrete quantities and not in a continuous manner. He gave the name "Quantum" to the smallest quantity of energy that can be emitted or absorbed in the form of electromagnetic radiation.** He arrived at the relation that energy of a quantum of radiation is proportional to its frequency. Expression: **E = hv** Here, **h** is the **Planck's Constant**. **h = 6.626 × 10-³⁴ J s**

Question 30

Explain Quantization given by Planck.

Answer 30

Quantization has been compared to standing on a staircase, a person can stand on any step of a staircase but it is not possible to stand in between any two steps. **The energy can take any one of the values from the following set but cannot take on any values between them: E = {0,hv,2hv,3hvm..........,nhv,........}**

Question 31

Explain Photoelectrical Effect.

Answer 31

**Hurtz performed an experiment where electrons were ejected when certain metals (Rb, Cs, K) were exposed to a beam of light. This phenomena was called photoelectrical effect.** Results: 1. Electrons are ejected from the metal surface as soon as the beam of light strikes the surface. **There is no time lag between the striking of the light beam and the ejection of electrons from the metal surface. 2.**The number of electrons ejected is proportional to the intensity or brightness of light.** 3. For each metal there is a **characteristic minimum frequency (Vo = Vnot) also known as threshold frequency below which photoelectrical effect is not observed.** At a frequency V>Vo, the ejected electrons come out with certain kinetic energy. **The kinetic energy of these electrons increase with the increase of frequency of light used.** Drawbacks: **it was later proved that though the number of electrons ejected does depend on the brightness of light, the kinetic energy of the ejected electron does not.** Expression for kinetic energy of the ejected electron during photoelectrical effect: **hV = hVo + 1/2mv²** Where, V is freq., Vo is Threshold Frequency, m is mass of electro and v is velocity associated with ejected electron. **Expression for Threshold Frequency: Work function = Planck's constant • Threshold Frequency W = hVo**

Question 32

State T/F: 1. Number of electrons ejected in photoelectrical effect depend upon the brightness of light. 2. Number of electrons ejected in photoelectrical effect depend upon the frequency of light used. 3. The kinetic energy of electrons ejected during photoelectrical effect depend upon the brightness of light. 4. The kinetic energy of electrons ejected during the photoelectrical effective end upon the frequency of light.

Answer 32

1. True 2. False: Only depends on incident light to cross threshold frequency, after that doesn't matter. 3. False 4. True

Question 33

True or False? Number of electrons ejected in photoelectrical effect depend upon the frequency of light used. State Reason.

Answer 33

False: Only depends on incident light to cross threshold frequency, after that doesn't matter.

Question 34

What is Threshold Frequency.

Answer 34

For each metal, there is a **characteristic minimum frequency Vnot (Vo) also known as threshold frequency, below which photoelectrical effect is not observed.** So, **threshold frequency is the minimum frequency of light required for a particular metal to show photoelectrical effect.**

Question 35

Define Work function.

Answer 35

Work function is the minimum energy required to eject the electron during photoelectrical effect. (Wo = Wnot) *Wo = hVo** Where, Vo is Threshold Frequency.

Question 36

Give the expression for kinetic energy of the ejected electron during photoelectrical effect.

Answer 36

**hV = hVo + 1/2mv²** Where, V is freq., Vo is Threshold Frequency, m is mass of electro and v is velocity associated with ejected electron.

Question 37

When a light is pass through a prism the wave with _________ bends more than the one with ________ .

Answer 37

Shorter Wavelength, Longer Wavelength

Question 38

Explain continuous spectrum or spectrum of white light.

Answer 38

Since ordinary white light consist of waves with all the wavelengths in the visible range. **A ray of white light is spread out in a series of continuous coloured bands called continuous spectrum.**Continuous because violet into blue blue into green and so on. **The spectrum of white light or the visible range that we can see ranges from Violet at 7.5 × 10¹⁴ Hz to Red at 4 ×10¹⁴ Hz**. VIBGYOR

Question 39

Describe Bohr's Model for Hydrogen Atom.

Answer 39

Bohr was the first to explain quantitatively the general features of the structure of hydrogen atom and spectrum. He used planck's concept of quantization of energy. His postulates for the hydrogen atom's model were: 1. Electron in the hydrogen atom can move around the nucleus in a circular path of fixed radius and energy, these paths are called Orbits, Stationery States or Allowed Energy States. These orbits are arranged concentrically around the nucleus. 2. Energy of an electron in the orbit does not change with the time. However, the electron will move from a lower energy state to high energy state when required amount of energy is absorbed by the electron or energy is emitted when an electron moves from higher state to lower energy state. This energy change does not take place in a continuous manner. 3. The frequency of radiation absorbed or emitted when transition occurs between two stationery states that differ in energy by ∆E is given by: **n = ∆E/h = E2-E1/h** Where E1 and E2 are the energy of the lower and higher allowed energy States respectively. This expression is commonly known as **Bohr's frequency rule**. 4. Angular momentum of an electron is quantised. In an stationary state it can be expressed as: **mvr = n•(h/2π)** n = 1,2,3.... Where: m = mass of electron, v = velocity, r = radius of orbit. Thus an electron can move only in this orbits for which its angular momentum is integral multiple (h/2π), that means angular momentum is quantized.

Question 40

Explain Angular momentum for electron.

Answer 40

Angular momentum is the product of moment of inertia (l) and angular velocity (w). For an electron of mass m moving in a circular path of radius around the nucleus: **angular momentum = l × w** **angular momentum = mvr Angular momentum of an electron is quantised. In an stationary state it can be expressed as: **mvr = n•(h/2π)** n = 1,2,3.... Where: m = mass of electron, v = velocity, r = radius of orbit. **Thus an electron can move only in this orbits for which its angular momentum is integral multiple (h/2π)**, that means angular momentum is quantized.

Question 41

Explain derivation of energy of the Stationery States used by Bohr.

Answer 41

1.The Stationery States for electron are numbered n = 1,2,3... ,these integral numbers are known as **Principal Quantum numbers.** 2.The radii of the stationery states are expressed as: **r = n²a** Where, a = ao = another = 52.9pm **The radius of the first stationary state called the Bohr Orbit is 52.9 pm.** Normally the electron in the hydrogen atom is found in the orbit (that is n=1). As n increases, the value of r will increase. In other words, the electron will present further away from the nucleus. 3.The most important property associated with the electron is the energy of a stationary state, given by: E = -R(1/n²) n = 1,2,3.... Where, **R = R(h) = Rydberg Constant = 2.18×10-¹⁸ J**. Also, E = E(n) **The energy of the lowest state, also called the ground state.** 4. Bohr's theory can also be applied to ions containing only one electron similar to that present in hydrogen atom. The energy of the Stationery States associated with these kind of ions is given by the expression: **E = -R(h) [(Z²/n²)] J** And, Radius by: **r = [52.9(n²)/Z]** pm Where, z = atomic no. It's clear that the value of energy becomes more negative and that of radium becomes smaller with the increase of atomic number z. 5. It's also possible to calculate the velocity of electron moving in these orbits. Qualitatively the magnitude of velocity of electron increases with increase of positive charge on the nucleus and decreases with increase of principal quantum number n.

Question 42

What is Bohr Orbit?

Answer 42

Radius of the first stationary state is called the Bohr Orbit. Value: 52.9 pm

Question 43

Define Bohr's Frequency Rule.

Answer 43

**n = ∆E/h = E2-E1/h** Where E1 and E2 are the energy of the lower and higher allowed energy States respectively. This expression is Bohr's Frequency Rule. It gives the frequency of radiation absorbed or emitted when transition occurs between two stationery states that differ in energy by ∆E

Question 44

What Is Rydberg Constant?

Answer 44

Rydberg constant denoted as R(h) has the value 2.18×10-¹⁸ J. It's used in expression for energy of stationary state of electron: E = -R(h)[(1/n²)] n = 1,2,3....

Question 45

What is Ionized Hydrogen Atom?

Answer 45

When the electron is free from the influence of nucleus, the energy taken as zero. Electron in this situation is associated with the stationary state of principal quantum number **n = ∞** and is called as ionized hydrogen atom.

Question 46

What is an energy level diagram?

Answer 46

.

Question 48

The lowest energy state is also called?

Answer 48

Ground state.

Question 49

The energy gap between the two orbits can be given by?

Answer 49

∆E = Ef - Ei

Question 50

Radiation is absorbed if the electron moves from the orbit of smaller principle quantum number to be orbit of higher principle quantum number where is the radiation is emitted if the electron moves from higher orbit to lower orbit. True or False?

Answer 50

True

Question 51

The brightness or intensity of spectral lines depend upon the number of photons of same wavelength or frequency absorbed or emitted. True/False?

Answer 51

True

Question 52

Limitations of Bohr's Model?

Answer 52

1. It fails to account for the finer details (doublet, thatis two closely spaced lines) of the hydrogen atom spectrum observed by using sophisticated spectroscopic techniques. 2. This model is also unable to explain the spectrum of atoms other than hydrogen. 3. It is unable to explain this splitting of spectral lines in the presence of magnetic fields (Zeeman effect) or an electrical field (Stark effect). 4. It could not explain the ability of atoms to form molecules by chemical bonds.

Question 53

Explain the development which contributed to the formation of Quantum Mechanical Model of Atom.

Answer 53

1. **Dual Behaviour of Matter:** de Broglie proposed that just like radiation matter also possesses dual nature of wave and particle properties. This means that just as Photon has momentum as well as wavelength, electrons should also have momentum as well as wavelength. He gave the relation: **λ = h/(mv) = h/p** Where, m is mass of particle and v it's velocity and p it's momentum. De broglie's prediction was confirmed experimentally when it was founded that an electron beam undergoes diffraction, a phenomena characteristic of waves. 2.**Heisenberg's Uncertainty Principle:** states that it is impossible to determine simultaneously the exact position and exact momentum (or velocity) of an electron. Expression: **∆x × ∆p ≥ h/(4π) ∆x × ∆v ≥ h/(4π) ∆x × ∆(mv) ≥ h/(4π)** If the **position** of electron is known with high degree of accuracy then the **velocity** of the electron will be highly uncertain and vice versa.

Question 55

Explain De Broglie's Dual Nature of Matter.

Answer 55

**Dual Behaviour of Matter:** de Broglie proposed that just like radiation matter also possesses dual nature of wave and particle properties. This means that just as Photon has momentum as well as wavelength, electrons should also have momentum as well as wavelength. He gave the relation: **λ = h/(mv) = h/p** Where, m is mass of particle and v it's velocity and p it's momentum. De broglie's prediction was confirmed experimentally when it was founded that an electron beam undergoes diffraction, a phenomena characteristic of waves.

Question 56

Explain Heisenberg's Uncertainty Principle.

Answer 56

**Heisenberg's Uncertainty Principle:** states that it is impossible to determine simultaneously the exact position and exact momentum (or velocity) of an electron. Expression: **∆x × ∆p ≥ h/(4π) ∆x × ∆v ≥ h/(4π) ∆x × ∆(mv) ≥ h/(4π)** If the **position** of electron is known with high degree of accuracy then the **velocity** of the electron will be highly uncertain and vice versa. Significance: **It rules out existence of definite paths or trajectories of electrons and other similar particles**. We know that the position of an object and its velocity fix its trajectory. Since for subatomic object such as an electronic is not possible simultaneously to determine both the position and the velocity at any given instant to an arbitrary degree of precision,** it is not possible to talk of the trajectory of an electron**. Imp: **the effect of Heisenberg Uncertainty Principle is significant only for motion of microscopic object and is neglegible for that of macroscopic objects.** **In dealing with milligram-sized or heavier objects, the associated uncertainty's are hardly of any real consequence**.

Question 57

Reasons for failure of Bohr's Model.

Answer 57

1. The wave character of the electron is not considered in Bohr's Model. 2.Orbit is a clearly defined path which was proven to be not scientifically possible due to Heisenberg's Uncertainty Principle.

Question 58

Which branch of science talks about microscopic objects.

Answer 58

The branch of science that takes into account the dual behaviour of matter, the microscopic objects and their properties is called quantum mechanics.

Question 59

What is the Schrodinger Equation?

Answer 59

This Schrodinger equation talks about total energy for a system whose energy does not change with the time while considering kinetic energies of subatomic particles, etc. Equation: **HΨ = EΨ** Where, H = Hamiltonian. They Schrodinger equation can** not** be solved exactly for multi electron atom.

Question 60

What is the natural consequence in the solution of Schrodinger Equation?

Answer 60

The quantized energy state and corresponding wave functions which are characterized by set of three Quantum numbers: **Principle quantum no. Azimuthal quantum no. Magnetic quantum no.** arise as an consequence in the solution of the Schrodinger equation.

Question 61

The wave function Ψ is a mathematical function whose value depends upon the coordinates of the electron in the atom and it does not carry any actual physical meaning. True or False

Answer 61

True.

Question 62

The probability of finding an electron at a point within an atom is proportional to the ______ at that point.

Answer 62

|Ψ|²

Question 63

Define Azimuthal Quantum no.

Answer 63

Azimuthal Quantum no. (l): is **also known as orbital angular momentum or subsidiary quantum number**. It **defines the three dimensional shape of ths orbital**. It can have **values ranging from 0 to n-1**. Each shell consists of one or more sub-shell or sub-level. **The number of subshells in a principle shell is equal to the value of n**. Each sub-shell is assigned an azimuthal quantum number. Ex: for n = 3 there are three subshells; (l = 0,1,2).

Question 64

Define Orbitals and Quantum Numbers.

Answer 64

An orbital can hold a max of 2 electrons, with opposite spin. An orbital of smaller size means there is more chance of finding the electron near the nucleus. Similarly, shape and orientation mean that there is more probability of finding the electron among certain directions and along others. Atomic orbitals are precisely distinguished by what are known as quantum numbers each orbital is designated by three quantum numbers labelled as n, l and m. 1.Principal Quantum Number (n): The principal quantum number is a positive integer n = 1,2,3... **It determines the size, and to a large extent the energy of the orbital**. The principle quantum number **also identifies the shell**. **With increase in value of n, the number of allowed orbital increases and are given by n²**. All the orbitals of a given value of n constitute a single shell of atom. Energy of orbital will increase with increase in value of n. **Value of n ∝ Energy of Orbital ∝ Size/Radius of Orbital**. 2. Azimuthal Quantum no. (l): is **also known as orbital angular momentum or subsidiary quantum number**. It **defines the three dimensional shape of ths orbital**. It can have **values ranging from 0 to n-1**. Each shell consists of one or more sub-shell or sub-level. **The number of subshells in a principle shell is equal to the value of n**. Each sub-shell is assigned an azimuthal quantum number. Ex: for n = 3 there are three subshells; (l = 0,1,2). 3. Magnetic Quantum no. (m): the magnetic orbital quantum number **gives information about the special orientation of the orbital with respect to standard set of co-ordinate axis**. **For any subshell; 2l+1 values are possible**. 4. Spin Quantum no. (s): Goudsmit proposed the presence of 4th quantum number known as the electron spin quantum number. **An electron has besides charge in mass intrinsic spin angular momentum number.** This quantum number can **have two orientations/values: +1/2 or -1/2** These are called the two spin states of the electron. **An orbital cannot hold more than two electrons and these two electrons should have opposite spins**. Therefore, it can be noted that values of m are derived from values of l and values of l are derived from values of n.

Question 65

Define Principal quantum no.

Answer 65

1.Principal Quantum Number (n): The principal quantum number is a positive integer n = 1,2,3... **It determines the size, and to a large extent the energy of the orbital**. The principle quantum number **also identifies the shell**. **With increase in value of n, the number of allowed orbital increases and are given by n²**. All the orbitals of a given value of n constitute a single shell of atom. Energy of orbital will increase with increase in value of n. **Value of n ∝ Energy of Orbital ∝ Size/Radius of Orbital**.

Question 66

Define Magnetic quantum no.

Answer 66

Magnetic Quantum no. (m): the magnetic orbital quantum number **gives information about the special orientation of the orbital with respect to standard set of co-ordinate axis**. **For any subshell; 2l+1 values are possible**.

Question 67

Define Spin Quantum no.

Answer 67

Spin Quantum no. (s): Goudsmit proposed the presence of 4th quantum number known as the electron spin quantum number. **An electron has besides charge in mass intrinsic spin angular momentum number.** This quantum number can **have two orientations/values: +1/2 or -1/2** These are called the two spin states of the electron. **An orbital cannot hold more than two electrons and these two electrons should have opposite spins**.

Question 68

For 1s orbital the probability density is maximum at the nucleus and decreases sharply as we move away from it. True or False?

Answer 68

True

Question 70

Explain Nodal Surfaces Or Nodes.

Answer 70

For 2s orbital the probability density first decreases sharply to zero and again starts increasing. After reaching small maximum it decreases again and approaches zero as the value of r increases further. **The region where this probability density function reduces to zero is called the nodal surfaces or simply nodes**. In general, **ns- orbital has (n-1) nodes**. That is, **number of nodes increase with number of principal quantum no.** So, no of nodes for 2s orbital will be 1, for 3s will be 2 and so on.

Question 71

A: Size of the s orbital increases with increase in principal quantum no. R: Number of nodes increase with increase in principal quantum no.

Answer 71

(b) Both are true but not correct reason.

Question 72

The correct statement about probability density (except at infinite distance from nucleus ) is : A: it can be zero for 1s orbital B: it can be negative for 2p orbital C: it can be zero for 3p orbital D: it can never be zero for 2s orbital

Answer 72

(c) as no of nodes in 3p orbital is 1.

Question 73

A: The probability density always has some value, however small it may be, at any finite distance from the nucleus. R: At infinite distance, it may have zero value.

Answer 73

(b) Both true but not correct reason.

Question 74

Draw boundary surface diagrams for 1s and 2s orbital.

Answer 74

.

Question 75

What are angular nodes?

Answer 75

Angular nodes are also referred to as the nodal planes. **Angular node refers to a plane that passes through the nucleus and has probability density of zero**. **No. of angular nodes is given by l orbital.** Ex: One angular node for p orbitals, two for d orbitals and so on.... The total no. of nodes is given by (n-1). Total no. of angular nodes is given by l. **For hybrid orbitals**: such as 4sp³, the orbital with higher hybridization energy is selected to choose l for angular node. Here, it will be p orbital, hence the the no. of angular nodes will be l = 1.

Question 76

Explain energies of orbitals.

Answer 76

The energy of an electron in a hydrogen atom is determine solly by the principle quantum number. Thus, Energy of orbitals in hydrogen atom increase in order: **1s < 2s = 2p < 3s = 3p < 4s = 4p = 4d = 4f<** **Orbital$ having the same energy are called degenerate orbitals. The 1s orbital in hydrogen atom corresponds to stable condition and is called the ground state. Electron held in this orbit is strongly held by the nucleus. Electron in 2s, 2p or higher orbitals in hydrogen atom is in excited state.

Question 77

What are degenerate orbitals?

Answer 77

Orbitals having the same energy.

Question 78

Define Effective Nuclear Charge.

Answer 78

Due to the presence of electrons in the inner shells, the electrons in the outer shells will not experience the full positive charge of the nucleus. The effect will be lowered due to partial screening of positive charge on the nucleus by the inner shell electrons. **This is known as the shielding of the outer shell electrons from the nucleus by the inner shell electrons, and **the net positive charge experienced by the outer electrons is known as the effective nuclear charge**.

Question 79

How do you determine which orbital has lower energy?

Answer 79

**The lower the value (n+l) of an orbital, the lower its energy.** **If two orbitals have the same value of (n+l), the orbital with the lower value n will have the lower energy.

Question 80

Filling of orbitals takes place through on which basis?

Answer 80

1. Aufbau's Principle: it states that, **in the ground state of the atoms the orbitals are filled in order of their increasing energies.** In other words, electrons occupy the lowest energy orbitals available to them first and enter into higher energy orbitals only after the lower energy orbitals are filled. **There is no single ordering of energies of orbitals which will be universally correct for all atoms** Trick: ssp sps dps dps fdps fdps 2. Pauli's Exclusion Principle: according to this principal, **no two electrons in an atom can have the same set of 4 quantum numbers**. It can also be stated as, **only two electrons may exist in the same orbital and these electrons must have opposite spin.** This means that two electrons can have the same value of three quantum numbers and l, m and nmust have the opposite spin quantum number. **The maximum number of electrons in the shell with principal quantum number and is equal to 2n²** 3. Hund's Rule of Maximum Multiplicity: this rule deals with the filling of electrons into the orbitals belonging to the same subshell (that is, orbitals of same energy, **degenerate orbitals**). It states: **pairing of electrons in the orbitals belonging to the same shell does not take place until each orbital belonging to the subshell has got one electron each (singly occupied). **It has been observed that half filled and fully filled degenerate set of orbitals acquire extra stability due to their symmetry.**

Question 81

The maximum number of electrons in the shell with principal quantum number and is equal to?

Answer 81

2n²

Question 82

State Aufbau's Principle.

Answer 82

Aufbau's Principle: it states that, **in the ground state of the atoms the orbitals are filled in order of their increasing energies.** In other words, electrons occupy the lowest energy orbitals available to them first and enter into higher energy orbitals only after the lower energy orbitals are filled. **There is no single ordering of energies of orbitals which will be universally correct for all atoms** Trick: ssp sps dps dps fdps fdps

Question 83

State Hund's Rule.

Answer 83

Hund's Rule of Maximum Multiplicity: this rule deals with the filling of electrons into the orbitals belonging to the same subshell (that is, orbitals of same energy, **degenerate orbitals**). It states: **pairing of electrons in the orbitals belonging to the same shell does not take place until each orbital belonging to the subshell has got one electron each (singly occupied). **It has been observed that half filled and fully filled degenerate set of orbitals acquire extra stability due to their symmetry.**

Question 84

State Pauli's Exclusion Principle.

Answer 84

Pauli's Exclusion Principle: according to this principal, **no two electrons in an atom can have the same set of 4 quantum numbers**. It can also be stated as, **only two electrons may exist in the same orbital and these electrons must have opposite spin.** This means that two electrons can have the same value of three quantum numbers and l, m and nmust have the opposite spin quantum number. **The maximum number of electrons in the shell with principal quantum number and is equal to 2n²**

Question 85

Define electronic configuration.

Answer 85

The distribution of electrons into orbitals of an atom is called its electronic configuration. It can be done in two ways: 1. spd... notation 2. Orbital Diagram. A +ve spin is upward arrow and -ve spin is downward arrow. **The electrons in completely filled shells are known as core elements.** **Electrons that are added to the electronic shell with the highest principle quantum number are called Valance Electrons.**

Question 86

What are core electrons?

Answer 86

The electrons in completely filled shells are known as core elements.

Question 87

What are Valence Electrons?

Answer 87

Electrons that are added to the electronic shell with the highest principle quantum number are called Valance Electrons.