NCERT Class 11 Chemistry - Unit 2
Structure of Atom
Simple Notes in British English for Easy Understanding
Hello students! I am your Chemistry teacher, and today we are going to learn about the Structure of Atom. Think of an atom as the smallest building block of everything around us. In this chapter, we will see how scientists slowly discovered what is inside an atom. Let's make it very simple and interesting!
1. Important Facts - 15 Simple Points to Remember
- Dalton's Atomic Theory said that atoms are solid, indivisible, and all atoms of one element are the same. But later we found that atoms can be divided.
- J.J. Thomson discovered the electron in 1897 using a cathode ray experiment. Cathode rays are streams of negatively charged particles.
- Anode rays or canal rays are positively charged particles discovered by E. Goldstein. They helped in the discovery of the proton.
- The proton was discovered by Ernest Rutherford. It is a positively charged particle found in the nucleus.
- The neutron was discovered by James Chadwick in 1932. It has no charge and is also present in the nucleus.
- Millikan's Oil Drop Experiment found the charge on one electron. It is 1.602 × 10⁻¹⁹ coulombs.
- The e/m ratio for an electron is 1.758820 × 10¹¹ C/kg. This means an electron is very, very light.
- According to Thomson's model, also called the 'plum pudding model', an atom is a ball of positive charge with electrons stuck in it like raisins.
- Rutherford's α-particle scattering experiment showed that an atom is mostly empty space with a small, heavy, positive nucleus at the centre.
- Isotopes are atoms of the same element with same number of protons but different number of neutrons. Example: ¹H, ²H, ³H.
- Isobars are atoms of different elements that have the same mass number but different atomic number. Example: ⁴⁰Ar and ⁴⁰Ca.
- Max Planck gave the Quantum Theory. He said that energy is given out or taken in as small packets called 'quanta'. For light, this packet is a 'photon'.
- The Photoelectric Effect is when light hits a metal and electrons come out. This proved that light behaves like a particle.
- Bohr's Model says that electrons move in fixed paths called 'orbits' or 'shells'. They do not lose energy while in these orbits.
- The energy and frequency of light are related by E = hν and the speed of light is c = νλ. Here h is Planck's constant.
Properties of Fundamental Particles
| Particle | Symbol | Charge (Coulombs) | Mass (kg) | Discoverer |
|---|---|---|---|---|
| Electron | e⁻ | -1.602 × 10⁻¹⁹ | 9.109 × 10⁻³¹ | J.J. Thomson |
| Proton | p or H⁺ | +1.602 × 10⁻¹⁹ | 1.672 × 10⁻²⁷ | E. Rutherford |
| Neutron | n | 0 | 1.675 × 10⁻²⁷ | J. Chadwick |
2. Multiple Choice Questions (MCQs)
Choose the correct option. Answers are given at the end of this section.
1. Who discovered the electron?
2. Cathode rays are made up of:
3. The charge on an electron was measured by:
4. Which model is also known as the 'plum pudding model'?
5. The nucleus of an atom was discovered by:
6. Which particle has no charge?
7. Isotopes have the same number of:
8. ¹⁴C and ¹⁴N are examples of:
9. The e/m ratio is maximum for:
10. Who gave the Quantum Theory of radiation?
11. The energy of a photon is given by:
12. The photoelectric effect proves that light has:
13. The minimum energy needed to eject an electron from a metal surface is called:
14. According to Bohr, the angular momentum of an electron is:
15. Which series of hydrogen spectrum lies in the visible region?
16. The Rydberg formula is used to calculate:
17. A limitation of Rutherford's model was that it could not explain:
18. Black body radiation is an example of:
19. The value of Planck's constant is:
20. Anode rays are also called:
21. The number of protons in an atom is its:
22. In the electromagnetic spectrum, which has the highest frequency?
23. The relation c = νλ shows that:
24. Bohr's model is applicable only to:
25. Which of the following is not a fundamental particle?
26. The mass of a neutron is nearly equal to the mass of a:
27. Dalton's theory could not explain:
28. In Rutherford's experiment, most α-particles passed through the gold foil without deflection because:
29. The Lyman series of hydrogen spectrum lies in which region?
30. If the frequency of radiation increases, its energy will:
Answer Key for MCQs
1-C, 2-C, 3-B, 4-C, 5-C, 6-C, 7-B, 8-B, 9-C, 10-C, 11-B, 12-B, 13-D, 14-B, 15-B, 16-B, 17-B, 18-B, 19-A, 20-B, 21-B, 22-C, 23-B, 24-C, 25-C, 26-B, 27-C, 28-B, 29-C, 30-B
3. Fill in the Blanks
Complete the sentences. Answers are given after each blank.
- Dalton said that atoms are ________. Indivisible
- The electron was discovered by ________. J.J. Thomson
- Cathode rays are deflected towards the ________ plate in an electric field. Positive
- The charge on one electron is ________ coulombs. -1.602 × 10⁻¹⁹
- The neutron was discovered by ________. James Chadwick
- Atoms with same atomic number but different mass number are called ________. Isotopes
- ⁴⁰Ar and ⁴⁰Ca are examples of ________. Isobars
- Plum pudding model was given by ________. J.J. Thomson
- Rutherford's experiment used ________ particles. Alpha (α)
- The central, positive part of an atom is called the ________. Nucleus
- According to Planck, energy is emitted in the form of small packets called ________. Quanta
- The relation between energy and frequency is E = ________. hν
- The emission of electrons when light falls on a metal is the ________. Photoelectric effect
- The minimum frequency needed for photoelectric effect is called ________ frequency. Threshold
- A body which can absorb and emit all radiations is a ________. Black body
- The speed of light, c, is equal to ________. νλ
- The Balmer series of hydrogen lies in the ________ region. Visible
- Bohr's model explains the spectrum of ________ atom. Hydrogen
- The angular momentum of an electron in an orbit is equal to ________. nh/2Ï€
- A limitation of Rutherford's model is that it cannot explain the ________ of an atom. Stability
4. Frequently Asked Questions (FAQs)
Q1. What is an atom?
Ans: An atom is the smallest particle of an element that can take part in a chemical reaction. It has a central nucleus with protons and neutrons, and electrons moving around it.
Q2. Why is Dalton's theory not fully correct today?
Ans: Dalton said atoms cannot be divided. But now we know atoms have smaller parts like electrons, protons, and neutrons. Also, he did not know about isotopes.
Q3. What are cathode rays?
Ans: Cathode rays are streams of fast-moving electrons. They are produced in a discharge tube at low pressure and high voltage. They travel in straight lines and have negative charge.
Q4. What are anode rays?
Ans: Anode rays, or canal rays, are positively charged particles. They are formed when gas atoms lose electrons and become positive ions. Their e/m ratio depends on the gas used.
Q5. Why did most α-particles pass through the gold foil in Rutherford's experiment?
Ans: Because most of the space in an atom is empty. Only very few particles hit the small, heavy nucleus and got deflected.
Q6. What is the main difference between isotopes and isobars?
Ans: Isotopes have same number of protons but different neutrons. Isobars have same mass number but different protons. So, isotopes are of same element, isobars are of different elements.
Q7. What did Millikan's oil drop experiment prove?
Ans: It proved that the charge on an electron is fixed. The value is 1.602 × 10⁻¹⁹ C. It also showed that charge is 'quantised', meaning it comes in small, fixed units.
Q8. What is Planck's Quantum Theory in simple words?
Ans: It says that energy is not given out smoothly. It is given out in tiny packets called 'quanta'. The energy of one packet is E = hν, where h is a constant and ν is frequency.
Q9. What is the photoelectric effect?
Ans: When light of enough energy falls on a clean metal surface, electrons are thrown out. This only happens if the light's frequency is above a 'threshold' value. This proves light is a particle.
Q10. What is work function?
Ans: Work function is the minimum energy needed to remove one electron from a metal surface. It is different for different metals. Its symbol is Φ or W₀.
Q11. What is black body radiation?
Ans: A black body is an object that absorbs all light and emits all wavelengths when heated. The light it gives out has a continuous spectrum. Its colour changes with temperature.
Q12. What does c = νλ mean?
Ans: This formula shows the relation for all electromagnetic waves. 'c' is the speed of light, 'ν' (nu) is frequency, and 'λ' (lambda) is wavelength. Speed is constant, so if wavelength increases, frequency decreases.
Q13. What is the hydrogen spectrum?
Ans: When hydrogen gas is given energy, it gives out light. This light has specific, bright lines. These lines are grouped into series like Lyman, Balmer, Paschen, etc.
Q14. What is the Rydberg formula used for?
Ans: The Rydberg formula, 1/λ = R(1/n₁² - 1/n₂²), is used to find the wavelength or wavenumber of the lines in the hydrogen spectrum. R is the Rydberg constant.
Q15. What are the main postulates of Bohr's model?
Ans: 1. Electrons move in fixed circular orbits. 2. While in these orbits, they do not lose energy. 3. Energy is absorbed or emitted when an electron jumps from one orbit to another.
Q16. Why was Rutherford's model unstable?
Ans: Rutherford said electrons move around the nucleus. But physics says a moving charged particle must lose energy. So, the electron should fall into the nucleus, but it does not. Hence, the model was unstable.
Q17. What is the difference between a continuous and a line spectrum?
Ans: A continuous spectrum has all wavelengths without gaps, like a rainbow. A line spectrum has only a few bright lines of specific wavelengths, like the hydrogen spectrum.
Q18. Which has more energy, a red photon or a blue photon?
Ans: A blue photon has more energy. Blue light has a higher frequency and shorter wavelength than red light. From E = hν, higher frequency means more energy.
Q19. Can Bohr's model explain the spectra of multi-electron atoms?
Ans: No, Bohr's model works well only for hydrogen and ions with one electron like He⁺ or Li²⁺. It fails for atoms with two or more electrons because it does not consider electron-electron repulsion.
Q20. What is the charge and mass of a proton compared to an electron?
Ans: A proton has a positive charge, equal but opposite to the electron's negative charge. A proton is about 1836 times heavier than an electron.
5. Short Questions and Answers (2-3 Lines Each)
Q1. State one main point of Dalton's atomic theory.
Ans: Dalton said that matter is made of very small, indivisible particles called atoms. All atoms of one element are identical in mass and properties.
Q2. How are cathode rays produced?
Ans: Cathode rays are produced in a discharge tube. When high voltage is applied at very low gas pressure, a stream of electrons comes out from the cathode.
Q3. Who discovered the neutron and when?
Ans: The neutron was discovered by James Chadwick in the year 1932. He did this by bombarding beryllium with α-particles.
Q4. Define isotopes with one example.
Ans: Isotopes are atoms of the same element having same atomic number but different mass number. For example, Protium ¹H, Deuterium ²H, and Tritium ³H.
Q5. What is the e/m ratio for an electron?
Ans: The charge to mass ratio (e/m) for an electron is 1.758820 × 10¹¹ C/kg. This value was found by J.J. Thomson.
Q6. What was the main conclusion of Rutherford's experiment?
Ans: Rutherford concluded that an atom has a very small, dense, positively charged centre called the nucleus. Most of the atom is empty space.
Q7. What is a 'quantum' of light called?
Ans: According to Planck's theory, a quantum of light energy is called a 'photon'. Each photon has energy E = hν.
Q8. Write the formula for the photoelectric effect.
Ans: Einstein's equation for the photoelectric effect is hν = Φ + K.E. Here, hν is the energy of the photon, Φ is the work function, and K.E. is kinetic energy of the electron.
Q9. What is the visible region of the electromagnetic spectrum?
Ans: The visible region is the small part of the spectrum that our eyes can see. Its wavelength range is about 400 nm (violet) to 750 nm (red).
Q10. Name the series of hydrogen spectrum in the UV region.
Ans: The Lyman series of the hydrogen spectrum lies in the ultraviolet (UV) region. It is formed when electrons fall to the first orbit (n=1).
Q11. What is the value of the Rydberg constant?
Ans: The value of the Rydberg constant (R) for hydrogen is 109,677 cm⁻¹ or 1.09677 × 10⁷ m⁻¹. It is used in the Rydberg formula.
Q12. What is meant by 'stationary state' in Bohr's model?
Ans: A stationary state is a fixed orbit where an electron can move without losing energy. Each orbit has a definite energy, so it is also called an energy level.
Q13. Why does a black body appear black?
Ans: An ideal black body absorbs all the light that falls on it. It does not reflect any light. That is why it appears black in colour.
Q14. What is the relationship between wavelength and energy?
Ans: Wavelength and energy are inversely proportional. This means if wavelength is short, energy is high. The formula is E = hc/λ.
Q15. Mention one limitation of Thomson's model.
Ans: Thomson's model could not explain the results of Rutherford's α-particle scattering experiment. It did not have the concept of a nucleus.
Q16. What are isobars? Give an example.
Ans: Isobars are atoms of different elements with the same mass number but different atomic number. Example: ⁴⁰₁₈Ar and ⁴⁰₂₀Ca.
Q17. What happens when an electron jumps from a higher to a lower orbit?
Ans: When an electron jumps from a higher to a lower orbit, it emits energy. This energy is given out as a photon of light with energy ΔE = E₂ - E₁ = hν.
Q18. What is the nature of anode rays?
Ans: Anode rays are made of positive ions of the gas in the discharge tube. So, their charge and mass depend on which gas is used in the tube.
Q19. Why is Bohr's model called a 'quantised' model?
Ans: It is called quantised because the electron can only move in orbits with certain fixed energy levels. The angular momentum is also fixed as mvr = nh/2Ï€.
Q20. What is the main failure of Rutherford's model?
Ans: The main failure was that it could not explain why atoms are stable. A revolving electron should lose energy and crash into the nucleus, but this does not happen.
6. Long Questions and Answers (Detailed but Simple)
Q1. Explain Dalton's Atomic Theory and its limitations.
Ans: My dear students, John Dalton gave his atomic theory in 1808. It had these main points:
1. All matter is made of very tiny particles called atoms.
2. Atoms are indivisible. They cannot be broken into smaller parts.
3. All atoms of one element are identical. They have the same mass and properties.
4. Atoms of different elements have different masses and properties.
5. Atoms combine in small, whole numbers to form compounds.
Limitations: This theory was very good for its time, but later science found some problems. First, atoms are divisible. We now know they have electrons, protons, and neutrons. Second, atoms of the same element are not always identical. We have isotopes like ¹²C and ¹⁴C which have different masses. Third, it could not explain why atoms combine and how they combine.
Q2. Describe J.J. Thomson's Cathode Ray experiment. What did he conclude?
Ans: J.J. Thomson did a famous experiment using a discharge tube. A discharge tube is a glass tube with gas at very low pressure. He applied a high voltage between two metal plates called the cathode (negative) and anode (positive).
He saw a stream of rays coming from the cathode. These rays were called cathode rays. He did three tests:
1. He put a small object in the path, and it made a shadow. This showed the rays travel in straight lines.
2. He put a small fan, and it started moving. This showed the rays have mass and energy.
3. He put an electric and magnetic field. The rays bent towards the positive plate. This showed the rays are negatively charged.
Conclusion: Thomson concluded that these rays are made of very small, negatively charged particles. He called them 'corpuscles', but now we call them electrons. All atoms contain electrons. He also found their e/m ratio, which was the same no matter which gas was used.
Q3. Explain Rutherford's α-particle scattering experiment and its results.
Ans: Ernest Rutherford and his students Geiger and Marsden did a very important experiment. They took a very thin sheet of gold foil. They hit this foil with fast-moving, positively charged α-particles from a radioactive source. They had a zinc sulphide screen around the foil to see where the particles went.
Observations:
1. Most of the α-particles passed straight through the foil without any deflection.
2. Some particles were deflected by small angles.
3. A very, very small number of particles (about 1 in 20,000) bounced back at angles greater than 90°.
Conclusions: From this, Rutherford made his model of the atom.
1. Since most particles went straight, he said most of the atom is empty space.
2. Since some were deflected, there must be a heavy, positively charged part in the atom. He called it the nucleus.
3. Since a few bounced back, the nucleus must be very small and dense. The electrons move around this nucleus in empty space, just like planets around the sun.
Q4. What is Planck's Quantum Theory? Explain black body radiation.
Ans: Before Max Planck, scientists thought energy is continuous, like water flowing from a tap. But this could not explain black body radiation. A black body is a perfect absorber and emitter of all radiation. When you heat an object, its colour changes from red to white to blue. The old theory failed here.
In 1900, Max Planck gave a new theory. He said:
1. Energy is not continuous. It is emitted or absorbed in small, separate packets.
2. Each packet of energy is called a 'quantum'. For light, it is called a 'photon'.
3. The energy of one quantum is directly proportional to its frequency. The formula is E = hν. Here, 'h' is Planck's constant (6.626 × 10⁻³⁴ Js) and 'ν' is frequency.
4. An object can emit energy only in whole number multiples, like 1hν, 2hν, 3hν, but not 1.5hν.
This theory perfectly explained black body radiation and was the start of quantum physics.
Q5. Explain the Photoelectric Effect. What did Einstein explain about it?
Ans: The photoelectric effect is a very interesting phenomenon. When light of a certain frequency shines on a clean metal surface, like potassium or cesium, electrons are emitted from the surface. These electrons are called photoelectrons.
Key observations that were confusing:
1. Electrons are emitted only if the light frequency is above a minimum value, called threshold frequency (ν₀). Below this, no electrons come out, no matter how bright the light is.
2. Once the frequency is above the threshold, increasing the brightness of light increases the number of electrons, but not their energy.
3. The kinetic energy of the electrons depends on the frequency of light, not its brightness.
Einstein's Explanation: In 1905, Albert Einstein used Planck's quantum theory to explain this. He said light is made of particles called photons.
1. One photon hits one electron. If the photon's energy (hν) is more than the energy holding the electron (work function, Φ), the electron is ejected.
2. The rest of the energy becomes the electron's kinetic energy.
3. His famous equation is: hν = Φ + ½mv². Here, Φ = hν₀.
This proved that light has a particle nature, and Einstein got the Nobel Prize for this.
Q6. State the postulates of Bohr's Model of Atom. What are its limitations?
Ans: Neils Bohr improved Rutherford's model in 1913. He used Planck's quantum theory to explain the stability of the atom and the hydrogen spectrum. His main postulates are:
Postulates:
1. Stationary Orbits: Electrons move around the nucleus in certain fixed circular paths called orbits or shells. Each orbit has a definite energy. While in these orbits, electrons do not radiate energy. These are called stationary states.
2. Quantisation of Angular Momentum: An electron can only move in those orbits where its angular momentum is a whole number multiple of h/2Ï€. The formula is mvr = nh/2Ï€, where n = 1, 2, 3... This 'n' is called the principal quantum number.
3. Energy Changes: Energy is absorbed when an electron jumps from a lower orbit to a higher orbit. Energy is emitted when it jumps from a higher to a lower orbit. The energy change is ΔE = E₂ - E₁ = hν.
Limitations: Bohr's model was very successful, but it had limitations.
1. It could not explain the spectra of atoms with more than one electron, like helium.
2. It could not explain the fine structure of the hydrogen spectrum. On high resolution, each line is actually many lines close together.
3. It could not explain the Zeeman effect (splitting of lines in a magnetic field) and Stark effect (splitting in an electric field).
4. It did not say anything about the shapes of molecules.
Q7. Explain the Hydrogen Spectrum and Rydberg Formula.
Ans: When you pass an electric discharge through hydrogen gas at low pressure, it gives out light. If you pass this light through a prism, you do not get a continuous rainbow. You get a series of sharp, bright lines. This is called a line spectrum or atomic spectrum. For hydrogen, these lines fall into different series.
Main Series of Hydrogen Spectrum:
1. Lyman Series: When electrons fall to the 1st orbit (n=1). This is in the Ultraviolet (UV) region.
2. Balmer Series: When electrons fall to the 2nd orbit (n=2). This is the only series in the Visible region. This is why we can see it.
3. Paschen, Brackett, Pfund Series: These are for jumps to n=3, n=4, n=5. They are in the Infrared (IR) region.
Rydberg Formula: In 1888, Johannes Rydberg gave a simple formula to calculate the wavelength of these lines. For hydrogen, the formula is:
1/λ = R(1/n₁² - 1/n₂²)
Here, λ is the wavelength, R is the Rydberg constant (109,677 cm⁻¹), n₁ is the lower orbit, and n₂ is the higher orbit. n₂ is always greater than n₁. This formula matched the experimental values perfectly and was a great success for Bohr's model.
Q8. Compare Thomson's Model and Rutherford's Model of Atom.
Ans: Let us compare these two early models of the atom in a simple table to understand the differences.
| Feature | Thomson's Model (Plum Pudding Model) | Rutherford's Model (Nuclear Model) |
|---|---|---|
| Year | 1898 | 1911 |
| Structure | Atom is a sphere of positive charge. Negative electrons are embedded in it like plums in a pudding. | Atom has a tiny, dense, positive nucleus at the centre. Electrons revolve around it in empty space. |
| Positive Charge | The positive charge is spread out evenly in the whole atom. | The positive charge is concentrated in the small nucleus. |
| Mass | Mass of the atom is uniformly distributed. | Almost all the mass of the atom is in the nucleus. |
| Main Evidence | Based on the discovery of the electron. It explained electrical neutrality. | Based on the α-particle scattering experiment. It showed a nucleus exists. |
| Main Failure | Could not explain the scattering of α-particles. Could not explain a nucleus. | Could not explain the stability of the atom. A revolving electron should lose energy and fall. |
So, Rutherford's model was a big improvement over Thomson's model because it introduced the nucleus. But it was still not perfect.
Q9. What are the limitations of Rutherford's Model? How did Bohr's model solve them?
Ans: Rutherford's nuclear model was very important, but it had two major problems according to classical physics.
Limitations of Rutherford's Model:
1. Stability of Atom: Rutherford said electrons revolve around the nucleus. But physics says that any charged particle moving in a circle must continuously lose energy in the form of electromagnetic radiation. If this happened, the electron would spiral into the nucleus in less than 10⁻⁸ seconds. But we know atoms are very stable. So, the model could not explain stability.
2. Atomic Spectrum: If an electron lost energy continuously, it should give a continuous spectrum. But experiments showed that atoms give a line spectrum with specific wavelengths. Rutherford's model could not explain why.
How Bohr Solved Them: Neils Bohr solved these by applying the idea of 'quantisation'.
1. He said electrons do not radiate energy when they are in specific 'stationary orbits'. This directly solved the stability problem.
2. He said energy is only emitted or absorbed when an electron jumps between these fixed orbits. The energy difference ΔE = hν gives a specific frequency. This explained why we see sharp lines in the spectrum, not a continuous band. So, Bohr's model successfully saved the nuclear model.
Q10. Explain the terms: Electromagnetic Spectrum, Wavelength, Frequency, and Wave Number.
Ans: These are very important terms to understand light and other radiations. Remember, light behaves as a wave.
1. Electromagnetic Spectrum: All types of electromagnetic radiation are arranged in order of their wavelength or frequency. This complete arrangement is the electromagnetic spectrum. It starts from very long radio waves, to microwaves, infrared, visible light, ultraviolet, X-rays, and finally very short gamma rays. Visible light is only a tiny part of this.
2. Wavelength (λ, Lambda): It is the distance between two nearest crests or two nearest troughs of a wave. Think of it as the length of one full wave. It is measured in metres (m), nanometres (nm), or Angstroms (Ã…). 1 nm = 10⁻⁹ m.
3. Frequency (ν, Nu): It is the number of waves that pass a given point in one second. Its unit is Hertz (Hz) or per second (s⁻¹). If frequency is high, many waves pass in one second.
4. Wave Number (ν̄, Nu bar): It is the number of waves or wavelengths per unit length, usually per centimetre. It is the reciprocal of wavelength. The formula is ν̄ = 1/λ. Its unit is cm⁻¹ or m⁻¹.
Relation: All three are related to the speed of light (c = 3.0 × 10⁸ m/s). The formula is c = νλ. This shows wavelength and frequency are inversely proportional. Also, energy E = hν = hc/λ.
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