In the first-year chemistry curriculum, Chapter 5 on Atomic Structure is a foundational and pivotal topic. This chapter delves into the fundamental building blocks of matter, the atoms, and their intricate internal structure. Students explore the historical development of atomic models, beginning with Dalton’s simple sphere and progressing through Thomson’s plum pudding model, Rutherford’s nuclear model, and ultimately Bohr’s quantum model. They learn about the discovery and properties of subatomic particles, including electrons, protons, and neutrons, as well as their organization within the atom.
The concept of energy levels and quantization of energy plays a crucial role in understanding electron configurations and the periodic table. Additionally, students are introduced to the concept of atomic number, isotopes, and atomic mass, all of which are vital in understanding the diversity and behavior of elements in chemical reactions. Chapter 5 serves as a foundational cornerstone in the study of chemistry, providing the necessary groundwork for comprehending chemical bonding, reactions, and the properties of matter at the atomic level.
Short Question Chemistry 1st Year Chapter 5 Atomic Structure
What are atoms made up of, according to Dalton’s theory?
Atoms were considered ultimate particles that could not be divided any further.
How has our understanding of the structure of atoms changed over the years?
Our understanding of the structure of atoms has undergone radical changes, with the discovery of subatomic particles.
Describe the experiment that led to the discovery of electrons (cathode rays).
In the experiment, a gas discharge tube is fitted with metallic electrodes (cathode and anode) filled with gas. When a high voltage is applied at low pressure, cathode rays are observed.
What happens when the pressure inside the gas discharge tube is reduced to about 0.01 torr?
When the pressure is reduced to about 0.01 torr, the original glow disappears, and cathode rays are produced, creating fluorescence on the glass wall opposite to the cathode.
What is the charge of cathode rays, and how was it established by J. Perrin and J. Thomson?
Cathode rays are negatively charged. J. Perrin demonstrated this by observing their curved path in a magnetic field, and J. Thomson confirmed it by deflecting them upwards with an electric field.
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What happens when cathode rays strike the walls of a glass tube, and what materials can they induce fluorescence in?
Cathode rays produce a greenish fluorescence when they strike the glass tube’s walls. They can also induce fluorescence in rare earths and minerals. Alumina glows red, and tin stone glows yellow in their presence.
How do cathode rays demonstrate their ability to travel in straight lines, and in what direction?
Cathode rays cast a shadow when an opaque object is placed in their path, proving that they travel in straight lines perpendicular to the cathode’s surface.
What evidence suggests that cathode rays are not just rays but material particles?
Cathode rays can drive a small paddle wheel placed in their path, indicating they possess momentum. This observation implies that cathode rays are material particles with definite mass and velocity.
How can cathode rays produce X-rays, and under what conditions is this more likely to occur?
Cathode rays can produce X-rays when they strike an anode, particularly when the anode has a large atomic mass.
What happens when cathode rays fall on matter, and can you provide an example?
When cathode rays fall on matter, they can produce heat. For instance, when cathode rays from a concave cathode are focused on a platinum foil, the foil begins to glow.
What are the ionizing capabilities of cathode rays?
Cathode rays can ionize gases.
What chemical effect can cathode rays have?
Cathode rays can cause a chemical change, as they have a reducing effect.
What materials can cathode rays pass through?
Cathode rays can pass through a thin metal foil, such as aluminum or gold foil.
What did J.J. Thomson conclude about the nature of cathode rays, and what did he name the particles within them?
J.J. Thomson concluded that cathode rays consist of streams of negatively charged particles. He named these particles electrons. Thomson also determined the charge-to-mass ratio (e/m) of electrons and found it to be consistent across different gases, leading to the conclusion that all atoms contain electrons.
Who is credited with the discovery of positive rays?
German physicist E. Goldstein discovered positive rays in 1886.
How are positive rays produced in a discharge tube?
Positive rays are produced when high-speed cathode rays (electrons) strike gas molecules in the discharge tube, knocking out electrons and creating positive ions.
Why are positive rays called “positive”?
Positive rays are named so because they carry a positive charge.
What are the properties of positive rays?
Positive rays are deflected by both electric and magnetic fields, travel in the opposite direction of cathode rays, produce flashes on ZnS plates, and have an e/m value that depends on the gas used in the discharge tube.
What particle was discovered to be the lightest among all positive particles?
The positive particle obtained from hydrogen gas, known as a proton, was found to be the lightest among all positive particles. It was suggested by Rutherford and is 1836 times more massive than an electron.
Who discovered the neutron and when?
James Chadwick discovered the neutron in 1932 and was awarded the Nobel Prize in Physics in 1935.
How was the neutron discovered experimentally?
Neutrons were discovered when a stream of alpha particles was directed at a beryllium (9Be) target, resulting in the production of neutrons in a nuclear reaction.
What are some properties of neutrons?
Neutrons decay into protons, electrons, and neutrinos; they cannot ionize gases; they are highly penetrating particles; and they can induce nuclear reactions when used as projectiles.
How are slow neutrons different from fast neutrons?
Slow neutrons have less energy (below 1 eV) and are more effective for nuclear fission purposes, whereas fast neutrons have higher energy (1.2 MeV or more).
What happens when slow neutrons hit copper (Cu) metal?
Slow neutrons hitting copper metal can emit gamma (γ) radiation, converting radioactive 66Cu into 66Zn through a nuclear reaction.
How are neutrons used in the treatment of cancer?
Neutrons, due to their intense biological effects, are used in cancer treatment, specifically for radiation therapy.
What is the purpose of J.J. Thomson’s apparatus for measuring the e/m value of electrons?
The apparatus is used to measure the e/m (charge-to-mass) value of electrons by passing cathode rays through electric and magnetic fields and observing their behavior.
How did Millikan determine the charge on an electron?
Millikan determined the charge on an electron using his oil drop experiment, where he observed the motion of charged oil droplets in an electric field and calculated the charge based on their behavior.
What is the smallest charge that Millikan found on any droplet in his experiment?
Millikan found the smallest charge on any droplet to be approximately 1.59 x 10^-19 coulombs, which is very close to the current accepted value of the charge of one electron.
What is the mass of an electron, and how was it determined?
The mass of an electron is approximately 9.1095 x 10^-31 kg. It was determined by using the known charge of an electron (1.602 x 10^-19 coulombs) and the measured e/m value (1.7588 x 10^11 coulombs kg^-1).
What did Lord Rutherford’s experiment with alpha particles and gold foil reveal about atomic structure?
Rutherford’s experiment showed that most alpha particles passed through the gold foil with little deflection, but some were deflected at large angles, and a few even bounced back. This led to the conclusion that atoms have a small, dense nucleus at the center and that most of the atom is empty space.
Who proposed the quantum theory in 1900?
Max Planck proposed the quantum theory in 1900.
What are the main points of Planck’s quantum theory?
The main points of Planck’s quantum theory are:
(i) Energy is emitted or absorbed in a discontinuous manner and in the form of wave packets called quanta.
(ii) The energy of a quantum is proportional to the frequency of the radiation.
(iii) Energy can only be emitted or absorbed in terms of quanta.
What is Planck’s constant, and what is its value?
Planck’s constant (h) is a constant that relates the energy and frequency of a photon. Its value is 6.626×10^-34 Js.
How is the energy of a photon related to its frequency and wavelength?
The energy of a photon (E) is related to its frequency (v) and wavelength (λ) by the equation E = hv = hc/λ, where c is the speed of light.
What did Bohr’s model of the atom propose?
Bohr’s model of the atom proposed that electrons can only exist in certain quantized energy levels, and they revolve in circular orbits around the nucleus. Energy is emitted or absorbed by electrons when they jump from one orbit to another.
What are the two components that make up the total energy of an electron in an orbit?
The total energy of an electron in an orbit is composed of kinetic energy (equal to 1/2mv^2) and potential energy.
How is the potential energy of an electron at a distance ‘r’ from the nucleus calculated?
The potential energy of an electron at a distance ‘r’ from the nucleus is calculated using the expression – (Ze^2 / 4πε₀r), where ‘Z’ is the atomic number, ‘e’ is the charge of an electron, ‘ε₀’ is the permittivity of free space, and ‘r’ is the distance.
What does the negative sign in the potential energy equation indicate?
The negative sign in the potential energy equation indicates that the potential energy of an electron decreases as it is brought closer to the nucleus. At infinity, the potential energy is zero, and it becomes increasingly negative as the electron approaches the nucleus.
How is the total energy of an electron in an orbit related to its kinetic and potential energies?
The total energy (E) of the electron is the sum of its kinetic energy and potential energy, given by E = (1/2)mv^2 – (Ze^2 / 4πε₀r).
What is the ionization energy of hydrogen according to Bohr’s model?
The ionization energy of hydrogen according to Bohr’s model is 1313.31 kJ/mol, which is the energy required to remove an electron from the hydrogen atom’s lowest energy level to infinity.
What happens to radiation when it passes through a prism?
Radiation undergoes refraction or bending.
How does the extent of bending of radiation in a prism depend on wavelength?
Longer wavelength radiation is bent to a smaller degree than shorter wavelength radiation.
What happens to white light when it passes through a prism?
White light is split up into radiations of different wavelengths.
What are the colors of the visible spectrum, and what is their wavelength range?
The colors of the visible spectrum are violet, indigo, blue, green, orange, yellow, and red, with wavelengths ranging from 400 nm to 750 nm.
Besides the visible spectrum, what are the other regions of the spectrum mentioned?
The other regions of the spectrum mentioned include ultraviolet, X-rays, gamma rays, cosmic rays (towards the lower wavelength end), as well as infrared, microwave, and radio frequency regions.
What is the term used to describe the visual display or dispersion of the components of white light when it passes through a prism?
It is called a spectrum.
How many types of spectra are there, and what are they?
There are two types of spectra: continuous spectrum and line spectrum.
Provide an example of a continuous spectrum.
The best example of a continuous spectrum is a rainbow, which is obtained from the light emitted by the sun or incandescent solids.
What is the characteristic of a continuous spectrum?
In a continuous spectrum, the boundary line between colors cannot be marked, and the colors diffuse into each other without any dark spaces.
What is atomic spectrum, and what does it depend on?
Atomic spectrum is the spectrum observed when an element or its compound is volatilized, and it depends on the element volatilized. The number of lines and their distances between them vary for different elements.
How can atomic spectrum be observed?
Atomic spectrum can be observed by volatilizing elements in their gaseous state at high temperatures or subjecting them to an electric discharge.
What are the two ways in which an atomic spectrum can be viewed?
An atomic spectrum can be viewed as atomic emission spectrum and atomic absorption spectrum.
Describe atomic emission spectrum.
Atomic emission spectrum is obtained when solids are volatilized or elements in their gaseous states are heated to high temperatures or subjected to an electrical discharge. It contains bright lines against a dark background.
What is an atomic absorption spectrum?
An atomic absorption spectrum is a spectrum of radiation produced when a beam of white light passes through a gaseous sample of an element, and certain wavelengths are absorbed by the element, resulting in dark lines on a bright background.
How are the lines in an atomic absorption spectrum different from those in an atomic emission spectrum?
In an atomic absorption spectrum, the lines appear dark because the wavelengths are absorbed by the element, whereas in an atomic emission spectrum, the lines appear bright because the corresponding wavelengths are emitted by the element.
What is the significance of the spectral lines in the hydrogen spectrum?
The spectral lines in the hydrogen spectrum represent the wavelengths of light emitted or absorbed by hydrogen atoms when electrons transition between energy levels. They are important for understanding the atomic structure of hydrogen.
How are the spectral lines in the hydrogen spectrum classified?
The spectral lines in the hydrogen spectrum are classified into five groups or series: Lyman series (UV region), Balmer series (visible region), Paschen series (IR region), Brackett series (IR region), and Pfund series (IR region).
How does Bohr’s model explain the origin of spectral lines in the hydrogen spectrum?
According to Bohr’s model, spectral lines in the hydrogen spectrum are produced when electrons transition between different energy levels within the atom. The specific wavelengths of light are emitted or absorbed during these transitions, leading to the observed spectral lines.
What is the equation for calculating wave numbers (v) in the text?
The equation for calculating wave numbers is v = a(Z-b), where ‘a’ and ‘b’ are constants characteristic of the metal under consideration.
What is the significance of Moseley’s Law?
Moseley’s Law establishes a relationship between the frequency of X-ray spectral lines and the atomic number (Z) of the element emitting them. It shows that atomic number, not atomic mass, determines an element’s characteristic properties.
How did Moseley’s Law impact the periodic table?
Moseley used his law to arrange elements like K and Ar, Ni and Co in a proper way in Mendeleev’s periodic table.
Can you explain the concept of Rydberg constant?
The Rydberg constant is mentioned as the value of the factor in equation (25), calculated to be 1.09678 x 10^7 m^-1. It’s a fundamental constant used in atomic physics to describe the wavelengths of spectral lines in hydrogen and hydrogen-like atoms.
What is the difference between Zeeman effect and Stark effect, and why can’t Bohr’s theory explain them?
The Zeeman effect involves the splitting of spectral lines in a magnetic field, while the Stark effect involves splitting in an electric field. Bohr’s theory cannot explain these effects as it was developed for hydrogen’s simple one-electron system and does not consider the complex interactions involved in multi-electron systems.
What are the K-series and L-series in X-ray spectral lines?
The K-series and L-series in X-ray spectral lines are distinct groups of lines with different wavelengths. K-series consists of shorter wavelengths, while L-series consists of relatively longer wavelengths.
How did Moseley’s research help in discovering new elements?
Moseley’s Law led to the discovery of new elements by allowing scientists to determine atomic numbers accurately, which helped in identifying missing elements in the periodic table. Examples mentioned in the text include the discovery of elements like Tc(43), Pr(59), and Rh(45).
What is the relationship between the frequency of X-ray spectral lines and atomic number according to Moseley’s Law?
According to Moseley’s Law, the frequency (v) of a particular line in an X-ray spectrum varies as the square of the atomic number (Z) of the element emitting it, expressed as v = a(Z-b).
What is wave-particle duality in matter?
Wave-particle duality in matter refers to the concept that matter particles, such as electrons, protons, neutrons, atoms, and molecules, exhibit both wave-like and particle-like characteristics simultaneously.
Who developed the mathematical equation that relates the wavelength of an electron to its momentum?
Louis de-Broglie developed the mathematical equation that relates the wavelength (λ) of an electron to its momentum (mv), expressed as λ = h / (mv).
What is the significance of the de-Broglie wavelength of an electron in the first orbit of a hydrogen atom?
The de-Broglie wavelength of an electron in the first orbit of a hydrogen atom is approximately 0.33 nanometers (nm). This value is comparable to the wavelength of X-rays and can be measured, demonstrating the wave-like nature of electrons.
What happens to the wavelength of a proton and an alpha particle when they move with the same velocity as an electron?
When a proton or an alpha particle moves with the same velocity as an electron, their wavelengths become significantly smaller compared to that of an electron. The wavelength of a proton is 1836 times smaller than that of an electron, and an alpha particle’s wavelength is 7344 times smaller.
What experiment was conducted by Davisson and Germer in 1927, and what did it verify?
Davisson and Germer conducted an experiment in 1927 to verify the wave nature of moving electrons. They showed that accelerated electrons undergo diffraction, similar to waves, when they strike a nickel crystal, thus providing experimental evidence for the wave nature of electrons.
What is Heisenberg’s Uncertainty Principle, and when is it applicable?
Heisenberg’s Uncertainty Principle states that it is impossible to simultaneously measure both the exact position (Δx) and momentum (Δp) of a microscopic particle, like an electron, with great accuracy. This principle is applicable only to microscopic particles and has no significance for macroscopic objects. The principle is mathematically expressed as Δx Δp ≥ h / 4π, where h is Planck’s constant.
Why does Bohr’s picture of an atom face criticism, and how does it relate to the uncertainty principle?
Bohr’s atom is criticized because it postulates specific electron orbits with defined velocities, which contradicts the uncertainty principle that prevents the simultaneous measurement of both position and velocity of electrons.
How did Schrodinger, Heisenberg, and Dirac address the limitations of Bohr’s atom model?
They developed wave theories of the atom, with Schrodinger’s wave equation being the most well-known, to describe electron behavior in terms of probability distributions rather than fixed orbits.
What is the maximum probability radius for finding an electron in a hydrogen atom’s ground state, and how does it compare to Bohr’s first orbit?
The maximum probability radius is 0.053 nm, which matches the radius calculated for Bohr’s first orbit.
What is an atomic orbital, and how does it differ from an orbit in Bohr’s theory?
An atomic orbital is a region in space where there is a 95% chance of finding an electron. It differs from Bohr’s orbits, as it represents a probability distribution rather than a fixed path.
What are quantum numbers, and how many are needed to describe an electron completely?
Quantum numbers are numerical values that describe an electron’s properties in an atom. Four quantum numbers are required to describe an electron completely.
What is the principal quantum number, and what does it represent in an atom?
The principal quantum number (n) represents the energy level or shell in which an electron revolves around the nucleus in an atom. It can take non-zero positive integer values.
How does the azimuthal quantum number relate to subshells, and what values can it have?
The azimuthal quantum number (λ) represents subshells within an energy level and can have values ranging from 0 to (n-1), where ‘n’ is the principal quantum number.
How are azimuthal quantum numbers related to subshell shapes?
The value of λ is related to the shape of the subshell. For example, λ = 0 corresponds to a spherical subshell, λ = 1 to a dumb-bell shape (p-subshell), and λ = 2 to a more complicated shape (d-subshell).
How can you calculate the total number of electrons in a subshell using the azimuthal quantum number?
The total number of electrons in a subshell can be calculated using the formula 2(2λ + 1).
What is the total number of electrons in the s-subshell, p-subshell, d-subshell, and f-subshell when λ = 0, 1, 2, and 3, respectively?
When λ = 0 (s-subshell): 2 electrons, λ = 1 (p-subshell): 6 electrons, λ = 2 (d-subshell): 10 electrons, λ = 3 (f-subshell): 14 electrons.
What are the possible values of the magnetic quantum number (m)?
The values of the magnetic quantum number (m) are m = 0, ±1, ±2, ±3, and so on.
How does the value of ‘m’ depend on the value of ‘λ’?
The value of ‘m’ depends on the value of ‘λ’, and for a given ‘λ’, there are (2λ + 1) possible values of ‘m’.
What does the magnetic quantum number (m) indicate?
The magnetic quantum number (m) indicates the degeneracy of orbitals in space, representing the number of different ways a subshell can be arranged in the presence of a magnetic field.
How many different space orientations does an s-subshell (λ = 0) have?
An s-subshell (λ = 0) has only one space orientation, and it can be arranged in space in only one way along the x, y, and z-axes.
How many space orientations does a p-subshell (λ = 1) have?
A p-subshell (λ = 1) has three space orientations, represented by m = 0, ±1, and these orientations correspond to the px, py, and pz orbitals.
What is the degeneracy of d-subshells (λ = 2) in the absence of a magnetic field?
In the absence of a magnetic field, d-subshells (λ = 2) have five-fold degenerate orbitals, represented by dxy, dyz, dzx, dx²-y², and dz².
How many different space orientations are there for an f-subshell (λ = 3)?
An f-subshell (λ = 3) has seven different space orientations, represented by m = 0, ±1, ±2, ±3, and these correspond to the complex shapes of f-orbitals.
What is the significance of the magnetic quantum number in atomic structure?
The magnetic quantum number determines the orientation of orbitals in space and is often referred to as the orbital orientation quantum number. It provides information about how subshells are arranged in the presence of a magnetic field.
How many electrons can an orbital accommodate?
An orbital can have at most two electrons.
What is the maximum number of electrons that can be accommodated in a shell, and what is the formula to calculate it?
The maximum number of electrons in a shell is given by the formula 2n^2, where n is the principal quantum number, and it cannot have a value of zero.
What are the three rules adopted to distribute electrons in subshells or orbitals?
The three rules are Aufbau principle, Pauli’s exclusion principle, and Hund’s rule.
According to the (n + λ) rule, how are subshells arranged in ascending order of energy?
Subshells are arranged in the increasing order of (n + λ) values, and if any two subshells have the same (n + λ) values, then the subshell with the smaller n value is placed first.
What is the order of arrangement for subshells in ascending order of energy?
The arrangement of subshells in ascending order of energy is as follows: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, and so on.
What is the Aufbau principle?
The Aufbau principle states that electrons should be filled in energy subshells in order of increasing energy values. Electrons are first placed in 1s, then 2s, 2p, and so on.
What is Pauli’s exclusion principle?
Pauli’s exclusion principle states that it is impossible for two electrons residing in the same orbital of a poly-electron atom to have the same values of four quantum numbers. Two electrons in the same orbital must have opposite spins (↑↓).
What is Hund’s rule?
Hund’s rule states that if degenerate orbitals are available and more than one electron is to be placed in them, they should be placed in separate orbitals with the same spin rather than putting them in the same orbital with opposite spins.
Why is it necessary to decrease the pressure in the discharge tube to get the cathode rays?
It is necessary to decrease the pressure in the discharge tube to get the cathode rays because at lower pressures, the mean free path of gas molecules becomes larger. In other words, the gas molecules are more spaced out in a low-pressure environment. This allows the high-speed electrons emitted from the cathode to travel longer distances without frequent collisions with gas molecules. As a result, they can form a distinct beam, which we observe as cathode rays.
Whichever gas is used in the discharge tube, the nature of the cathode rays remains the same. Why?
The nature of cathode rays remains the same regardless of the gas used in the discharge tube because cathode rays are actually streams of electrons. Electrons have a negative charge and a very small mass compared to gas molecules. Therefore, the behavior of cathode rays is primarily determined by the properties of electrons rather than the nature of the gas in the tube.
Why e/m value of the cathode rays is just equal to that of electron?
The e/m (charge-to-mass) value of cathode rays is equal to that of electrons because cathode rays are composed of electrons. The e/m value is a fundamental property of charged particles, and since cathode rays are essentially a stream of electrons, their e/m value matches that of electrons.
How the bending of the cathode rays in the electric and magnetic fields shows that they are negatively charged?
The bending of cathode rays in electric and magnetic fields shows that they are negatively charged because negatively charged particles (such as electrons) experience deflection in the presence of electric and magnetic fields. The direction of the deflection and the amount of bending can be used to determine the charge-to-mass ratio (e/m) of the particles, confirming their negative charge.
Why the positive rays are also called canal rays?
Positive rays are also called canal rays because they pass through a canal or slit in the anode of the discharge tube. These rays are composed of positively charged ions formed when high-speed electrons from the cathode collide with gas atoms or molecules in the tube, leading to the removal of one or more electrons from the atoms or molecules.
The e/m value of positive rays for different gases are different but those for cathode rays the e/m values are the same. Justify it.
The e/m values for positive rays from different gases are different because the mass of the positive ions formed in the discharge tube depends on the specific gas used. Positive ions can have different masses, resulting in variations in their charge-to-mass ratios (e/m values). In contrast, cathode rays are composed of electrons, which have a constant mass and charge, so their e/m value remains the same regardless of the gas.
The e/m value for positive rays obtained from hydrogen gas is 1836 times less than that of cathode rays. Justify it.
The e/m value for positive rays obtained from hydrogen gas being 1836 times less than that of cathode rays can be justified by considering the difference in mass between electrons and protons. Hydrogen is the lightest element, and its positive ions are essentially single protons. The mass of a proton is approximately 1836 times greater than that of an electron. Since the e/m value is inversely proportional to mass, the e/m value for hydrogen ions is 1836 times less than that of electrons.
Long Question Chemistry 1st Year Chapter 5 Atomic Structure
Question: Explain Millikan’s oil drop experiment to determine the charge of an electron.
Answer: Millikan’s oil drop experiment was a groundbreaking experiment conducted by American physicist Robert A. Millikan in 1909. This experiment aimed to determine the fundamental charge of an electron and contributed significantly to our understanding of atomic and subatomic particles.
Here’s a step-by-step explanation of Millikan’s oil drop experiment
Apparatus: Millikan used a vertical chamber with transparent sides filled with air. Inside the chamber, he introduced tiny oil droplets by spraying oil into the chamber. These oil droplets became charged as they passed through an atomizer, acquiring excess electrons.
Electric Field: Millikan introduced a uniform electric field in the chamber by applying a voltage across two parallel metal plates placed at the top and bottom of the chamber. This electric field caused the charged oil droplets to move either upward (opposite to gravity) or downward (along with gravity), depending on the sign of the charge on the droplets.
Observations: Millikan observed the motion of the charged oil droplets through a microscope. Some droplets would hang suspended in the chamber, while others would fall under the influence of gravity. The key to the experiment was to find oil droplets that were stationary or moved at constant velocities.
Measurement: By carefully adjusting the voltage applied to the plates and calculating the electric field strength, Millikan was able to measure the force of gravity acting on each oil droplet and the electrical force acting on them due to the electric field.
Equilibrium: Millikan observed that some oil droplets would remain suspended in mid-air, indicating that the electrical force (due to the excess electrons on the droplets) balanced the gravitational force acting on them. In these cases, the oil droplets were neither rising nor falling but remained stationary.
Calculations: By equating the electrical force to the gravitational force and knowing the mass of the oil droplets (measured from their terminal velocity) and the acceleration due to gravity, Millikan was able to determine the charge (q) on each oil droplet.
Quantization of Charge: Millikan repeated the experiment multiple times, measuring the charges on different oil droplets. He observed that the charges on the droplets were always integer multiples of a fundamental charge (e), which led him to conclude that the charge on each oil droplet and, by extension, the charge on an electron was quantized.
Determination of the Electron’s Charge: After analyzing numerous data points, Millikan calculated the charge of an electron to be approximately -1.6 x 10^-19 coulombs, a value that is remarkably close to the currently accepted value.
Millikan’s oil drop experiment provided strong experimental evidence for the quantization of electric charge and helped establish the charge of the electron as a fundamental constant in physics. This experiment played a crucial role in advancing our understanding of the structure of atoms and subatomic particles.
Question: What is J.J Thomson’s experiment for determining e/m value of electron?
Answer: J.J. Thomson’s experiment to determine the charge-to-mass ratio (e/m ratio) of electrons is known as the Cathode Ray Tube (CRT) experiment. This experiment was conducted by Thomson in the late 19th century and played a pivotal role in discovering the existence of electrons and characterizing their properties. Here’s an overview of the experiment:
Apparatus
Cathode Ray Tube (CRT): A vacuum tube with electrodes sealed inside. It consists of a cathode (negatively charged electrode) and an anode (positively charged electrode) at either end, with a narrow, centrally located region called the “anode-cathode gap.”
High Voltage Power Supply: To create a potential difference between the cathode and anode.
Magnetic Field Source: A magnet or electromagnet is placed perpendicular to the path of the electron beam.
Procedure
Evacuation: The CRT is evacuated to create a vacuum inside, so there are no gas molecules to interfere with the experiment.
Electron Emission: When a high voltage is applied between the cathode and anode, it creates an electric field. This electric field exerts an attractive force on the negatively charged electrons in the cathode, causing them to be emitted from the cathode surface. These emitted electrons are referred to as “cathode rays.”
Path of Cathode Rays: The cathode rays travel in a straight line from the cathode to the anode. However, when a magnetic field is applied perpendicular to their path, the cathode rays are deflected.
Deflection of Cathode Rays: The deflection of cathode rays in the presence of a magnetic field is a key observation. By adjusting the strength of the magnetic field and measuring the extent of deflection, Thomson was able to study the behavior of the cathode rays.
Observations and Analysis
Thomson observed that the cathode rays were deflected by the magnetic field, and the degree of deflection depended on the strength of the magnetic field and the speed of the cathode rays.
From these observations, Thomson was able to derive the following key information:
The cathode rays were negatively charged because they were deflected in a direction consistent with the known behavior of negatively charged particles in a magnetic field.
The degree of deflection of the cathode rays was proportional to the charge-to-mass ratio (e/m) of the particles. By measuring the angle of deflection and the known strength of the magnetic field, Thomson could calculate the e/m ratio of the electrons.
Thomson’s experiments showed that the e/m ratio of electrons was the same, regardless of the type of metal used for the cathode or the gas inside the CRT, indicating that electrons were a fundamental particle present in all matter.
J.J. Thomson’s CRT experiment provided crucial evidence for the existence of electrons, their charge, and their charge-to-mass ratio, contributing significantly to the development of the atomic model and our understanding of subatomic particles.
Question: Discuss Chadwick’s experiment for the discovery of neutron. Compare the properties of electron, proton and neutron.
Answer: James Chadwick’s experiment in the early 1930s led to the discovery of the neutron, a subatomic particle with no electric charge and nearly equal in mass to a proton. Here’s an overview of Chadwick’s experiment and a comparison of the properties of electrons, protons, and neutrons:
Chadwick’s Experiment for the Discovery of the Neutron
Apparatus
Polonium Source: Chadwick used polonium-210 (²¹⁰Po) as a source of alpha particles.
Beryllium Target: He directed alpha particles emitted by the polonium source onto a target made of beryllium (⁹Be).
Paraffin Wax Block: Placed around the beryllium target to slow down the emitted particles.
Procedure
Alpha particles (helium nuclei) emitted from the polonium source were directed at the beryllium target.
When the alpha particles struck the beryllium nuclei in the target, they caused the beryllium nuclei to emit particles.
These emitted particles were observed and found to be electrically neutral (no deflection in electric fields) and had roughly the same mass as a proton.
Discovery of the Neutron
Chadwick concluded that the emitted neutral particles were previously unknown subatomic particles, which he named “neutrons.” These neutrons had a mass similar to that of protons but carried no electric charge, which explained why they were not deflected in electric fields.
Comparison of Properties
Electron
Charge: Negative (-1 elementary charge, approximately -1.602 x 10^-19 coulombs).
Mass: Much smaller than protons and neutrons (approximately 9.109 x 10^-31 kilograms).
Location: Found in electron shells surrounding the atomic nucleus.
Role: Involved in chemical reactions and electricity.
Proton
Charge: Positive (+1 elementary charge, approximately +1.602 x 10^-19 coulombs).
Mass: Approximately equal to a neutron (approximately 1.673 x 10^-27 kilograms).
Location: Located in the atomic nucleus.
Role: Determines the element’s identity and is involved in nuclear interactions.
Neutron
Charge: Electrically neutral (no charge).
Mass: Approximately equal to a proton (approximately 1.675 x 10^-27 kilograms).
Location: Located in the atomic nucleus.
Role: Helps stabilize the atomic nucleus by balancing the repulsive forces between positively charged protons.
Question: Rutherford’s atomic model is based on the scattering of a-particles from a thin gold foil. Discuss it and explain the conclusions.
Answer: Ernest Rutherford’s atomic model, often referred to as the Rutherford model or planetary model, was developed based on the famous alpha particle scattering experiment conducted in 1909 in collaboration with Hans Geiger and Ernest Marsden. This experiment provided crucial insights into the structure of the atom and led to the rejection of the earlier plum pudding model proposed by J.J. Thomson.
The Alpha Particle Scattering Experiment
Apparatus
Alpha Particle Source: Rutherford used a radioactive source of alpha particles (helium nuclei), typically radium or polonium.
Gold Foil: A thin gold foil, only a few atoms thick, was used as the target.
Fluorescent Screen: Placed around the gold foil, the screen emitted flashes of light when struck by alpha particles.
Detecting System: Geiger and Marsden designed a detecting system to record the deflection and scattering of alpha particles.
Procedure
Alpha particles were emitted from the radioactive source and directed towards the gold foil.
A fluorescent screen surrounded the gold foil, and the location of the flashes of light on the screen was recorded.
Observations and Conclusions
Rutherford and his team made the following observations during the experiment:
Most Alpha Particles Passed Through: The majority of alpha particles passed through the gold foil without significant deflection. This observation was unexpected, as the plum pudding model suggested that electrons were distributed uniformly throughout the atom.
Some Alpha Particles Were Deflected: However, a small fraction of alpha particles experienced significant deflection or even bounced back in the direction from which they came. This result was astonishing, as it implied that the positive charge in the atom was concentrated in a tiny, dense nucleus.
Conclusions
Based on the alpha particle scattering experiment, Rutherford drew several important conclusions;
The Nuclear Model: Rutherford proposed a new atomic model in which the atom consists of a small, positively charged nucleus at its center. This nucleus contains most of the atom’s mass and is where the protons are located. The electrons, which are negatively charged, orbit the nucleus at a distance.
Mostly Empty Space: The fact that most alpha particles passed through the gold foil without deflection suggested that atoms are mostly empty space. The alpha particles interacted with the small, dense nucleus only rarely.
Explanation for Scattering: The significant deflections and reversals of some alpha particles were explained by the presence of the positively charged nucleus. When alpha particles came close to the nucleus, they experienced strong repulsive forces due to the positive charge, causing them to be deflected or even bounce back.
Rejection of the Plum Pudding Model: Rutherford’s model rejected J.J. Thomson’s plum pudding model, which proposed that positive and negative charges were evenly distributed throughout the atom. The scattering experiment’s results were inconsistent with this model.
Rutherford’s nuclear model of the atom laid the foundation for our modern understanding of atomic structure. It was later refined with the incorporation of quantum mechanics, leading to the development of the Bohr model and, eventually, the quantum mechanical model of the atom.
Question: Give the postulates of Bohr’s atomic model. Which postulate tells us that orbits are stationary and energy is quantized?
Answer: Niels Bohr’s atomic model, known as the Bohr model or the planetary model, was proposed in 1913 to explain the structure of the hydrogen atom and the spectral lines observed in hydrogen’s emission spectrum. The key postulates of Bohr’s atomic model are as follows:
Electron Orbits: Electrons in an atom revolve around the nucleus in well-defined orbits, similar to planets orbiting the sun. These orbits are also referred to as electron shells or energy levels.
Stationary Orbits: Electrons can only exist in certain, specific orbits with fixed radii, known as stationary orbits. In these orbits, electrons do not emit or absorb electromagnetic radiation, and they experience no acceleration. This postulate explains why electrons in Bohr’s model do not continuously spiral into the nucleus, which classical physics would predict.
Quantization of Angular Momentum: Bohr introduced the concept of quantization of angular momentum, stating that the angular momentum of an electron in an orbit is an integral multiple of Planck’s constant divided by 2π (h/2π). Mathematically, it can be expressed as L = n(h/2π), where L is the angular momentum, n is the principal quantum number, and h is Planck’s constant.
Quantization of Energy: Electrons in Bohr’s model occupy specific energy levels or orbits with quantized energies. An electron can only exist in one of these energy levels, and the energy of an electron in a particular orbit is fixed and cannot change spontaneously.
Radiation of Light: Electrons can transition from one energy level to another by either absorbing or emitting energy in the form of electromagnetic radiation (photons). When an electron transitions to a lower energy level, it emits energy in the form of light, and when it absorbs energy, it moves to a higher energy level.
The postulate that tells us that orbits are stationary and energy is quantized is the second postulate: “Stationary Orbits.” According to this postulate, electrons revolve in stable orbits without emitting radiation in the form of electromagnetic waves, as long as they remain in these stationary orbits. When electrons transition between these orbits, energy is either emitted or absorbed in discrete amounts, resulting in the quantization of energy levels in the atom. This quantization of energy levels is a fundamental feature of Bohr’s atomic model and is responsible for the observed line spectra in the emission and absorption spectra of hydrogen and other elements.
Question: What are X-rays? What is their origin? How was the idea of atomic number derived from the discovery of X-rays?
Answer: X-rays are a form of electromagnetic radiation that have higher energy and shorter wavelengths than visible light. They were discovered by the German physicist Wilhelm Conrad Roentgen in 1895. X-rays have various applications in fields such as medicine, industry, and scientific research due to their ability to penetrate matter and produce images of the internal structures of objects.
Origin of X-rays
X-rays are generated when high-energy electrons strike a target material, typically a heavy metal such as tungsten. The process by which X-rays are produced is called X-ray emission or bremsstrahlung. It occurs as follows:
High-energy electrons are accelerated toward the metal target by an electric potential difference (voltage). When these fast-moving electrons approach the positively charged atomic nuclei in the target material, they experience strong electrostatic attraction.
As a result, the electrons undergo rapid deceleration, and their kinetic energy is converted into electromagnetic radiation, including X-rays. This emission of X-rays is a consequence of the sudden change in the velocity of the electrons as they interact with the atomic nuclei in the target.
The emitted X-rays have a continuous spectrum of energies, ranging from low to high, depending on the energy of the incident electrons. The characteristic properties of the X-ray spectrum depend on the specific target material and the energy of the incident electrons.
Derivation of Atomic Number from X-ray Discoveries
The discovery of X-rays had a profound impact on our understanding of atomic structure and led to the derivation of the concept of atomic number. Here’s how it happened:
X-ray Spectra: Scientists observed that the X-ray spectra of different elements were unique. Each element produced a characteristic X-ray spectrum, which included both a continuous spectrum (due to bremsstrahlung) and sharp lines at specific wavelengths (characteristic X-rays).
Moseley’s Law: In 1913, Henry Moseley, a British physicist, conducted experiments on the X-ray spectra of various elements. He found that the square root of the frequency (or reciprocal wavelength) of the characteristic X-rays was directly proportional to the atomic number (Z) of the element.
Mathematically, Moseley’s law can be expressed as:
√ν = k(Z – b)
Where ν is the frequency of the characteristic X-rays, k and b are constants, and Z is the atomic number of the element.
Derivation of Atomic Number: Moseley’s law provided a direct experimental method for determining the atomic number of elements. By measuring the frequencies of characteristic X-rays and applying Moseley’s law, scientists could accurately determine the atomic number of an unknown element.
This discovery was a significant breakthrough because it provided a more precise and systematic method for organizing the elements in the periodic table based on their atomic numbers rather than their atomic masses. It resolved some anomalies in the periodic table and led to the modern understanding of the arrangement of elements in the periodic table, where elements are ordered by increasing atomic number, and their properties are periodic.
Question: How does the Bohr’s model justify the Moseley’s equation?
Answer: Bohr’s model of the hydrogen atom can be used to provide a qualitative justification for Moseley’s empirical equation, which relates the frequency of characteristic X-rays to the atomic number of an element. Although Bohr’s model specifically applies to hydrogen, it provides insights into the behavior of electrons in atoms and their energy levels, which can be extended to explain certain aspects of X-ray spectra and Moseley’s law. Here’s how Bohr’s model justifies Moseley’s equation:
Bohr’s Model of the Hydrogen Atom
Bohr’s model incorporates several key principles;
Electrons in an atom occupy discrete energy levels or orbits.
Electrons can transition between these energy levels by absorbing or emitting energy in the form of photons of specific frequencies.
The energy of an electron in a particular orbit is quantized and can be calculated using the following equation for the hydrogen atom:
E = -13.6 eV / n²
Where:
E is the energy of the electron in electronvolts (eV).
n is the principal quantum number, which represents the energy level (n = 1, 2, 3, …).
Justification of Moseley’s Equation:
Emission Spectra: Bohr’s model explains the line spectra observed when electrons transition between energy levels in hydrogen. When an electron drops from a higher energy level (n₂) to a lower energy level (n₁), it emits a photon of specific frequency (ν) given by:
ΔE = E₂ – E₁ = hν
Where ΔE is the energy difference between the two levels, h is Planck’s constant, and ν is the frequency of the emitted photon.
Connection to X-ray Spectra: Moseley’s law is derived from the observation that the frequency of characteristic X-rays produced by different elements is directly proportional to the square root of their atomic number (Z).
√ν = k(Z – b)
Bohr’s model can be connected to this by considering that the X-ray emission lines result from transitions of inner-shell electrons (e.g., K and L shells) to outer vacant electron orbits. When an inner-shell electron falls to a lower energy level (closer to the nucleus), it emits an X-ray photon. The energy of this X-ray photon is determined by the difference in energy levels, which is related to the atomic number.
Qualitative Explanation: Bohr’s model provides a qualitative explanation for why X-ray frequencies increase with atomic number. As the atomic number (Z) increases, the effective nuclear charge experienced by inner-shell electrons increases. This increased attraction causes inner-shell electrons to be bound more tightly, resulting in higher energy transitions and, consequently, higher frequency X-rays.
While Bohr’s model does not provide precise quantitative predictions for X-ray frequencies, it illustrates the fundamental concept that X-ray emissions are related to the electronic structure of atoms and the quantization of energy levels, which is consistent with Moseley’s empirical observations. Moseley’s law, derived from experiments, quantitatively connects the X-ray frequencies to the atomic number, providing a powerful tool for characterizing elements based on their X-ray spectra.