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Structure of Atom Class 11 Chemistry Chapter 2 One Shot | New NCERT CBSE

LearnoHub - Class 11, 12・2 minutes read

The video explains the structure of an atom, detailing the nucleus with protons and neutrons and the orbiting electrons. It also covers the historical development of atomic models by scientists and the limitations of the Bohr model, leading to the quantum mechanical model of the atom.

Insights

Atoms consist of protons, neutrons, and electrons, with protons and neutrons in the nucleus and electrons orbiting around it.
The historical development of atomic models involved scientists like John Dalton, Thomson, Rutherford, Bohr, and the quantum mechanical model.
Subatomic particles, including electrons, protons, and neutrons, were discovered through experiments like the cathode ray discharge tube and the Oil Drop Experiment.
Rutherford's model of the atom, emphasizing a positively charged nucleus surrounded by electrons, lacked explanations for atom stability and electronic structure.
The dual nature of electromagnetic radiation as both particles (photons) and waves is evident in phenomena like the photoelectric effect and atomic spectra.
The Bohr Model explained the hydrogen spectrum and circular electron orbits but was limited to species with one electron, leading to the quantum mechanical model's development.
The quantum mechanical model describes electron behavior in three-dimensional space, with atomic orbitals defined by wave functions and quantum numbers determining energy levels and orbital characteristics.

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Recent questions

What are atoms made of?
Atoms are composed of protons, neutrons, and electrons.

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Summary

00:00

"Atom Structure Simplified for Learners"

Matter is composed of molecules, which are in turn made up of atoms.
The video discusses the structure of an atom, aiming to simplify the concept for learners.
The narrator, Roshni from Ln Hub, introduces the free learning platform for various subjects.
Atoms are defined as the basic unit of matter, containing a nucleus with positively charged protons and neutral neutrons.
Electrons, negatively charged particles, orbit around the nucleus within an atom.
The nucleus of an atom is incredibly small, with a comparison made to a ball's size and a 5 km radius for the atom.
The lesson delves into the historical development of atomic models by scientists like John Dalton, Thomson, Rutherford, Bohr, and the quantum mechanical model.
The focus of the lesson is on understanding the structure of atoms, the constituent particles within atoms, and subatomic particles.
The discovery of subatomic particles began with experiments on electric discharge through gases, leading to the identification of electrons.
The experiment to discover electrons involved a cathode ray discharge tube, where observations revealed negatively charged particles moving from the cathode to the anode.

16:39

"Cathode Rays: Charge to Mass Ratio Discovery"

Television screens had a coating of fluorescent or phosphorescent material that glowed when cathode rays struck them.
Images on television were visible due to the glowing of the screen caused by cathode rays passing through the tube.
Cathode rays moved in a straight line from cathode to anode without any external fields applied.
Applying electric or magnetic fields caused deviation in the path of cathode rays towards the positive side due to their negative charge.
Cathode particles, regardless of the nature of the electrodes or gas inside the tube, were negatively charged electrons.
J.J. Thomson calculated the charge to mass ratio of electrons experimentally by manipulating the strength of electric and magnetic fields.
Thomson observed that the deviation of cathode ray particles depended on the negative charge, mass of the particle, and the strength of electric and magnetic fields.
The hypothetical story of Alec illustrated how the deviation of particles depended on their negative charge, mass, and the strength of electric and magnetic fields.
Thomson's experiments led to the determination of the charge to mass ratio of electrons, approximately 1.75 * 10^11 coulombs per kilogram.
R.A. Millikan conducted the Oil Drop Experiment to determine the charge on an electron, utilizing electric fields to observe the motion of oil droplets between electrodes.

32:13

"Milikan's Oil Drop Experiment and Atomic Models"

Milikan conducted experiments on oil droplets to determine their charges.
The charge on oil droplets was always an integral multiple of a specific value, denoted as e.
The value of e was found to be 1.6 * 10^-19 coulombs, representing the charge on an electron.
Through the Milikan Oil Drop experiment, the charge-to-mass ratio of electrons was determined.
The discovery of the electron and its charge was followed by the discovery of the proton.
Protons were found to be positively charged atomic particles, distinct from electrons.
The discovery of the proton involved a setup similar to the electron discovery, using a cathode ray tube.
Positively charged particles, known as canal rays, were observed to depend on the gas present in the tube.
Neutrons, neutral subatomic particles, were discovered by James Chadwick in 1932.
The atomic models proposed by Thomson and Rutherford depicted the distribution of positive charges and electrons within atoms.

48:11

Rutherford's Alpha Particle Experiment and Model

Alpha particles have the lowest penetrating power among Alpha, Beta, and Gamma particles.
Rutherford's experiment involved alpha particles, a thin gold foil, and a fluorescent screen.
The setup included a radioactive source emitting alpha particles towards the gold foil.
Observations from the experiment showed most alpha particles went undeflected, some deflected, and a few bounced back.
Rutherford's model of the atom suggested a positively charged nucleus at the center with electrons orbiting around it.
The model lacked explanations for atom stability, electronic structure, and the dual nature of matter.
Atomic number (protons) and mass number (protons + neutrons) are crucial in identifying elements.
Isotopes are elements with the same atomic number but different mass numbers, while isobars have different atomic numbers but the same mass number.
Isotopes find applications in nuclear reactors, cancer treatment, and goiter treatment.
Rutherford's model emphasized the positive charge and mass concentration in the nucleus of an atom.

01:02:54

Evolution from Rutherford to Bohr's Atom

Rutherford's model of the atom placed electrons moving around the nucleus in circular paths, similar to the solar system.
The model failed to explain the stability of the atom due to the electrostatic force of attraction between the negatively charged electrons and positively charged nucleus.
Maxwell's theory suggested that accelerating charged particles emit radiation, which should cause electrons to spiral into the nucleus, contradicting Rutherford's model.
Rutherford's model couldn't explain the stability of the atom, the electronic structure of atoms with multiple electrons, or atomic spectral lines.
The limitations of Rutherford's model led to its failure, paving the way for Bohr's model of the atom.
Bohr's model was based on the dual behavior of electromagnetic radiation and atomic spectra, addressing the shortcomings of Rutherford's model.
Wavelength, frequency, wave number, amplitude, and period are crucial parameters in understanding electromagnetic waves.
Electromagnetic waves, as described by Maxwell, consist of alternating electric and magnetic fields that propagate without the need for a medium.
Electromagnetic waves, like light, are represented by properties such as wavelength, frequency, and wave number.
Different types of electromagnetic radiation exist based on varying wavelengths, collectively forming the electromagnetic spectrum.

01:20:01

"Understanding the Electromagnetic Spectrum and Radiation"

Visible light can be dispersed into different colors by passing it through a prism, creating a spectrum of violet, indigo, blue, green, yellow, orange, and red.
The electromagnetic spectrum consists of different types of electromagnetic radiation, arranged based on their frequency and wavelength.
Radiations like gamma rays, x rays, and ultraviolet radiation have high frequencies and short wavelengths, while radio waves have low frequencies and long wavelengths.
The visible spectrum includes colors from violet to red, with violet having a shorter wavelength and red having a longer wavelength.
As wavelengths decrease in the electromagnetic spectrum, the penetration power of radiations increases, with X-rays and gamma rays having very small wavelengths.
The scale of wavelengths ranges from atomic and molecular levels for X-rays and gamma rays to microorganism levels for infrared and radio waves.
Visible light cannot penetrate certain materials due to its longer wavelength, unlike X-rays and gamma rays with shorter wavelengths.
X-rays, with their small wavelengths, can penetrate materials like bags in airport security checks, revealing their contents.
The particle nature of electromagnetic radiation is evident in phenomena like black body radiation, the photoelectric effect, and atomic spectra, which wave theory cannot fully explain.
Heating solids causes them to emit radiation, with the emitted radiation changing in color and frequency as the solid heats up, demonstrating the particle nature of radiation.

01:35:22

Black Body Radiation: Temperature and Quanta

Ordinary objects are imperfect absorbers of radiation, not absorbing all radiation perfectly.
An ideal body, a black body, absorbs and emits all frequencies of radiation perfectly.
Experimental observations showed that the black body spectrum depends only on temperature, not on the material it is made of.
As temperature increases, a black body emits more energy at all wavelengths.
The peak wavelength of the black body spectrum becomes shorter as temperature increases.
The black body spectrum always becomes smaller at the left-hand side as temperature increases.
Classical theory predicted that at very high temperatures, radiation energy would go to infinity, but this was not observed experimentally.
Planck's quantum theory explained the black body radiation results by stating that energy is emitted or absorbed in discrete quantities called quanta.
At high temperatures, black bodies prefer to emit a large number of low-energy quanta rather than a single high-energy quanta.
The preference for emitting low-energy quanta at high temperatures explains the shape of the black body spectrum, with a drop in the low-wavelength region and a peak shifting towards lower wavelengths as temperature increases.

01:52:34

"Photoelectric Effect: Light's Particle-Wave Duality"

Photoelectric effect explained by Hearts through an experiment where electrons were observed to be emitted from a metal surface when exposed to light.
The phenomenon was named the photoelectric effect, where electrons were ejected from the metal surface upon light exposure.
Observations showed that the number of ejected electrons was proportional to the intensity and brightness of the light.
Different metals had unique minimum frequencies, termed threshold frequencies, below which the photoelectric effect was not observed.
Experimentation revealed that emitting light of different frequencies on the same metal surface could induce or inhibit the photoelectric effect.
Electrons ejected from the metal surface due to light exposure carried kinetic energy, dependent on the frequency of the light.
Einstein's explanation of the photoelectric effect introduced the concept of photons as particles of light interacting with electrons to release them from the metal surface.
The energy of a photon falling on a metal surface was denoted by E = hν, with a minimum energy requirement known as the work function for each metal.
The photoelectric equation, hν = h0 + 1/2 mv^2, illustrated the relationship between the energy of the photon, work function, and kinetic energy of the ejected electron.
Electromagnetic radiation exhibited dual behavior, acting as both particles and waves, with interactions with matter showcasing its particle-like characteristics and propagation displaying wave-like properties.

02:08:48

"Calculating Energy and Velocity in Spectroscopy"

The relation between frequency and wavelength is given by the formula nya = c / ladda.
The value of h is 6.62 * 10^-34 joules second, and the value of c is 3 * 10^8 meters per second.
Energy can be calculated by dividing 3 * 10^8 by 4 * 10^-7 meters.
Energy can be converted to electron volts by dividing by 1.6 * 10^-19 joules.
The kinetic energy of emitted electrons can be calculated using the photoelectric equation.
The kinetic energy is determined by subtracting the energy values of h nya and h.
The velocity of the photoelectron can be calculated using the kinetic energy and mass of an electron.
The mass of an electron is 9.1 * 10^-31 kg.
The spectrum of white light passing through a prism results in seven colors due to different wavelengths bending at different angles.
The emission spectrum is created when atoms or molecules emit radiation after returning from an excited state to the ground state, resulting in colored lines.

02:24:47

Comparing Absorption and Emission Spectra in Spectroscopy

Absorption and emission spectra are compared by matching photos to reveal similarities.
Spectroscopy involves studying emissions and absorption spectra.
Spectroscopic methods are valuable for obtaining detailed information about elements.
Visible light passing through a prism creates a continuous spectrum.
Emission spectrum is observed when hot gas emits radiation.
Cold gas produces an absorption spectrum when radiation passes through a prism.
Line emission spectrum is unique to each element, acting as a fingerprint for identification.
Spectroscopic methods have led to the discovery of elements like helium, rubidium, gallium, and scandium.
Hydrogen's spectrum is simpler compared to other elements, with distinct series of spectral lines.
Balmer series in hydrogen's spectrum falls in the visible light region, with specific transitions between energy levels.

02:41:02

Bohr Model: Energy States and Orbits

Energy states in orbits are named by Bor, who described them as concentric circles with different radii.
Energy states are numbered starting from n=1, known as the ground state, with subsequent states at n=2, n=3, n=4, and so on.
The radius of each orbit increases with higher energy states, calculated as rn=a*n^2, with a value of 52.9 picometers for the first stationary state.
Electrons absorb energy to move from lower to higher energy states and release energy when transitioning back.
The energy of an electron in an energy state increases as it moves away from the nucleus due to the negative sign in the formula.
The difference in energy between two states can be calculated using the formula delta e = -rh1/a * (1/nf^2 - 1/ns^2).
The Bohr Model applies not only to hydrogen but also to hydrogen-like species, with energy calculations adjusted by multiplying by the atomic number squared.
Transitioning electrons emit radiation with a frequency calculated as nya = delta e/h, where h is Planck's constant.
Bohr postulated that the angular momentum of an electron in an orbit is quantized, with values being integral multiples of h/2π.
The Bohr Model can be used to explain the hydrogen spectrum and potentially derive the Rydberg formula by manipulating energy gap equations.

02:57:19

"Rydberg Formula and Quantum Mechanics Explained"

The constant value of r is 2.18 * 10^18, divided by h, where h is Planck's constant (6.62 * 10^-34) and c is the speed of light (3 * 10^8).
Calculating these values results in 1.09 677 * 10^7, which is crucial for the Rydberg formula.
The Rydberg formula involves r * 1 / n1^2 - 1 / n2^2, specifically for the hydrogen line spectrum.
The formula for the maximum number of emission lines possible when an excited electron in a hydrogen atom at n = 6 drops to the ground state is n * (n - 1) / 2, resulting in 15 possible transitions.
To calculate the wave number for the longest wavelength transition in the Balmer series of atomic hydrogen, the formula 1 / λ = RH * (1 / n1^2 - 1 / n2^2) is used, with n1 = 2 and n2 = 3 resulting in 1.52 * 10^6 meters.
Bohr's model for hydrogen explains the emission spectrum and the circular orbits of electrons around the nucleus.
The relationship between r and Plank's constant and the speed of light is r = r / (hc^3).
Bohr's model was limited to explaining spectra for species with only one electron, leading to its eventual failure.
The development of the quantum mechanical model of the atom was influenced by the dual behavior of matter and the Heisenberg Uncertainty Principle.
De Broglie's hypothesis in 1924 proposed that all matter exhibits both particle and wave behavior, with larger objects having shorter wavelengths making their wave nature undetectable.

03:13:14

Heisenberg's Uncertainty Principle and Quantum Mechanics

Heisenberg's Uncertainty Principle states that the exact position and momentum of an electron cannot be determined simultaneously.
Mathematically, the principle can be represented as delta x * delta px = h / 4p, where delta x is the uncertainty in position and delta px is the uncertainty in momentum.
The uncertainty in position and momentum is inversely proportional, meaning if one is small, the other is large.
The principle is significant for microscopic objects but negligible for macroscopic objects due to their mass.
An example problem involves calculating the uncertainty in position of a golf ball with given mass and speed.
Another problem involves calculating the wavelength of an electron with given kinetic energy.
The limitations of the Bohr model led to the development of the quantum mechanical model of the atom.
The quantum mechanical model considers the dual behavior of matter and describes the movement of electrons in three-dimensional space.
Quantum mechanics is a branch of science that considers the wave-particle duality of matter.
The Schrödinger equation is a fundamental equation in quantum mechanics that describes the behavior of matter in three-dimensional space.

03:29:44

Electron Energy Levels and Atomic Orbitals

Possible values of e satisfy the equation for energy levels of electrons in atoms, indicating quanta-based energy levels.
The wave nature of electrons is evidenced by elevated energy levels and the uncertainty principle regarding electron paths in an atom.
Atomic orbitals are defined by the wave function psi for n electrons in an atom, with each orbital having a specific energy level.
Multi-electron atoms have fixed electrons in various orbitals in increasing energy order, ensuring stability.
The probability of finding an electron at a point within an atom is proportional to the square of the orbital wave function, indicating the probability density.
The Schrödinger equation provides solutions for energy levels and atomic orbitals of electrons in atoms.
Quantum numbers, including the principal quantum number (n), azimuthal quantum number (l), and magnetic quantum number (m), define the size, shape, and orientation of orbitals.
The principal quantum number (n) determines the shell of the electron, with different values indicating different shells.
The azimuthal quantum number (l) determines the shape of the orbital, with values ranging from 0 to n-1.
The magnetic quantum number (m) indicates the spatial orientation of the orbital with respect to coordinate axes, defining the shape and orientation of the orbital.

03:45:22

Quantum Numbers Determine Electron Orientations in Atoms

The principal quantum number, denoted as l or azimuthal quantum number, can have values ranging from 0 to n - 1.
For a specific value of l, the magnetic quantum number, denoted as ml, can have 2l + 1 possible values, determining the orientations.
The ml values for different l values vary, with ml ranging from -l to +l, such as -1, 0, +1 for l = 1.
The number of possible orientations increases with higher quantum numbers, with s having one orientation, p having three, and d having five possible orientations.

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