Login Get started

Plus One Chemistry | Chapter 2 | Structure Of Atom | Oneshot | Exam Winner

Exam Winner Plus One・152 minutes read

The chapter delves into the Structure of Atom and the properties and behaviors of Cathode and Anode Races in understanding subatomic particles, with specific experiments and models discussed by Thomson and Rutherford. It also covers topics such as the hydrogen spectrum, electromagnetic radiation, the photoelectric effect, and the quantum mechanical model, emphasizing the significance of quantum numbers and electron configurations.

Insights

Part One of the text covers foundational Chemistry concepts, including the properties of Cathode and Anode Races, subatomic particles, and the Rutherford model of the atom.
The Cathode Ray Discharge Tube experiment revealed key characteristics of cathode rays, such as their straight-line travel and deflection by electric and magnetic fields, leading to the discovery of electrons by JJ Thomson.
Anode Races, studied similarly to Cathode Races, exhibit positive charges and are influenced by the nature of the gas, unlike Cathode Races, which are gas-independent.
The text explores the evolution of atomic models, from Thomson and Rutherford's contributions to the Bohr model's limitations and the subsequent development of the quantum mechanical model.
Quantum theory, as introduced by Planck, explains the quantized energy of radiation through photons and the photoelectric effect, demonstrating the relationship between light intensity, frequency, and electron emission.
Understanding quantum numbers, electron configuration rules, and the behavior of subatomic particles is crucial in comprehending the structure of atoms and their electron orbitals, with practical applications in determining electron locations and properties.

Get key ideas from YouTube videos. It’s free

Recent questions

What are the subatomic particles?
Electrons, protons, and neutrons.
How did Rutherford discover the nucleus?
Gold foil experiment with alpha particles.
What is the photoelectric effect?
Emission of electrons when hit by light.
What is the Bohr model of the atom?
Electrons orbit nucleus in fixed paths.
What is the quantum mechanical model?
Modern approach to electron behavior in atoms.

Related videos

Summary

00:00

"Structure of Atom: Cathode and Anode Races"

The chapter being discussed is the Structure of Atom, divided into three parts: Part One, Part Two, and Part Three.
Part One covers basic concepts of Chemistry, including Properties of Cathode Rise, Anode Rise, subatomic particles, and the Rutherford model of the atom.
Specific topics in Part One include the hydrogen spectrum, Quantum Theory, Photoelectric Effect, and the experiment conducted by JJ Thomson with the Cathode Ray Discharge Tube.
The Cathode Ray Discharge Tube experiment involved a cathode and anode connected by a wire, high voltage, gas fill, and a fluorescent material like zinc sulfide.
Cathode races travel in a straight line but can be deflected by electric and magnetic fields, with a negative charge on the cathode and a positive charge on the anode.
The properties of Cathode Rise are gas-independent, with a constant Charge to Mass Ratio for all gases, reflecting the discovery of electrons by JJ Thomson.
Anode Rise is studied similarly to Cathode Rise, with an experiment conducted by Goldstein involving the movement of anode races towards the cathode.
Anode races are produced when gas is struck by electric fields, moving towards the cathode and deflecting in the presence of electric and magnetic fields.
The charge on the Anode Race is positive, opposite to the negative charge on the Cathode Race, with both races depending on the gas and traveling in straight lines.
Understanding the properties and behaviors of Cathode and Anode Races is crucial in comprehending the Structure of Atom and its subatomic particles.

14:57

"Gas Determines Anode Race Properties"

Cathode race does not depend on the nature of gas, while anode race does.
The properties of the anode race change based on the nature of the gas taken.
The E by M ratio will vary depending on the gas taken for the anode race.
Charge to mass ratio differs for cathode races and properties of the anode race.
Cathode race is positive and negative, while anode race depends on the cathode E by M ratio.
Protons race through the cathode, and electrons through the anode race.
Subatomic particles include electrons, protons, and neutrons.
The electron was discovered by J. Thomson, the neutron by Chadwick, and the proton by Rutherford.
The mass of an electron is 9.1 * 10^-31, while a proton's mass is 1.6 * 10^-27.
Isotope, isobar, and isotone are terms related to changes in atomic structure, with examples provided for better understanding.

30:42

Rutherford's Atom Model and Limitations

A new model of the atom was presented, with Rutherford and JJ Thomson as key figures.
Rutherford experimented with gold foil and alpha particles to observe their behavior.
Alpha particles passed through the gold foil without deflection, were deflected at small angles, and very few were deflected back.
Rutherford's observations led to three conclusions: atoms have empty space, positive charges are concentrated in a small nucleus, and there is an electrostatic force of attraction between the nucleus and electrons.
Rutherford's model included three postulates: electrons surround the nucleus, positive charge and mass are concentrated in the nucleus, and there is an electrostatic force of attraction between the nucleus and electrons.
Maxwell challenged Rutherford's model, questioning the stability of electrons orbiting the nucleus.
Rutherford admitted he did not know if electrons would fall into the nucleus due to energy loss from radiation emission.
Rutherford's model had limitations in explaining the stability of atoms, the hydrogen spectrum, and the electronic structure of atoms.
The chapter covered topics such as the properties of cathode rays, subatomic particles, models by Thomson and Rutherford, and the limitations of Rutherford's model.
Equations were introduced to explain the decrease in energy of electrons as they orbit the nucleus and emit radiation, leading to the concept of electromagnetic radiation.

47:05

Understanding Electromagnetic Radiation and Photoelectric Effect

Maxwell's theory explains the energy of electrons falling to the center and emitting radiation.
Radiation has three properties: frequency, wavelength, and wave number.
The relationship between frequency, wavelength, and wave number is crucial in understanding radiation.
The equations c = new lambda, new bar = 1/lambda, and new bar = 1/lambda are fundamental in studying electromagnetic radiation.
Planck's quantum theory introduces the concept of quantized energy in radiation, with photons representing packets of energy.
The photoelectric effect involves photons hitting certain metals and emitting electrons, showing only for specific metals like potassium and cesium.
The threshold frequency is essential for electrons to be emitted from metals when hit by light.
The intensity of light hitting metals affects the number of emitted electrons, with kinetic energy increasing as frequency rises.
Kinetic energy is directly proportional to frequency in the photoelectric effect.
Metals like potassium, rubidium, and cesium exhibit the photoelectric effect, showing the relationship between light intensity and electron emission.

01:04:10

"Exploring Energy and Spectra in Physics"

Minimum frequency required for the energy of a photon is called H New Zero
The energy needed to overcome the force between the nucleus and electron is H New Zero
Kinetic energy is half mv squared to save the electron from the nucleus's grip
The equation for total energy is H New Zero plus half mv squared
The dual nature of electromagnetic radiation is explained by Planck's Quantum Theory and the photoelectric effect
The absorption spectrum shows dark lines in a colored background
The emission spectrum displays colored lines in a dark background
The spectrum of hydrogen is observed when light passes through it, causing excitation of electrons
Excitation occurs when electrons absorb energy and jump to higher energy levels
Excited electrons eventually return to lower energy levels, releasing energy in the process

01:21:16

Hydrogen Spectrum: Energy Levels and Radiation

Energy levels in hydrogen atoms are affected by radiation when electrons jump between them.
The resulting series of lines in the hydrogen spectrum are formed by adding radiations.
Different series within the hydrogen spectrum include Lyman, Balmer, Passion, Bracket, and Fund.
Electrons in hydrogen atoms absorb and emit radiation when struck by light.
Jumping between energy levels in hydrogen atoms results in the emission of different spectral lines.
The Lyman series involves electron jumps to the first energy level.
Balmer, Passion, Bracket, and Fund series correspond to jumps to higher energy levels.
The wavelength of the first line in the Lyman series can be calculated using the formula new bar = 1/λ = RH * 1/n1² - 1/n2.
The threshold frequency for a metal can be determined to calculate the kinetic energy of emitted electrons.
Observations related to the wave nature of electromagnetic radiation cannot be explained by particle nature.

01:36:52

"Atomic Structure and Spectral Lines Explained"

To watch the series, go to Very Simple N One and add the number of lines.
The first line should be added to Apo N One, and the result should be added to the first line.
The answer to 10967 * 1/n1 square minus 1/n2 is obtained by adding one plus one to so one by so.
Lambda is equal to wavelength, and the answer should be taken reciprocally.
The Lyman Series is discussed, and the second line of the Balmer series is mentioned.
The number of protons, electrons, and neutrons in a species are equal to 17, 18, and 18, respectively.
The Bohr Model is detailed, emphasizing the fixed circular path of electrons around the nucleus.
The angular momentum of the electron is explained as MVR equal to NH bi 2π.
The equation to find the frequency of radiation is new equal to delta e by h, following the Bohr frequency rule.
The energy of the orbit is calculated using the equation -RH by N square, with the radius equal to A zero in n square by sed.

01:52:43

Quantum Mechanics: Wave Nature and Uncertainty

The text discusses the Seaman effect and the Stark effect in relation to the hydrogen spectrum, mentioning Lyman, Balmer, Paschen, Brackett, and Pfund series.
It explains the splitting of spectral lines in an electric field and magnetic field, detailing the impact of mass on wavelength according to the de Broglie equation.
The text delves into the wave nature of particles, emphasizing that microscopic particles have longer wavelengths due to their negligible mass.
It introduces the de Broglie equation, lambda equals h by mv, highlighting the inverse relationship between mass and wavelength.
The text explores Heisenberg's Uncertainty Principle, stating that it is impossible to determine both the position and velocity of a microscopic particle simultaneously.
It elaborates on the uncertainty in position and velocity, emphasizing the limitations in determining these parameters accurately.
The text concludes by discussing the failure of Bohr's atomic model due to contradictions with Heisenberg's principle and de Broglie's wave nature theory.
It mentions the rejection of the Bohr model by J.J. Thomson, Rutherford, and Maxwell, leading to the development of the quantum mechanical model.
The quantum mechanical model is introduced as a new approach to understanding the behavior of electrons in atoms, replacing the Bohr model.
The text emphasizes the significance of the quantum mechanical model in modern physics, highlighting its role in explaining the behavior of electrons in atoms.

02:08:49

"Quantum Numbers Define Electron's Location and Orientation"

The search for the electron is conducted using the Schödinger equation.
The Schödinger equation provides the wave function, represented by psi squared, to determine the electron's probable location.
The orbital, indicated by psi squared, is where the electron is most likely to be found in space.
Quantum numbers, including principal, azimuthal, magnetic, and spin quantum numbers, define the electron's location and orientation.
The principal quantum number (n) determines the shell in which the electron resides.
The azimuthal quantum number (l) specifies the subshell within the shell where the electron is located.
The magnetic quantum number (m) indicates the orbital's orientation within the subshell.
The spin quantum number (s) describes the electron's spin direction.
The values of quantum numbers are interconnected, with specific constraints on their possible combinations.
The connection between the values of N and L, as well as the restrictions on M values based on L, are crucial in determining the electron's location and orientation.

02:31:08

Orbitals and Electronic Configuration in Chemistry

L value said M value is zero
If M value is zero, then L is zero
If the value is one, then the M value is one
Look at the L value of one
If M value is minus one zero one, then M value is one and L value is one
Each value is set from zero to one
Are they connected?
If the value is two, then it is now L
If the value is three, then the M value is minus Three -2 -1 Zero One Two Three
S Value Total Two are S value total two are two
Only two values s have a total of s
Either clockwise or anti-clockwise
N Value L Value N Value S Value
Values also said that he will give gave
If you get the value, you get the L value and so on
After getting a room, it is in each room of chairs
The number is the room inside the house
The chair is a room in the first house
The second house has two rooms and one of them
There are five chairs in the given room
So many orbitals are possible for it
Shell Subshell Orbital
Now each In the chair, two people each spin one clockwise and one anti-clockwise
S Orbital Shape of Spherical P Orbital Shape of Dumbbell
S Orbital has one chair, P has three, D has five, and F has seven
D has five orbitals and one d XY, one DY Z, one DX Zed, and DX Y
The shape of the orbital is double
The trick to draw is very simple
Remember that s has an orbital, three orbitals for p, and five for d
The graph of S will also be set
The electronic configuration is going to go
The electronic configuration is going to go
You need to know three rules, one off
The Woe Principle
What is the Offbo Principle of Energies?
Orbitals are filled in order of What is increasing energies off?
Principle orbitals are filled in order of energy increasing energy
Orbitals must be filled in order
You are the increasing order of orbitals
One was the first to study in a small class
Will fill in S then fill in two S
Will do and fill in two p
One has only one S and two has two
There are three out of three rooms, two S and two P
There are three S and three P rooms
D and then three S and three P and three
For S before writing D
Write three d because four 3D has more energy than S
K so energy should be written in increasing order

02:47:00

"Quantum Numbers and Electron Configuration Essentials"

The text discusses the order of energy increasing, emphasizing the importance of learning the OneS2S2P Four before three S three P three D S sequence.
It introduces the Police Exclusion Principle, stating that no two electrons in the same atom can have the same set of four quantum numbers.
The text explains the concept of quantum numbers, detailing the N, L, M, and S values of electrons.
It highlights the rule that no more than two electrons can occupy the same orbital, emphasizing the importance of electron pairing.
The text delves into the electron configuration of atoms, particularly focusing on chromium and copper, explaining the significance of half-filled and completely filled subshells for stability.
It discusses the stability of completely filled and half-filled subshells, attributing it to greater exchange energy and symmetry.
The text touches on the importance of understanding points and quantum numbers, emphasizing the Police Exclusion Principle and Hun's rule of maxima.
It encourages students to practice questions related to the topics discussed to enhance their understanding and familiarity with the concepts.
The text provides examples of quantum numbers for the valence electron of a sodium atom, illustrating the values of N, L, M, and S.
It presents a question on quantum numbers, challenging students to identify which set of quantum numbers is not allowed based on the rules discussed.

03:01:54

Quantum Numbers and Electron Configuration Explained

The electronic configuration for an atom with n=4 and l=0 is discussed, emphasizing the importance of following Hund's Rule and the Pauli Exclusion Principle.
A microscope using suitable photons is employed to locate an electron in an atom at a distance of 4.4 angstroms, highlighting the uncertainty involved in measuring its velocity.
The number of unpaired electrons in nickel is determined based on its atomic number, with the application of Hund's Rule to identify the unpaired electrons.
The quantum numbers are explained as crucial in determining the address and size of an electron from the nucleus, with the principal quantum number providing the distance and the orbital angular momentum indicating the shape of the electron's orbit.

Try it yourself — It’s free.