Text Book
Modern's abc of Chemistry by Dr. S.P. Jauhar for Class XI CBSE
Sections in the Chapter
3.1 Fundamental particles
3.2 Thomson’s Atomic Model
3.3 Rutherford’s Scattering Experiment
3.4 Concept of Atomic Number and Discovery of Neutron
3.5 Developments Leading to the Bohr Model of Atom
3.6 Nature of Light and Electromagnetic Radiation
3.7 Particle Nature of Electromagnetic Radiation and Planck’s Quantum Theory
3.8 Atomic Spectra
3.9 Failure of Rutherford Model
3.10 Concept of Energy Levels or Orbits
3.11 Modern Concept of Structure of an Atom: Quantum Mechanical Model
3.12 Wave Mechanical Model of Atom and Concept of Atomic Orbital
3.13 Quantum Numbers
3.14 Pauli’s Exclusion Principle
3.15 Shapes of Orbitals or Boundary Surface Diagrams
3.16 Energy Level Diagram for Electrons in an Atom
3.17 Electronic Configuration of Atoms
Conceptual Questions with Answers: 35
Additional Numerical Problems for Practice: 8
Revision Exercises
Very Short Answer questions 16
Short Answer Questions 42
Long Answer Questions 3
Competition File
Numerical Problems 20
Objective Questions: 47
Fill in the blanks: 10
True or False: 10
Study Plan
Day 1
3.1 Fundamental particles
3.2 Thomson’s Atomic Model
3.3 Rutherford’s Scattering Experiment
Day 2
3.4 Concept of Atomic Number and Discovery of Neutron
Ex. 3.1 to 3.3
Practice Problems 3.1 to 3.9
Day 3
3.5 Developments Leading to the Bohr Model of Atom
3.6 Nature of Light and Electromagnetic Radiation
Ex. 3.4, 3.5,3.6
P.P. 3.10 to 3.15
Day 4
3.7 Particle Nature of Electromagnetic Radiation and Planck’s Quantum Theory
Ex. 3.7 to 3.16
Day 5
P.P 3.16 to 3.24
Day 6
3.8 Atomic Spectra
Ex. 3.17, 3.18
P.P. 3.25 t 3.27
3.9 Failure of Rutherford Model
Day 7
3.9 Failure of Rutherford Model
3.10 Concept of Energy Levels or Orbits
Ex. 3.19 to 3.23
P.P 3.28 to 3.32
Day 8
Bohr’s Theory and Concept of Quantisation
3.11 Modern Concept of Structure of an Atom: Quantum Mechanical Model
Ex. 3.24 to 3.29
P.P. 3.33, 3.34
Day 9
3.12 Wave Mechanical Model of Atom and Concept of Atomic Orbital
3.13 Quantum Numbers
3.14 Pauli’s Exclusion Principle
Day 10
Ex. 3.30 to 3.33
P.P. 3.35 to 3.37
3.15 Shapes of Orbitals or Boundary Surface Diagrams
Day 11
3.16 Energy Level Diagram for Electrons in an Atom
3.17 Electronic Configuration of Atoms
Day 12
P.P 3.38 to 3.46
Revision Period
Day 13
Conceptual Questions with Answers: 1 to 1835
Day 14
Conceptual Questions with Answers: 19 to 35
Day 15
Additional Numerical Problems for Practice: 8
Day 16
Very Short Answer questions 16
Day 17
Short Answer Questions 42
Day 18
Study Competition File
Day 19
Numerical Problems 1 to 10
Day 20
Numerical Problems 19 to 20
Day 21
Objective Questions: 1 to 24
Day 22
Objective Questions: 25 to 47
Day 23
Fill in the blanks: 10
Day 24
True or False: 10
Tampilkan postingan dengan label Atomic structure. Tampilkan semua postingan
Tampilkan postingan dengan label Atomic structure. Tampilkan semua postingan
Rabu, 11 Maret 2009
IIT JEE Chemistry-12th Portion - Study Guide - 1. Atomic Structure and Chemical Bonding
IIT JEE Chemistry 12th class portion starts from here.
Based on Book by Jauhar CBSE 12th class
Sections in the Chapter
1.1 Dual Nature of Radiation
1.2 Dual Nature of Matter - de-Broglie Equation
1.3 Heisengberg's Uncertainty Principle
1.4 Wave Mechanical Model of Atom and Concept of Atomic Orbital
1.5 Quantum Numbers
1.6 Pauli's Exclusion Principles
1.7 Orbital Wave Functions and Shapes of Orbitals
1.8 Electronic Configurations of Atoms
1.9 Chemical Bonding
1.10 REview of Valency Bond Theory
1.11 Molecular Orbital Theory
1.12 Linear Combination of Atomic Orbitals (LCAO) Method
1.13 Relative Energies of Bonding and Antibonding Molecular Orbitals
1.14 Combination of 2s and 2p Atomic Orbitals to form Molecular Orbitals
1.15 Conditions for the Combination of Atomic Orbitals
1.16 Energy Level diagram for Molecular Orbitals
1.17 Rules for Filling Molecualr Orbitals
1.18 Electronic Configurations and Molecular Behavior
1.19 Bonding in Some Diatomic Molecules
1.20 Metallic Bond
1.21 Hybridisation
1.22 Intermolecular Forces
1.23 Hydrogen Bonding
Study Plan
Day 1
1.1 Dual Nature of Radiation
Pratice Problems 1.1 to 1.12
Day 2
1.2 Dual Nature of Matter - de-Broglie Equation
P.P. 1.13 to 1.22
Day 3
1.3 Heisengberg's Uncertainty Principle
P.P. 1.23 to 1.29
Day 4
1.4 Wave Mechanical Model of Atom and Concept of Atomic Orbital
P.P. 130 to 1.31
Day 5
1.5 Quantum Numbers
1.6 Pauli's Exclusion Principles
P.P. 1.32 to 1.35
Day 6
1.7 Orbital Wave Functions and Shapes of Orbitals
1.8 Electronic Configurations of Atoms
Day 7
1.9 Chemical Bonding
1.10 REview of Valency Bond Theory
1.11 Molecular Orbital Theory
1.12 Linear Combination of Atomic Orbitals (LCAO) Method
1.13 Relative Energies of Bonding and Antibonding Molecular Orbitals
Day 8
1.14 Combination of 2s and 2p Atomic Orbitals to form Molecular Orbitals
1.15 Conditions for the Combination of Atomic Orbitals
P.P. 1.36 to 1.37
1.16 Energy Level diagram for Molecular Orbitals
1.17 Rules for Filling Molecualr Orbitals
1.18 Electronic Configurations and Molecular Behavior
Day 9
1.19 Bonding in Some Diatomic Molecules
P.P. 1.38 to 1.45
Day 10
1.20 Metallic Bond
Day 11
1.21 Hybridisation
1.22 Intermolecular Forces
Day 12
1.23 Hydrogen Bonding
P.P. 1.44 to 1.47
Day 13
Additional numerical problems 1 to 20
day 14
Conceptual questions 1 to 40
Day 15
Key facts to remember
Revision exercises very short answer questions 1 to 30
Revision Period (30 minutes per day)
Day 16
Revision exercises very short answer questions 31 to 50
day 17
Revision exercises very short answer questions 51 to 65
Day 18
Revision exercises short answer questions 1 to 15
Day 19
Revision exercises short answer questions 16 to 30
Day 20
Revision exercises short answer questions 31 to 45
Day 21
Revision exercises short answer questions 46 to 60
Day 22
Revision exercises short answer questions 61 to 78
Day 23
Some useful facts for competitive examinations
Nemerical problems for competitive examinations 1 to 5
Day 24
Nemerical problems for competitive examinations 6 to 13
Day 25
Multiple choice questions 1 to 20
Day 26
Multiple choice questions 21 to 40
Day 27
Multiple choice questions 41 to 63
Day 28
Fill in the blanks 1 to 15
Day 29
True or false 15
Day 30
Matching
Concept revision
Based on Book by Jauhar CBSE 12th class
Sections in the Chapter
1.1 Dual Nature of Radiation
1.2 Dual Nature of Matter - de-Broglie Equation
1.3 Heisengberg's Uncertainty Principle
1.4 Wave Mechanical Model of Atom and Concept of Atomic Orbital
1.5 Quantum Numbers
1.6 Pauli's Exclusion Principles
1.7 Orbital Wave Functions and Shapes of Orbitals
1.8 Electronic Configurations of Atoms
1.9 Chemical Bonding
1.10 REview of Valency Bond Theory
1.11 Molecular Orbital Theory
1.12 Linear Combination of Atomic Orbitals (LCAO) Method
1.13 Relative Energies of Bonding and Antibonding Molecular Orbitals
1.14 Combination of 2s and 2p Atomic Orbitals to form Molecular Orbitals
1.15 Conditions for the Combination of Atomic Orbitals
1.16 Energy Level diagram for Molecular Orbitals
1.17 Rules for Filling Molecualr Orbitals
1.18 Electronic Configurations and Molecular Behavior
1.19 Bonding in Some Diatomic Molecules
1.20 Metallic Bond
1.21 Hybridisation
1.22 Intermolecular Forces
1.23 Hydrogen Bonding
Study Plan
Day 1
1.1 Dual Nature of Radiation
Pratice Problems 1.1 to 1.12
Day 2
1.2 Dual Nature of Matter - de-Broglie Equation
P.P. 1.13 to 1.22
Day 3
1.3 Heisengberg's Uncertainty Principle
P.P. 1.23 to 1.29
Day 4
1.4 Wave Mechanical Model of Atom and Concept of Atomic Orbital
P.P. 130 to 1.31
Day 5
1.5 Quantum Numbers
1.6 Pauli's Exclusion Principles
P.P. 1.32 to 1.35
Day 6
1.7 Orbital Wave Functions and Shapes of Orbitals
1.8 Electronic Configurations of Atoms
Day 7
1.9 Chemical Bonding
1.10 REview of Valency Bond Theory
1.11 Molecular Orbital Theory
1.12 Linear Combination of Atomic Orbitals (LCAO) Method
1.13 Relative Energies of Bonding and Antibonding Molecular Orbitals
Day 8
1.14 Combination of 2s and 2p Atomic Orbitals to form Molecular Orbitals
1.15 Conditions for the Combination of Atomic Orbitals
P.P. 1.36 to 1.37
1.16 Energy Level diagram for Molecular Orbitals
1.17 Rules for Filling Molecualr Orbitals
1.18 Electronic Configurations and Molecular Behavior
Day 9
1.19 Bonding in Some Diatomic Molecules
P.P. 1.38 to 1.45
Day 10
1.20 Metallic Bond
Day 11
1.21 Hybridisation
1.22 Intermolecular Forces
Day 12
1.23 Hydrogen Bonding
P.P. 1.44 to 1.47
Day 13
Additional numerical problems 1 to 20
day 14
Conceptual questions 1 to 40
Day 15
Key facts to remember
Revision exercises very short answer questions 1 to 30
Revision Period (30 minutes per day)
Day 16
Revision exercises very short answer questions 31 to 50
day 17
Revision exercises very short answer questions 51 to 65
Day 18
Revision exercises short answer questions 1 to 15
Day 19
Revision exercises short answer questions 16 to 30
Day 20
Revision exercises short answer questions 31 to 45
Day 21
Revision exercises short answer questions 46 to 60
Day 22
Revision exercises short answer questions 61 to 78
Day 23
Some useful facts for competitive examinations
Nemerical problems for competitive examinations 1 to 5
Day 24
Nemerical problems for competitive examinations 6 to 13
Day 25
Multiple choice questions 1 to 20
Day 26
Multiple choice questions 21 to 40
Day 27
Multiple choice questions 41 to 63
Day 28
Fill in the blanks 1 to 15
Day 29
True or false 15
Day 30
Matching
Concept revision
Minggu, 28 Desember 2008
Atomic structure and chemical bonding- Study Guide - IIT JEE
Bohr model, spectrum of hydrogen atom, quantum numbers; Wave-particle duality, de Broglie hypothesis; Uncertainty principle; Quantum mechanical picture of hydrogen atom (qualitative treatment), shapes of s, p and d orbitals; Electronic configurations of elements (up to atomic number 36); Aufbau principle; Pauli's exclusion principle and Hund's rule;
Orbital overlap and covalent bond; Hybridisation involving s, p and d orbitals only; Orbital energy diagrams for homonuclear diatomic species; Hydrogen bond; Polarity in molecules, dipole moment (qualitative aspects only); VSEPR model and shapes of molecules (linear, angular, triangular, square planar, pyramidal, square pyramidal, trigonal bipyramidal, tetrahedral and octahedral).
Orbital overlap and covalent bond; Hybridisation involving s, p and d orbitals only; Orbital energy diagrams for homonuclear diatomic species; Hydrogen bond; Polarity in molecules, dipole moment (qualitative aspects only); VSEPR model and shapes of molecules (linear, angular, triangular, square planar, pyramidal, square pyramidal, trigonal bipyramidal, tetrahedral and octahedral).
Senin, 04 Februari 2008
IIT JEE Revision - Ch.3. ATOMIC STRUCTURE Core Points
JEE Syllabus
Atomic structure
Bohr model, spectrum of hydrogen atom, quantum numbers;
Wave-particle duality, de Broglie hypothesis;
Uncertainty principle;
Quantum mechanical picture of hydrogen atom (qualitative treatment), shapes of s, p and d orbitals;
Electronic configurations of elements (up to atomic number 36);
Aufbau principle;
Pauli's exclusion principle and Hund's rule;
In 1913, Niels Bohr proposed a model of the atom. He proposed that the electrons in an atom could only be in certain orbits, or energy levels, around the nucleus. Refinement of Bohr theory led to the modern theory of atomic structure based on quantum mechanics.
Bohr's model is based on particle theory.
As wave-particle duality which says that all micromatter particle exhibit dualitya new model is proposed.
Debroglie p = h/λ
Uncertainty principle Δp* Δx ≥h/4π
The quantum mechanics model
Principal quantum number-Shell,Azimuthal quantum number-sublevel,Magnetic quantum number-orbital, spin quantum number
The orbits are called as shells. The energy level of orbits or shells increases as they increase in distance from the nucleus of the atom. The orbits or shells are represented by numbers as 1,2,3,4,5,6 or 7. They are represented by letters as K,L,M,N,O,P,Q.
Sublevel of an Orbit
The energy levels, or orbits or shells are further divided into sublevels, or subshells. These subshells are designated by letters: s for the first possible sublevel, p for the second possible sublevel, d for the third, f for the fourth, g for the fifth, and from here on they simply go in alphabets.
The number of sublevels of each energy level is equal to the number of the energy levels. This means energy level 1, the K shell will have only one sub levels – s sublevel. The energy level 2, the L shell will have 2 sub levels – s and p.
Orbitals
Sublevels have further divisions called orbitals. Electrons are found in these orbitals. Each orbital contains two electrons.
“s” sublevel has only one orbital. “p” sublevel has 3 orbitals. “d” sublevel has 5 orbitals. “f” sublevel has 7 orbitals.
The two electons in each orbital spin in different directions.
Shape of Orbitals
1. Spherical shape for s.
2. Dumbbell shape for orbitals of p.
3. Four-lobed shape for orbitals of d.
4. Complex shape for all orbitals of higher sublevels.
Pauli's exclusion principle: No two electrons can have all four same quantum numbers
Electrons occupy the lowest energy sublevels that are available. This is known as ‘aufbau’ order or principles.
Hund’s rule says that, for any set of orbitals of equal energy say p orbitals of orbit 2, there is one electron is each orbital before the second electron enters or occupies an orbital.
The energy level of some sublevels at higher orbits is less than the some sublevels at lower orbitals. This fact is to be kept in mind when electron configuration is determined for any atom. The increasing order of energy levels of sublevels is:
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f*, 5d, 6p, 7s, 5f*, 6d, 7p, 8s
Atomic structure
Bohr model, spectrum of hydrogen atom, quantum numbers;
Wave-particle duality, de Broglie hypothesis;
Uncertainty principle;
Quantum mechanical picture of hydrogen atom (qualitative treatment), shapes of s, p and d orbitals;
Electronic configurations of elements (up to atomic number 36);
Aufbau principle;
Pauli's exclusion principle and Hund's rule;
In 1913, Niels Bohr proposed a model of the atom. He proposed that the electrons in an atom could only be in certain orbits, or energy levels, around the nucleus. Refinement of Bohr theory led to the modern theory of atomic structure based on quantum mechanics.
Bohr's model is based on particle theory.
As wave-particle duality which says that all micromatter particle exhibit dualitya new model is proposed.
Debroglie p = h/λ
Uncertainty principle Δp* Δx ≥h/4π
The quantum mechanics model
Principal quantum number-Shell,Azimuthal quantum number-sublevel,Magnetic quantum number-orbital, spin quantum number
The orbits are called as shells. The energy level of orbits or shells increases as they increase in distance from the nucleus of the atom. The orbits or shells are represented by numbers as 1,2,3,4,5,6 or 7. They are represented by letters as K,L,M,N,O,P,Q.
Sublevel of an Orbit
The energy levels, or orbits or shells are further divided into sublevels, or subshells. These subshells are designated by letters: s for the first possible sublevel, p for the second possible sublevel, d for the third, f for the fourth, g for the fifth, and from here on they simply go in alphabets.
The number of sublevels of each energy level is equal to the number of the energy levels. This means energy level 1, the K shell will have only one sub levels – s sublevel. The energy level 2, the L shell will have 2 sub levels – s and p.
Orbitals
Sublevels have further divisions called orbitals. Electrons are found in these orbitals. Each orbital contains two electrons.
“s” sublevel has only one orbital. “p” sublevel has 3 orbitals. “d” sublevel has 5 orbitals. “f” sublevel has 7 orbitals.
The two electons in each orbital spin in different directions.
Shape of Orbitals
1. Spherical shape for s.
2. Dumbbell shape for orbitals of p.
3. Four-lobed shape for orbitals of d.
4. Complex shape for all orbitals of higher sublevels.
Pauli's exclusion principle: No two electrons can have all four same quantum numbers
Electrons occupy the lowest energy sublevels that are available. This is known as ‘aufbau’ order or principles.
Hund’s rule says that, for any set of orbitals of equal energy say p orbitals of orbit 2, there is one electron is each orbital before the second electron enters or occupies an orbital.
The energy level of some sublevels at higher orbits is less than the some sublevels at lower orbitals. This fact is to be kept in mind when electron configuration is determined for any atom. The increasing order of energy levels of sublevels is:
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f*, 5d, 6p, 7s, 5f*, 6d, 7p, 8s
JEE ervision - Atomic structure - Basics
Thomson Model
Rutherford model
In 1913, Niels Bohr proposed a model of the atom. He proposed that the electrons in an atom could only be in certain orbits, or energy levels, around the nucleus. Refinement of Bohr theory led to the modern theory of atomic structure based on quantum mechanics.
Wave mechanical model
Rutherford model
In 1913, Niels Bohr proposed a model of the atom. He proposed that the electrons in an atom could only be in certain orbits, or energy levels, around the nucleus. Refinement of Bohr theory led to the modern theory of atomic structure based on quantum mechanics.
Wave mechanical model
JEE Revision - Bohr model
The most important properties of atomic and molecular structure may be exemplified using a simplified picture of an atom that is called the Bohr Model.
This model was proposed by Niels Bohr in 1915;
It was revised later on as it is not completely correct, but it has many features that are approximately correct and they were used in the later models.
The correct theory of the atom is called quantum mechanics;
The Bohr Model is an approximation to quantum mechanics that has the virtue of being much simpler.
A Planetary Model of the Atom
The Bohr Model is probably familar as the "planetary model" of the atom. In the Bohr Model the neutrons and protons occupy a dense central region called the nucleus, and the electrons orbit the nucleus much like planets orbiting the Sun (but the orbits are not confined to a plane).
This similarity between a planetary model and the Bohr Model of the atom ultimately arises because the attractive gravitational force in a solar system and the attractive Coulomb (electrical) force between the positively charged nucleus and the negatively charged electrons in an atom are mathematically of the same form. (The form is the same, but the intrinsic strength of the Coulomb interaction is much larger than that of the gravitational interaction; in addition, there are positive and negative electrical charges so the Coulomb interaction can be either attractive or repulsive, but gravitation is always attractive in our present Universe.)
Orbits Are Quantized
Quantized energy levels in hydrogen
The basic feature of quantum mechanics that is incorporated in the Bohr Model and that is completely different from the analogous planetary model is that the energy of the particles in the Bohr atom is restricted to certain discrete values.
One says that the energy is quantized. This means that only certain orbits with certain radii are allowed; orbits in between simply don't exist.
These levels are labeled by an integer n that is called a quantum number.
The lowest energy state is generally termed the ground state. The states with successively more energy than the ground state are called the first excited state, the second excited state, and so on.
Beyond an energy called the ionization potential the single electron of the hydrogen atom is no longer bound to the atom. Then the energy levels form a continuum.
In the case of hydrogen, this continuum starts at 13.6 eV above the ground state ("eV" stands for "electron-Volt", a common unit of energy in atomic physics).
This model was proposed by Niels Bohr in 1915;
It was revised later on as it is not completely correct, but it has many features that are approximately correct and they were used in the later models.
The correct theory of the atom is called quantum mechanics;
The Bohr Model is an approximation to quantum mechanics that has the virtue of being much simpler.
A Planetary Model of the Atom
The Bohr Model is probably familar as the "planetary model" of the atom. In the Bohr Model the neutrons and protons occupy a dense central region called the nucleus, and the electrons orbit the nucleus much like planets orbiting the Sun (but the orbits are not confined to a plane).
This similarity between a planetary model and the Bohr Model of the atom ultimately arises because the attractive gravitational force in a solar system and the attractive Coulomb (electrical) force between the positively charged nucleus and the negatively charged electrons in an atom are mathematically of the same form. (The form is the same, but the intrinsic strength of the Coulomb interaction is much larger than that of the gravitational interaction; in addition, there are positive and negative electrical charges so the Coulomb interaction can be either attractive or repulsive, but gravitation is always attractive in our present Universe.)
Orbits Are Quantized
Quantized energy levels in hydrogen
The basic feature of quantum mechanics that is incorporated in the Bohr Model and that is completely different from the analogous planetary model is that the energy of the particles in the Bohr atom is restricted to certain discrete values.
One says that the energy is quantized. This means that only certain orbits with certain radii are allowed; orbits in between simply don't exist.
These levels are labeled by an integer n that is called a quantum number.
The lowest energy state is generally termed the ground state. The states with successively more energy than the ground state are called the first excited state, the second excited state, and so on.
Beyond an energy called the ionization potential the single electron of the hydrogen atom is no longer bound to the atom. Then the energy levels form a continuum.
In the case of hydrogen, this continuum starts at 13.6 eV above the ground state ("eV" stands for "electron-Volt", a common unit of energy in atomic physics).
Spectrum of Hydrogen Atom
All electromagnetic waves have the same speed (3.0*10^8 m s^-1).
But different types of radiations differ from one another in their wavelengths and therefore, in frequency.
v = c/λ
C = speed = constant
λ = wavelength
v = frequency
The arrangement of different types of electromagnetic radiations in the order of increasing wavelengths (or decreasing frequencies) is known as electromagnetic spectrum.
When white light from sun is passed through a prism, it splits into a series of colour bands know as rainbow of colours: violet, indigo, blue, green, yellow, orange and red (VIBGYOR). This series of colour bands obtained is called a spectrum and in this spectrum of visible light there is continuity of colours, one colour merges into the other without any gap.
We can obtain spectrum of atoms. When the gases of vapour of a chemical substance are heated by electric spark, light is emitted. From such light emitted, the spectrum can be obtained using spectroscope.
It is observed that when radiations emitted by different substances are analysed, the spectrum obtained consists of share well defined lines each corresponding to a definited frequency (or wave length). The spectrum obtained from light emitted by chemical substances is called emission spectrum.
When electromagnetic radiation for example say white light is allowed to pass through a gas or a solution of some salt and the light that comes out is sent through a spectroscope, we obtain a spectrum. In this spectrum some dark lines are observed, in an otherwise continuous spectrum. This indicates that some radiations of specific wavelengths are absorbed by the substance. This spectrum is called absorption spectrum
Spectrum of hydrogen atom
The spectrum of hydrogen atom can be obtained by passing an electric discharge through the hydrogen gas in a discharge tube under low pressure.
It is observed that the spectrum consists of a large number of lines appearing in different regions of wavelengths. Some of the lines are present in visible region while others in ultra-violet and infra-red regions.
In 1885, J.J. Balmer developed a simple relationship among the different wavelengths of the series of visible lines in the hydrogen spectrum.
The relationship developed by Balmer for frequencies observed in hydrogen spe trum is:
1/ λ = v (in cm^-1) = 109677 (1/2² - 1/n²)
Where n is an integer that takes values equal to or greater than 3.
109677 cm^-1 is called Rydberg constant.
The lines in the hydrogen spectrum in various regions of the electromagnetic spectrum are given different names.
Lyman series, Balmer series, paschen series, Brackett series, and Pfund series
Lyman series appears in the ultraviolet region.
Balmer series appear in the visible region.
Other three series appear in the infra red region.
A more general formula for hydrogen spectral lines is
1/ λ = v (in cm^-1) = 109677 (1/nf² - 1/ni²)
ni>nf
ni = the initial level of the electron
nf = the final level of the electron
The above equation is called Rydberg equation.
For lyman series nf = 1
For Balmer series nf = 2
For Paschen series nf = 3
For Brackett series nf = 4
For Pfund series nf = 5
We have to note that above equation is valid for hydrogen spectrum lines only.
Emission of radiation occurs from atoms when electrons in the atom goes into an excited state and then returns to a lower energy state.
For example, if an hydrogen atom goes into an excited state, say the electron goes to fourth energy level and then returns to second energy level, radiation is emitted. The frequency of the emitted radiation can be found out by
1/ λ = v (in cm^-1) = 109677 (1/2² - 1/4²) = 486 nm.
This frequency corresponds to bluish green in visible region.
But different types of radiations differ from one another in their wavelengths and therefore, in frequency.
v = c/λ
C = speed = constant
λ = wavelength
v = frequency
The arrangement of different types of electromagnetic radiations in the order of increasing wavelengths (or decreasing frequencies) is known as electromagnetic spectrum.
When white light from sun is passed through a prism, it splits into a series of colour bands know as rainbow of colours: violet, indigo, blue, green, yellow, orange and red (VIBGYOR). This series of colour bands obtained is called a spectrum and in this spectrum of visible light there is continuity of colours, one colour merges into the other without any gap.
We can obtain spectrum of atoms. When the gases of vapour of a chemical substance are heated by electric spark, light is emitted. From such light emitted, the spectrum can be obtained using spectroscope.
It is observed that when radiations emitted by different substances are analysed, the spectrum obtained consists of share well defined lines each corresponding to a definited frequency (or wave length). The spectrum obtained from light emitted by chemical substances is called emission spectrum.
When electromagnetic radiation for example say white light is allowed to pass through a gas or a solution of some salt and the light that comes out is sent through a spectroscope, we obtain a spectrum. In this spectrum some dark lines are observed, in an otherwise continuous spectrum. This indicates that some radiations of specific wavelengths are absorbed by the substance. This spectrum is called absorption spectrum
Spectrum of hydrogen atom
The spectrum of hydrogen atom can be obtained by passing an electric discharge through the hydrogen gas in a discharge tube under low pressure.
It is observed that the spectrum consists of a large number of lines appearing in different regions of wavelengths. Some of the lines are present in visible region while others in ultra-violet and infra-red regions.
In 1885, J.J. Balmer developed a simple relationship among the different wavelengths of the series of visible lines in the hydrogen spectrum.
The relationship developed by Balmer for frequencies observed in hydrogen spe trum is:
1/ λ = v (in cm^-1) = 109677 (1/2² - 1/n²)
Where n is an integer that takes values equal to or greater than 3.
109677 cm^-1 is called Rydberg constant.
The lines in the hydrogen spectrum in various regions of the electromagnetic spectrum are given different names.
Lyman series, Balmer series, paschen series, Brackett series, and Pfund series
Lyman series appears in the ultraviolet region.
Balmer series appear in the visible region.
Other three series appear in the infra red region.
A more general formula for hydrogen spectral lines is
1/ λ = v (in cm^-1) = 109677 (1/nf² - 1/ni²)
ni>nf
ni = the initial level of the electron
nf = the final level of the electron
The above equation is called Rydberg equation.
For lyman series nf = 1
For Balmer series nf = 2
For Paschen series nf = 3
For Brackett series nf = 4
For Pfund series nf = 5
We have to note that above equation is valid for hydrogen spectrum lines only.
Emission of radiation occurs from atoms when electrons in the atom goes into an excited state and then returns to a lower energy state.
For example, if an hydrogen atom goes into an excited state, say the electron goes to fourth energy level and then returns to second energy level, radiation is emitted. The frequency of the emitted radiation can be found out by
1/ λ = v (in cm^-1) = 109677 (1/2² - 1/4²) = 486 nm.
This frequency corresponds to bluish green in visible region.
quantum numbers
According to quantum mechanical model or wave mechanical model of atom, orbitals represent regions in space around the nucleus where the probability of finding electrons is maximum. A large number of orbitals are possible in an atom.
To describe each electron in an atom in different orbitals, four quantum numbers are used. They are designated as n,l,ml, and ms.
1. Principal quantum number (n) This quantum number determines the main energy shell or level in which the electron is present. It can have whole number values starting from 1 in an atom.
The principle quantum number indicates the average distance of the electron from the nucleus. If n = 1, it is closest to the nucleus and has lowest energy.
Eariest practice was to number shells as K,L,M,N etc.
Shell with principal quantum number n = 1 is called K.
Shell with principal quantum number n = 2 is called etc.
2. Azimuthal quantum number or angular quantum number (l): This number determines the angular momentum of the electron.
It can have positive integer values from zero to (n-1) where n is the principal quantum number. For each value of n, there are n possible values of l.
For n =3, l has three values: l = 0,1,2
The earlier practice is to designate l as subshell and refer it by letters s,p,d,f,….
l=0 = s; l=1=p; l=2=d, l=3=f etc.
The energy of subshell increases with increasing value of l.
3. Magnetic quantum number ( ml): Magnetic field acts on moving electrical charges. ( from chapters on magnetism in physics syllabus). On revolving electrons external magnetic field of the earth acts. Therefore, the electrons in a given subshell orient themselves in certain preferred regions space around the nucleus. These are called orbitals. This quantum number gives the number of orbitals for given angular quantum number l or in a given subshell.
The allowed values of ml are –l through 0 to +l.
There are (2l+1) values of ml for each value of l.
If l = 0, ml has only one value. ml = 0.
If l = 3, ml has 7 values.
ml = -3,-2,-1,0,1,2,3
4. Spin quantum number (ms) : It is observed that the electron in an atom is not only revolving around the nucleus but is also spinning around its own axis. This quantum number describes the spin orientation of the electron.
The electron can spin in two ways – clockwise and anticlockwise.
Values of +1/2 and -1/2 are given to this quantum number. Its value is not dependent on other quantum numbers.
The orientations of spin are also designated by up and down arrows ↑ ↓.
To describe each electron in an atom in different orbitals, four quantum numbers are used. They are designated as n,l,ml, and ms.
1. Principal quantum number (n) This quantum number determines the main energy shell or level in which the electron is present. It can have whole number values starting from 1 in an atom.
The principle quantum number indicates the average distance of the electron from the nucleus. If n = 1, it is closest to the nucleus and has lowest energy.
Eariest practice was to number shells as K,L,M,N etc.
Shell with principal quantum number n = 1 is called K.
Shell with principal quantum number n = 2 is called etc.
2. Azimuthal quantum number or angular quantum number (l): This number determines the angular momentum of the electron.
It can have positive integer values from zero to (n-1) where n is the principal quantum number. For each value of n, there are n possible values of l.
For n =3, l has three values: l = 0,1,2
The earlier practice is to designate l as subshell and refer it by letters s,p,d,f,….
l=0 = s; l=1=p; l=2=d, l=3=f etc.
The energy of subshell increases with increasing value of l.
3. Magnetic quantum number ( ml): Magnetic field acts on moving electrical charges. ( from chapters on magnetism in physics syllabus). On revolving electrons external magnetic field of the earth acts. Therefore, the electrons in a given subshell orient themselves in certain preferred regions space around the nucleus. These are called orbitals. This quantum number gives the number of orbitals for given angular quantum number l or in a given subshell.
The allowed values of ml are –l through 0 to +l.
There are (2l+1) values of ml for each value of l.
If l = 0, ml has only one value. ml = 0.
If l = 3, ml has 7 values.
ml = -3,-2,-1,0,1,2,3
4. Spin quantum number (ms) : It is observed that the electron in an atom is not only revolving around the nucleus but is also spinning around its own axis. This quantum number describes the spin orientation of the electron.
The electron can spin in two ways – clockwise and anticlockwise.
Values of +1/2 and -1/2 are given to this quantum number. Its value is not dependent on other quantum numbers.
The orientations of spin are also designated by up and down arrows ↑ ↓.
Wave-particle duality
The earliest view of light, due to Newton, regarded light as made up of particles (commonly termed as corpuscles of light). The particle nature of light explained some of the experimental facts such as reflection and refraction of light. However, using particle nature of light, phenomena of interference and diffraction could not be explained. (See relevant chapters in Physics text and link that material with this idea).
Huygens proposed wave like character of light. With the help of wave theory of light, Huygens explained the phenomena of interference and diffraction as well as reflection and refraction. Hence the corpuscular theory was discarded and wave theory was adopted.
But once again, phenomena such black body radiation and photoelectric effect needed explanation. Wave theory could not explain these phenomena.
Einstein has suggested that light can behave as a wave as well as a particle i.e., it has dual character.
Characteristics of electromagnetic radiations
(i) The consists of electric and magnetic fields that oscillate in the directions perpendicular to each other and perpendicular to the direction in which the wave is traveling
(ii) These electromagnetic radiations do not require any medium for propagation.
(iii) All waves are characterized in terms of their wavelengths, frequency, velocity and amplitude.
Velocity of a wave = wavelength * frequency
Huygens proposed wave like character of light. With the help of wave theory of light, Huygens explained the phenomena of interference and diffraction as well as reflection and refraction. Hence the corpuscular theory was discarded and wave theory was adopted.
But once again, phenomena such black body radiation and photoelectric effect needed explanation. Wave theory could not explain these phenomena.
Einstein has suggested that light can behave as a wave as well as a particle i.e., it has dual character.
Characteristics of electromagnetic radiations
(i) The consists of electric and magnetic fields that oscillate in the directions perpendicular to each other and perpendicular to the direction in which the wave is traveling
(ii) These electromagnetic radiations do not require any medium for propagation.
(iii) All waves are characterized in terms of their wavelengths, frequency, velocity and amplitude.
Velocity of a wave = wavelength * frequency
JEE Revision - de Broglie hypothesis
de Broglie hypothesis
Wave nature of light could not explain the photoelectric effect. Einstein has suggested that light can behave as a wave as well as a particle i.e., it has dual character. In 1924, de-Broglie extended the wave nature to matter. He suggested that all microscopic particles such as electrons, protons, atoms and molecules have dual character similar to light.
Thus according to de Broglie hypothesis, “all material particles in motion possess wave characteristics.”
The wavelength(λ), mass of the particle (m) and velocity (v) of the particle are related to each other by the relation
λ = h/mv
where
h = Planck’s constant = 6.63 *10^-34 kg m² s^-1
The equation was experimentally verified by Davisson and Germer.
The wave like chacter of electron helped in making electron microscope. The electron microscope utilizes the wave like behaviour of electrons.
The dual nature of matter is applicable to all material objects but it is significant for microscopic bodies only. For bigger bodies, wave lengths calculated using the above formula are so small that they have no significance.
Wave nature of light could not explain the photoelectric effect. Einstein has suggested that light can behave as a wave as well as a particle i.e., it has dual character. In 1924, de-Broglie extended the wave nature to matter. He suggested that all microscopic particles such as electrons, protons, atoms and molecules have dual character similar to light.
Thus according to de Broglie hypothesis, “all material particles in motion possess wave characteristics.”
The wavelength(λ), mass of the particle (m) and velocity (v) of the particle are related to each other by the relation
λ = h/mv
where
h = Planck’s constant = 6.63 *10^-34 kg m² s^-1
The equation was experimentally verified by Davisson and Germer.
The wave like chacter of electron helped in making electron microscope. The electron microscope utilizes the wave like behaviour of electrons.
The dual nature of matter is applicable to all material objects but it is significant for microscopic bodies only. For bigger bodies, wave lengths calculated using the above formula are so small that they have no significance.
JEE revision - Heisenberg Uncertainty principle
Uncertainty principle
In 1927, Heisenberg put forward a principle known as Heisenberg’s uncertainty principle.
According it, “it is not possible to measure simultaneously both the position and momentum (or velocity) of a microscopic particle, with absolute accuracy.”
Mathematically, this principle is expressed as:
∆x * ∆p = h/4 π
Where
∆x = uncertainty in position
∆p = uncertainty in momentum
The constancy of the product of uncertainties means that, if the position of the particle is known with more accuracy, there will be large uncertainty in momentum and vice versa.
This uncertainty arises, as all observations are made by impact of light, the microscopic objects suffer a change in position or velocity as a result of impact of light. So there is a disturbance in them due to the measurement.
The principle does not affect the measurement of large objects as in these cases impact of light does not created any appreciable change in their position or velocity.
In 1927, Heisenberg put forward a principle known as Heisenberg’s uncertainty principle.
According it, “it is not possible to measure simultaneously both the position and momentum (or velocity) of a microscopic particle, with absolute accuracy.”
Mathematically, this principle is expressed as:
∆x * ∆p = h/4 π
Where
∆x = uncertainty in position
∆p = uncertainty in momentum
The constancy of the product of uncertainties means that, if the position of the particle is known with more accuracy, there will be large uncertainty in momentum and vice versa.
This uncertainty arises, as all observations are made by impact of light, the microscopic objects suffer a change in position or velocity as a result of impact of light. So there is a disturbance in them due to the measurement.
The principle does not affect the measurement of large objects as in these cases impact of light does not created any appreciable change in their position or velocity.
JEE Revision -Shapes of s, p and d orbitals
Shape of Orbitals
1. Spherical shape for s.
2. Dumbbell shape for orbitals of p.
3. Four-lobed shape for orbitals of d.
4. Complex shape for all orbitals of higher sublevels.
1. Spherical shape for s.
2. Dumbbell shape for orbitals of p.
3. Four-lobed shape for orbitals of d.
4. Complex shape for all orbitals of higher sublevels.
JEE Revision - Pauli's exclusion principle
Pauli's exclusion principle: No two electrons can have all four same quantum numbers
JEE Revision - Hund's rule
Hund’s rule says that, for any set of orbitals of equal energy say p orbitals of orbit 2, there is one electron is each orbital before the second electron enters or occupies an orbital.
Sabtu, 19 Januari 2008
IIT JEE Ch.3. ATOMIC STRUCTURE Core Points for Revision
JEE Syllabus
Atomic structure
Bohr model, spectrum of hydrogen atom, quantum numbers;
Wave-particle duality, de Broglie hypothesis;
Uncertainty principle;
Quantum mechanical picture of hydrogen atom (qualitative treatment), shapes of s, p and d orbitals;
Electronic configurations of elements (up to atomic number 36);
Aufbau principle;
Pauli's exclusion principle and Hund's rule;
----------------
In 1913, Niels Bohr proposed a model of the atom. He proposed that the electrons in an atom could only be in certain orbits, or energy levels, around the nucleus. Refinement of Bohr theory led to the modern theory of atomic structure based on quantum mechanics.
Bohr's model is based on particle theory.
As wave-particle duality which says that all micromatter particle exhibit dualitya new model is proposed.
Debroglie p = h/λ
Uncertainty principle Δp* Δx ≥h/4π
The quantum mechanics model
Principal quantum number-Shell,Azimuthal quantum number-sublevel,Magnetic quantum number-orbital, spin quantum number
The orbits are called as shells. The energy level of orbits or shells increases as they increase in distance from the nucleus of the atom. The orbits or shells are represented by numbers as 1,2,3,4,5,6 or 7. They are represented by letters as K,L,M,N,O,P,Q.
Sublevel of an Orbit
The energy levels, or orbits or shells are further divided into sublevels, or subshells. These subshells are designated by letters: s for the first possible sublevel, p for the second possible sublevel, d for the third, f for the fourth, g for the fifth, and from here on they simply go in alphabets.
The number of sublevels of each energy level is equal to the number of the energy levels. This means energy level 1, the K shell will have only one sub levels – s sublevel. The energy level 2, the L shell will have 2 sub levels – s and p.
Orbitals
Sublevels have further divisions called orbitals. Electrons are found in these orbitals. Each orbital contains two electrons.
“s” sublevel has only one orbital. “p” sublevel has 3 orbitals. “d” sublevel has 5 orbitals. “f” sublevel has 7 orbitals.
The two electons in each orbital spin in different directions.
Shape of Orbitals
1. Spherical shape for s.
2. Dumbbell shape for orbitals of p.
3. Four-lobed shape for orbitals of d.
4. Complex shape for all orbitals of higher sublevels.
Pauli's exclusion principle: No two electrons can have all four same quantum numbers
Electrons occupy the lowest energy sublevels that are available. This is known as ‘aufbau’ order or principles.
Hund’s rule says that, for any set of orbitals of equal energy say p orbitals of orbit 2, there is one electron is each orbital before the second electron enters or occupies an orbital.
The energy level of some sublevels at higher orbits is less than the some sublevels at lower orbitals. This fact is to be kept in mind when electron configuration is determined for any atom. The increasing order of energy levels of sublevels is:
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f*, 5d, 6p, 7s, 5f*, 6d, 7p, 8s
Atomic structure
Bohr model, spectrum of hydrogen atom, quantum numbers;
Wave-particle duality, de Broglie hypothesis;
Uncertainty principle;
Quantum mechanical picture of hydrogen atom (qualitative treatment), shapes of s, p and d orbitals;
Electronic configurations of elements (up to atomic number 36);
Aufbau principle;
Pauli's exclusion principle and Hund's rule;
----------------
In 1913, Niels Bohr proposed a model of the atom. He proposed that the electrons in an atom could only be in certain orbits, or energy levels, around the nucleus. Refinement of Bohr theory led to the modern theory of atomic structure based on quantum mechanics.
Bohr's model is based on particle theory.
As wave-particle duality which says that all micromatter particle exhibit dualitya new model is proposed.
Debroglie p = h/λ
Uncertainty principle Δp* Δx ≥h/4π
The quantum mechanics model
Principal quantum number-Shell,Azimuthal quantum number-sublevel,Magnetic quantum number-orbital, spin quantum number
The orbits are called as shells. The energy level of orbits or shells increases as they increase in distance from the nucleus of the atom. The orbits or shells are represented by numbers as 1,2,3,4,5,6 or 7. They are represented by letters as K,L,M,N,O,P,Q.
Sublevel of an Orbit
The energy levels, or orbits or shells are further divided into sublevels, or subshells. These subshells are designated by letters: s for the first possible sublevel, p for the second possible sublevel, d for the third, f for the fourth, g for the fifth, and from here on they simply go in alphabets.
The number of sublevels of each energy level is equal to the number of the energy levels. This means energy level 1, the K shell will have only one sub levels – s sublevel. The energy level 2, the L shell will have 2 sub levels – s and p.
Orbitals
Sublevels have further divisions called orbitals. Electrons are found in these orbitals. Each orbital contains two electrons.
“s” sublevel has only one orbital. “p” sublevel has 3 orbitals. “d” sublevel has 5 orbitals. “f” sublevel has 7 orbitals.
The two electons in each orbital spin in different directions.
Shape of Orbitals
1. Spherical shape for s.
2. Dumbbell shape for orbitals of p.
3. Four-lobed shape for orbitals of d.
4. Complex shape for all orbitals of higher sublevels.
Pauli's exclusion principle: No two electrons can have all four same quantum numbers
Electrons occupy the lowest energy sublevels that are available. This is known as ‘aufbau’ order or principles.
Hund’s rule says that, for any set of orbitals of equal energy say p orbitals of orbit 2, there is one electron is each orbital before the second electron enters or occupies an orbital.
The energy level of some sublevels at higher orbits is less than the some sublevels at lower orbitals. This fact is to be kept in mind when electron configuration is determined for any atom. The increasing order of energy levels of sublevels is:
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f*, 5d, 6p, 7s, 5f*, 6d, 7p, 8s
Langganan:
Postingan (Atom)