The document discusses spectroscopic notations used to describe the quantum states of atoms and ions. It introduces the principal, azimuthal, magnetic, and spin quantum numbers that are used to quantitatively describe observed atomic transitions. The spectroscopic notation describes the atomic state using these quantum numbers, written as 2S+1LJ, where S, L, and J are the spin, orbital, and total angular momentum quantum numbers. Examples are given for the ground and excited states of helium.
Basic operating principle and instrumentation of photo-luminescence technique. Brief description about interpretation of a photo-luminescence spectrum. Applications, advantages and disadvantages of photo-luminescence.
Energy bands consisting of a large number of closely spaced energy levels exist in crystalline materials. The bands can be thought of as the collection of the individual energy levels of electrons surrounding each atom. The wavefunctions of the individual electrons, however, overlap with those of electrons confined to neighboring atoms. The Pauli exclusion principle does not allow the electron energy levels to be the same so that one obtains a set of closely spaced energy levels, forming an energy band. The energy band model is crucial to any detailed treatment of semiconductor devices. It provides the framework needed to understand the concept of an energy bandgap and that of conduction in an almost filled band as described by the empty states.
Basic operating principle and instrumentation of photo-luminescence technique. Brief description about interpretation of a photo-luminescence spectrum. Applications, advantages and disadvantages of photo-luminescence.
Energy bands consisting of a large number of closely spaced energy levels exist in crystalline materials. The bands can be thought of as the collection of the individual energy levels of electrons surrounding each atom. The wavefunctions of the individual electrons, however, overlap with those of electrons confined to neighboring atoms. The Pauli exclusion principle does not allow the electron energy levels to be the same so that one obtains a set of closely spaced energy levels, forming an energy band. The energy band model is crucial to any detailed treatment of semiconductor devices. It provides the framework needed to understand the concept of an energy bandgap and that of conduction in an almost filled band as described by the empty states.
NQR - DEFINITION - ELECTRIC FIELD GRADIENT - NUCLEAR QUADRUPOLE MOMENT - NUCLEAR QUADRUPOLE COUPLING CONSTANT - PRINCIPLE OF NQR - ENERGY OF INTERACTION - SELECTION RULE - FREQUENCY OF TRANSITION - APPLICATIONS
In 1916, Sommerfeld extended Bohr's atomic model with the assumption of elliptical electron paths to explain the fine splitting of the spectral lines in the hydrogen atom. It is known as the Bohr-Sommerfeld model.
For more information on this concept, kindly visit our blog article at;
https://jayamchemistrylearners.blogspot.com/2022/04/bohr-sommerfeld-model-chemistrylearners.html
Spectroscopy is the study of the interaction of electromagnetic radiation in all its forms with the matter. The interaction might give rise to electronic excitations, (e.g. UV), molecular vibrations (e.g. IR) or nuclear spin orientations (e.g. NMR). Thus Spectroscopy is the science of the interaction of energy, in the form of electromagnetic radiation (EMR), acoustic waves, or particle beams, with the matter.
Here in this article, the matter is studied in further detail.
Photoluminescence Spectroscopy for studying Electron-Hole pair recombination ...RunjhunDutta
Description of Photoluminescence Spectroscopy: Principle, Instrumentation & Application.
Three research papers have been summarized which lay stress on Photoluminescence Study for Electron-Hole Pair Recombination for characterizing the properties of semiconductors used in Photoelectrochemical Splitting of Water.
Quantum Numbers and Atomic Orbitals By solving t.pdfarasanlethers
Quantum Numbers and Atomic Orbitals By solving the Schrödinger equation (Hy
= Ey), we obtain a set of mathematical equations, called wave functions (y), which describe the
probability of finding electrons at certain energy levels within an atom. A wave function for an
electron in an atom is called an atomic orbital; this atomic orbital describes a region of space in
which there is a high probability of finding the electron. Energy changes within an atom are the
result of an electron changing from a wave pattern with one energy to a wave pattern with a
different energy (usually accompanied by the absorption or emission of a photon of light). Each
electron in an atom is described by four different quantum numbers. The first three (n, l, ml)
specify the particular orbital of interest, and the fourth (ms) specifies how many electrons can
occupy that orbital. Principal Quantum Number (n): n = 1, 2, 3, …, 8 Specifies the energy of
an electron and the size of the orbital (the distance from the nucleus of the peak in a radial
probability distribution plot). All orbitals that have the same value of n are said to be in the same
shell (level). For a hydrogen atom with n=1, the electron is in its ground state; if the electron is in
the n=2 orbital, it is in an excited state. The total number of orbitals for a given n value is n2.
Angular Momentum (Secondary, Azimunthal) Quantum Number (l): l = 0, ..., n-1. Specifies the
shape of an orbital with a particular principal quantum number. The secondary quantum number
divides the shells into smaller groups of orbitals called subshells (sublevels). Usually, a letter
code is used to identify l to avoid confusion with n: l 0 1 2 3 4 5 ... Letter s p d f g h ... The
subshell with n=2 and l=1 is the 2p subshell; if n=3 and l=0, it is the 3s subshell, and so on. The
value of l also has a slight effect on the energy of the subshell; the energy of the subshell
increases with l (s < p < d < f). Magnetic Quantum Number (ml): ml = -l, ..., 0, ..., +l. Specifies
the orientation in space of an orbital of a given energy (n) and shape (l). This number divides the
subshell into individual orbitals which hold the electrons; there are 2l+1 orbitals in each subshell.
Thus the s subshell has only one orbital, the p subshell has three orbitals, and so on. Spin
Quantum Number (ms): ms = +½ or -½. Specifies the orientation of the spin axis of an electron.
An electron can spin in only one of two directions (sometimes called up and down). The Pauli
exclusion principle (Wolfgang Pauli, Nobel Prize 1945) states that no two electrons in the same
atom can have identical values for all four of their quantum numbers. What this means is that no
more than two electrons can occupy the same orbital, and that two electrons in the same orbital
must have opposite spins. Because an electron spins, it creates a magnetic field, which can be
oriented in one of two directions. For two electrons in the same orbital, the spins must be
opposite to each oth.
NQR - DEFINITION - ELECTRIC FIELD GRADIENT - NUCLEAR QUADRUPOLE MOMENT - NUCLEAR QUADRUPOLE COUPLING CONSTANT - PRINCIPLE OF NQR - ENERGY OF INTERACTION - SELECTION RULE - FREQUENCY OF TRANSITION - APPLICATIONS
In 1916, Sommerfeld extended Bohr's atomic model with the assumption of elliptical electron paths to explain the fine splitting of the spectral lines in the hydrogen atom. It is known as the Bohr-Sommerfeld model.
For more information on this concept, kindly visit our blog article at;
https://jayamchemistrylearners.blogspot.com/2022/04/bohr-sommerfeld-model-chemistrylearners.html
Spectroscopy is the study of the interaction of electromagnetic radiation in all its forms with the matter. The interaction might give rise to electronic excitations, (e.g. UV), molecular vibrations (e.g. IR) or nuclear spin orientations (e.g. NMR). Thus Spectroscopy is the science of the interaction of energy, in the form of electromagnetic radiation (EMR), acoustic waves, or particle beams, with the matter.
Here in this article, the matter is studied in further detail.
Photoluminescence Spectroscopy for studying Electron-Hole pair recombination ...RunjhunDutta
Description of Photoluminescence Spectroscopy: Principle, Instrumentation & Application.
Three research papers have been summarized which lay stress on Photoluminescence Study for Electron-Hole Pair Recombination for characterizing the properties of semiconductors used in Photoelectrochemical Splitting of Water.
Quantum Numbers and Atomic Orbitals By solving t.pdfarasanlethers
Quantum Numbers and Atomic Orbitals By solving the Schrödinger equation (Hy
= Ey), we obtain a set of mathematical equations, called wave functions (y), which describe the
probability of finding electrons at certain energy levels within an atom. A wave function for an
electron in an atom is called an atomic orbital; this atomic orbital describes a region of space in
which there is a high probability of finding the electron. Energy changes within an atom are the
result of an electron changing from a wave pattern with one energy to a wave pattern with a
different energy (usually accompanied by the absorption or emission of a photon of light). Each
electron in an atom is described by four different quantum numbers. The first three (n, l, ml)
specify the particular orbital of interest, and the fourth (ms) specifies how many electrons can
occupy that orbital. Principal Quantum Number (n): n = 1, 2, 3, …, 8 Specifies the energy of
an electron and the size of the orbital (the distance from the nucleus of the peak in a radial
probability distribution plot). All orbitals that have the same value of n are said to be in the same
shell (level). For a hydrogen atom with n=1, the electron is in its ground state; if the electron is in
the n=2 orbital, it is in an excited state. The total number of orbitals for a given n value is n2.
Angular Momentum (Secondary, Azimunthal) Quantum Number (l): l = 0, ..., n-1. Specifies the
shape of an orbital with a particular principal quantum number. The secondary quantum number
divides the shells into smaller groups of orbitals called subshells (sublevels). Usually, a letter
code is used to identify l to avoid confusion with n: l 0 1 2 3 4 5 ... Letter s p d f g h ... The
subshell with n=2 and l=1 is the 2p subshell; if n=3 and l=0, it is the 3s subshell, and so on. The
value of l also has a slight effect on the energy of the subshell; the energy of the subshell
increases with l (s < p < d < f). Magnetic Quantum Number (ml): ml = -l, ..., 0, ..., +l. Specifies
the orientation in space of an orbital of a given energy (n) and shape (l). This number divides the
subshell into individual orbitals which hold the electrons; there are 2l+1 orbitals in each subshell.
Thus the s subshell has only one orbital, the p subshell has three orbitals, and so on. Spin
Quantum Number (ms): ms = +½ or -½. Specifies the orientation of the spin axis of an electron.
An electron can spin in only one of two directions (sometimes called up and down). The Pauli
exclusion principle (Wolfgang Pauli, Nobel Prize 1945) states that no two electrons in the same
atom can have identical values for all four of their quantum numbers. What this means is that no
more than two electrons can occupy the same orbital, and that two electrons in the same orbital
must have opposite spins. Because an electron spins, it creates a magnetic field, which can be
oriented in one of two directions. For two electrons in the same orbital, the spins must be
opposite to each oth.
95electrons in the same orbital have different rus values .docxfredharris32
95
electrons in the same orbital have different rus values (one is +Yz and another -%),they
are said to be paired.
Electron Configuration
The energy of an electron in a hydrogen (H) atom is determined solely by its principal
quantum number n. However for many-electron atoms the orbital energies depend on
both the principal quantum number n andthe angular momenfum quantum number /.
Thus the energy of the orbitals in a many-electron atom increases in the order: ls < 2s <
2p < 3s a 3p < 4s < 3d < 4p < 5s, and so on. This order is also the order of filling
electrons into the orbitals in a many-electron atom. The guiding principle in assigning
electrons to the orbitals in a many-electron atom contains a set of three ru1es called the
Aufbau principle:
1. Lower-energy orbitals f,rll before higher-energy orbitals.
2. An atomic orbital can contain only two electrons, which must have opposite spins.
(Pauli exclusion principle: no two electrons in an atom can have the same four
quantum numbers.)
3 . When electrons are assigne d to p, d, or f orbitals, each successive electron enters a
different orbital of the subshell, each electron having the same spin as the previous
one; this proceeds until the subshell is half-full, after which electrons pair in the
orbitals one by one. (Hund's rule: the most stable arrangement of electrons in the
subshell is that with the maximum number of unpaired electrons, all with the same
spin.)
Flame Test
The resultant lowest-energy electron configuration is called the ground-state
configrnation of the atom. The electrons in the atom's outermost shell are called valance
electrons. When the atom absorbs enough energy, one or more of the valance electrons
move to a higher energy orbital, and the atom is said to be in an excited state. The excited
states are generally short-lived and rapidly decay back to the ground state by releasing
radiant energy in the form of light. The energy and frequency of the light that is released
during the decay transition depend on the difference in energy between the ground state
and the excited state. The energy difference (AQ, the frequency (v), and the wavelength
(2) of the light during emission are related by the equation, LE : hv: hclTwhere h ts
Planck's constant and c is the speed of light. When the wavelengths of the light emitted
fall in the visible region (400-800 nm), colors willbe observed.
Atoms of certain elements emit light
u'hen the elements or their
compounds are heated in a gas flame.
The flame takes on a distinctive
color detemined by the particular
element (flame test). Each atom has
its characteristic emission lines,
therefore flame tests can be used to
detect certain elements in unknown
cornpounds.
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the energy of the electron increases, and the electron is farther away from the nucleus. A
coliection of orbitals with the same n is called an electr.
In this presentation you will be able to, describe how atomic orbitals arise from the Schrodinger's equation, relate orbital shapes to electron density distribution and interpret the information obtained from a four set of quantum numbers.
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Operation “Blue Star” is the only event in the history of Independent India where the state went into war with its own people. Even after about 40 years it is not clear if it was culmination of states anger over people of the region, a political game of power or start of dictatorial chapter in the democratic setup.
The people of Punjab felt alienated from main stream due to denial of their just demands during a long democratic struggle since independence. As it happen all over the word, it led to militant struggle with great loss of lives of military, police and civilian personnel. Killing of Indira Gandhi and massacre of innocent Sikhs in Delhi and other India cities was also associated with this movement.
Model Attribute Check Company Auto PropertyCeline George
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Students, digital devices and success - Andreas Schleicher - 27 May 2024..pptxEduSkills OECD
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The Roman Empire A Historical Colossus.pdfkaushalkr1407
The Roman Empire, a vast and enduring power, stands as one of history's most remarkable civilizations, leaving an indelible imprint on the world. It emerged from the Roman Republic, transitioning into an imperial powerhouse under the leadership of Augustus Caesar in 27 BCE. This transformation marked the beginning of an era defined by unprecedented territorial expansion, architectural marvels, and profound cultural influence.
The empire's roots lie in the city of Rome, founded, according to legend, by Romulus in 753 BCE. Over centuries, Rome evolved from a small settlement to a formidable republic, characterized by a complex political system with elected officials and checks on power. However, internal strife, class conflicts, and military ambitions paved the way for the end of the Republic. Julius Caesar’s dictatorship and subsequent assassination in 44 BCE created a power vacuum, leading to a civil war. Octavian, later Augustus, emerged victorious, heralding the Roman Empire’s birth.
Under Augustus, the empire experienced the Pax Romana, a 200-year period of relative peace and stability. Augustus reformed the military, established efficient administrative systems, and initiated grand construction projects. The empire's borders expanded, encompassing territories from Britain to Egypt and from Spain to the Euphrates. Roman legions, renowned for their discipline and engineering prowess, secured and maintained these vast territories, building roads, fortifications, and cities that facilitated control and integration.
The Roman Empire’s society was hierarchical, with a rigid class system. At the top were the patricians, wealthy elites who held significant political power. Below them were the plebeians, free citizens with limited political influence, and the vast numbers of slaves who formed the backbone of the economy. The family unit was central, governed by the paterfamilias, the male head who held absolute authority.
Culturally, the Romans were eclectic, absorbing and adapting elements from the civilizations they encountered, particularly the Greeks. Roman art, literature, and philosophy reflected this synthesis, creating a rich cultural tapestry. Latin, the Roman language, became the lingua franca of the Western world, influencing numerous modern languages.
Roman architecture and engineering achievements were monumental. They perfected the arch, vault, and dome, constructing enduring structures like the Colosseum, Pantheon, and aqueducts. These engineering marvels not only showcased Roman ingenuity but also served practical purposes, from public entertainment to water supply.
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Lecture 02.; spectroscopic notations by Dr. Salma Amir
1. Lecture No. 02
Couse title:
Atomic Spectroscopy
Topic: Spectroscopic Notations
Course instructor: Dr. Salma Amir
GFCW Peshawar
2. Spectroscopic Notations
Spectroscopists of the late 19th and early 20th century created a system of
spectroscopic notation to describe the observed line spectra. Quantum
numbers were invented to provide an quantitative description of observed
(and unobserved) transitions. These together provide a short-hand
description of the state of the electrons in an atom or ion.
The spectroscopic notion is a method of describing the quantum state of an
atom using the principal quantum number, the orbital quantum number and
the total angular momentum quantum number. The total angular momentum
can be determined by taking the sum of the orbital quantum number and the
spin quantum number.
3. Quantum Numbers
Four quantum numbers suffice to describe any electron in an atom.
These are:
n, the principal quantum number., n takes on integral values
1, 2, 3, ... .
l, the azimuthal quantum number., l takes on the integral values
0, 1, 2, ... , n-2, n-1.
m, the magnetic quantum number. m takes on the integral values
-l , -(l-1), ..., -1, 0, 1, ..., (l-1), l.
s, the spin quantum number. This describes the spin of the electron,
and is either +1/2 or -1/2.
4. Quantum Numbers for Atoms
As with electrons, 4 quantum numbers suffice to describe the electronic state
of an atom or ion.
L is the total orbital angular momentum. L corresponds to the term of the ion
(S terms have L=0, P terms have L=1, etc.). In the case of more than one
electron in the outer shell, the value of L takes on all possible values of L=Σ li
The quantum number S is the absolute value of the total electron spin, S= Σsi.
Each electron has a spin of +/- 1/2. S is integral for an even number of
electrons, and half integral for an odd number. S=0 for a closed shell.
J represents the total angular momentum of the atom of ion. It is the vector
sum of L and S. For a hydrogenic ion, L=0, S=1/2, and J=1/2. For more
complex atom, J takes on the values L+S…….L-S,
M, the Magnetic quantum number, takes on values of J, J-1, ..., 0, ..., -J-1, -J.
5. Spectroscopic notations
The atomic level is described as
n 2S+1LJ or
2S+1LJ
where S, n, and J are the quantum numbers defined earlier, and L is the term
(S,P,D,F,G, etc). 2S+1 is the multiplicity
The multiplicity of a term is given by the value of 2S+1. A term with S=0 is
a singlet term; S=1/2 is a doublet term; S=1 is a triplet term; S=3/2 is a
quartet term, etc.
6. Spectroscopic notation for Helium (He)
No. of electrons=2
Electrons in ground state
S= s1+s2= ½-½=0
2S+1= 1 (Singlet state)
L=0 for electron in s
orbital
J=L+S= 0+0=0
Spectroscopic notation=
2S+1LJ
1So
Electrons in excited state
(1s and 2s)
S= s1+s2= ½-½=0
S= s1+s2= ½+½=1
2S+1= 1 (Singlet state) and
2S+1= 3 (Triplet state)
L=0 for electron in s orbital
J=L+S= 0+0=0
J=L+S= 0+1=1
Spectroscopic notation=
2S+1LJ
1So ,
3S1
