This document provides an overview of some foundational concepts in inorganic chemistry, including:
1. It defines inorganic chemistry as the study of inorganic substances excluding organic compounds.
2. It discusses atomic structure, including the components of atoms like protons, neutrons, and electrons. Molecular structures and bonding are also introduced.
3. Key concepts in molecular structure and bonding are covered, including the octet rule, Lewis diagrams, VSEPR model for predicting molecular shapes, valence bond theory involving hybrid atomic orbitals, and molecular orbital theory involving delocalized molecular orbitals.
The polarity of a molecule is determined by its molecular structure and the distribution of electrons within that structure. Polarity arises from differences in electronegativity between the atoms in a molecule. Electronegativity is a measure of an atom's ability to attract and hold onto electrons. When two atoms with different electronegativities bond together, the electrons in the bond are not shared equally, leading to an uneven distribution of charge within the molecule.
Polar Molecules: When there is an uneven distribution of charge within a molecule due to differences in electronegativity, the molecule is said to be polar. This results in a separation of charges, with one end of the molecule having a partial positive charge (δ+) and the other end having a partial negative charge (δ-).
Nonpolar Molecules: Nonpolar molecules have an even distribution of charge, meaning there are no significant differences in electronegativity between the atoms. As a result, there is no separation of charges within the molecule.
Electronegativity: The electronegativity of an atom is determined by the periodic table, and elements with higher electronegativities tend to attract electrons more strongly. The electronegativity difference between atoms in a bond is a key factor in determining the molecule's polarity.
Symmetry: In some cases, a molecule may have polar bonds but still be nonpolar overall due to its molecular geometry. If the polar bonds are arranged symmetrically so that the dipole moments cancel each other out, the molecule is nonpolar.
Dipole Moment: The dipole moment of a molecule is a measure of its polarity. It is a vector quantity that points from the positive end (δ+) to the negative end (δ-) of the molecule. A larger dipole moment indicates a more polar molecule.
Examples:
Water (H2O) is a polar molecule because oxygen is more electronegative than hydrogen, creating a significant dipole moment.
Carbon tetrachloride (CCl4) is a nonpolar molecule even though it has polar C-Cl bonds because the tetrahedral arrangement of the chlorine atoms results in cancellation of the dipole moments.
Solubility and Intermolecular Interactions: The polarity of a molecule plays a crucial role in its interactions with other molecules. Polar molecules tend to be soluble in polar solvents, while nonpolar molecules are more soluble in nonpolar solvents. Additionally, polar-polar interactions (dipole-dipole interactions) and nonpolar-nonpolar interactions (Van der Waals forces) are significant in determining the physical properties of substances.
Understanding the polarity of molecules is important in various fields, including chemistry, biology, and materials science, as it helps explain and predict the behavior of substances in chemical reactions and physical processes.
The polarity of a molecule is determined by its molecular structure and the distribution of electrons within that structure. Polarity arises from differences in electronegativity between the atoms in a molecule. Electronegativity is a measure of an atom's ability to attract and hold onto electrons. When two atoms with different electronegativities bond together, the electrons in the bond are not shared equally, leading to an uneven distribution of charge within the molecule.
Polar Molecules: When there is an uneven distribution of charge within a molecule due to differences in electronegativity, the molecule is said to be polar. This results in a separation of charges, with one end of the molecule having a partial positive charge (δ+) and the other end having a partial negative charge (δ-).
Nonpolar Molecules: Nonpolar molecules have an even distribution of charge, meaning there are no significant differences in electronegativity between the atoms. As a result, there is no separation of charges within the molecule.
Electronegativity: The electronegativity of an atom is determined by the periodic table, and elements with higher electronegativities tend to attract electrons more strongly. The electronegativity difference between atoms in a bond is a key factor in determining the molecule's polarity.
Symmetry: In some cases, a molecule may have polar bonds but still be nonpolar overall due to its molecular geometry. If the polar bonds are arranged symmetrically so that the dipole moments cancel each other out, the molecule is nonpolar.
Dipole Moment: The dipole moment of a molecule is a measure of its polarity. It is a vector quantity that points from the positive end (δ+) to the negative end (δ-) of the molecule. A larger dipole moment indicates a more polar molecule.
Examples:
Water (H2O) is a polar molecule because oxygen is more electronegative than hydrogen, creating a significant dipole moment.
Carbon tetrachloride (CCl4) is a nonpolar molecule even though it has polar C-Cl bonds because the tetrahedral arrangement of the chlorine atoms results in cancellation of the dipole moments.
Solubility and Intermolecular Interactions: The polarity of a molecule plays a crucial role in its interactions with other molecules. Polar molecules tend to be soluble in polar solvents, while nonpolar molecules are more soluble in nonpolar solvents. Additionally, polar-polar interactions (dipole-dipole interactions) and nonpolar-nonpolar interactions (Van der Waals forces) are significant in determining the physical properties of substances.
Understanding the polarity of molecules is important in various fields, including chemistry, biology, and materials science, as it helps explain and predict the behavior of substances in chemical reactions and physical processes.
ELEMENTARY PARTICLES OF MATTER
Matter is made up of discrete particles, the main ones are:-
1. Atoms 2. Ions 3. Molecules - (AIM)
DEFINITION OF ATOMS
An atom is the smallest particle of an element which can take part in a chemical reaction
THE CONSTITUENTS OF ATOMS
Rutherford in 1911 threw more light on the nature of the atom. He demonstrated that atom is made up of sub-particles which are called:
1. Proton 2. Neutron 3. Electron.
He discovered that the protons and neutrons are concentrated in the nucleus of an atom, while the electrons are revolving round the nucleus.
J.J THOMPSON’S MODEL
J.J Thompson described the atom as being made up of a mixture of positive (Protons) and negative (Electrons) charges.
LORD RUTHERFORD’S MODEL
Lord Rutherford described the atom as being made up of Positive (Protons) and Neutral (Neutrons) charges in its centre (nucleus) while the negative charges (electrons) rotates around its orbit.
He used the planetary bodies rotating around the sun to describe the structure of atoms.
CHARACTERISTICS OF PROTON, ELECTRON AND NEUTRON
DALTON’S ATOMIC THEORY
In 1808, John Dalton proposed the Atomic Theory which can be summarised as follows:
All elements are made up of small, indivisible particles called atoms.
Atoms can neither be created nor destroyed.
Atoms of the same elements are alike in every aspect, and differ from atoms of all other elements.
When atoms combine with other atoms, they do so in simple ratios.
All chemical changes result from the combination or the separation of atoms.
The Atomic Theory was partially supported by experimental evidences deduced from the Law of Conservation of Mass, the Law of Definite Proportions, the Law of Multiple Proportions and so on. It could not explain electrolysis and certain other phenomena. As a result of new discoveries, Dalton’s original Atomic Theory has undergone several modifications but the principal aspects as outlined above are still useful in the study of chemistry.
MODIFICATIONS OF DALTON’S ATOMIC THEORY
All Elements Are Made Up Of Small Indivisible Particles Called Atoms: This statement has been proven wrong by Rutherford’s discovery – the atom is built up of three main types of sub-particles: the proton, the electron, the neutron. It is not an indivisible solid piece.
The Atom Can Neither Be Created Nor Destroyed: This statement still holds good for ordinary chemical reactions and is embodied in the basic Law of Conservation of Mass. During a nuclear reaction, such as the fission of Uranium – 235, the nucleus is broken up into smaller units which form simpler atoms while a tremendous amount of heat energy is released. These changes that occur during nuclear fission destroy the atoms of the element involved.
The Atoms Of The Same Elements Are Alike In Every Aspect And Differ From Atoms Of All Other Elements: The discovery of isotopes makes this statement unacceptable. Chlorine, for example has two different atom
Professional air quality monitoring systems provide immediate, on-site data for analysis, compliance, and decision-making.
Monitor common gases, weather parameters, particulates.
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ELEMENTARY PARTICLES OF MATTER
Matter is made up of discrete particles, the main ones are:-
1. Atoms 2. Ions 3. Molecules - (AIM)
DEFINITION OF ATOMS
An atom is the smallest particle of an element which can take part in a chemical reaction
THE CONSTITUENTS OF ATOMS
Rutherford in 1911 threw more light on the nature of the atom. He demonstrated that atom is made up of sub-particles which are called:
1. Proton 2. Neutron 3. Electron.
He discovered that the protons and neutrons are concentrated in the nucleus of an atom, while the electrons are revolving round the nucleus.
J.J THOMPSON’S MODEL
J.J Thompson described the atom as being made up of a mixture of positive (Protons) and negative (Electrons) charges.
LORD RUTHERFORD’S MODEL
Lord Rutherford described the atom as being made up of Positive (Protons) and Neutral (Neutrons) charges in its centre (nucleus) while the negative charges (electrons) rotates around its orbit.
He used the planetary bodies rotating around the sun to describe the structure of atoms.
CHARACTERISTICS OF PROTON, ELECTRON AND NEUTRON
DALTON’S ATOMIC THEORY
In 1808, John Dalton proposed the Atomic Theory which can be summarised as follows:
All elements are made up of small, indivisible particles called atoms.
Atoms can neither be created nor destroyed.
Atoms of the same elements are alike in every aspect, and differ from atoms of all other elements.
When atoms combine with other atoms, they do so in simple ratios.
All chemical changes result from the combination or the separation of atoms.
The Atomic Theory was partially supported by experimental evidences deduced from the Law of Conservation of Mass, the Law of Definite Proportions, the Law of Multiple Proportions and so on. It could not explain electrolysis and certain other phenomena. As a result of new discoveries, Dalton’s original Atomic Theory has undergone several modifications but the principal aspects as outlined above are still useful in the study of chemistry.
MODIFICATIONS OF DALTON’S ATOMIC THEORY
All Elements Are Made Up Of Small Indivisible Particles Called Atoms: This statement has been proven wrong by Rutherford’s discovery – the atom is built up of three main types of sub-particles: the proton, the electron, the neutron. It is not an indivisible solid piece.
The Atom Can Neither Be Created Nor Destroyed: This statement still holds good for ordinary chemical reactions and is embodied in the basic Law of Conservation of Mass. During a nuclear reaction, such as the fission of Uranium – 235, the nucleus is broken up into smaller units which form simpler atoms while a tremendous amount of heat energy is released. These changes that occur during nuclear fission destroy the atoms of the element involved.
The Atoms Of The Same Elements Are Alike In Every Aspect And Differ From Atoms Of All Other Elements: The discovery of isotopes makes this statement unacceptable. Chlorine, for example has two different atom
Professional air quality monitoring systems provide immediate, on-site data for analysis, compliance, and decision-making.
Monitor common gases, weather parameters, particulates.
This pdf is about the Schizophrenia.
For more details visit on YouTube; @SELF-EXPLANATORY;
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Thanks...!
The increased availability of biomedical data, particularly in the public domain, offers the opportunity to better understand human health and to develop effective therapeutics for a wide range of unmet medical needs. However, data scientists remain stymied by the fact that data remain hard to find and to productively reuse because data and their metadata i) are wholly inaccessible, ii) are in non-standard or incompatible representations, iii) do not conform to community standards, and iv) have unclear or highly restricted terms and conditions that preclude legitimate reuse. These limitations require a rethink on data can be made machine and AI-ready - the key motivation behind the FAIR Guiding Principles. Concurrently, while recent efforts have explored the use of deep learning to fuse disparate data into predictive models for a wide range of biomedical applications, these models often fail even when the correct answer is already known, and fail to explain individual predictions in terms that data scientists can appreciate. These limitations suggest that new methods to produce practical artificial intelligence are still needed.
In this talk, I will discuss our work in (1) building an integrative knowledge infrastructure to prepare FAIR and "AI-ready" data and services along with (2) neurosymbolic AI methods to improve the quality of predictions and to generate plausible explanations. Attention is given to standards, platforms, and methods to wrangle knowledge into simple, but effective semantic and latent representations, and to make these available into standards-compliant and discoverable interfaces that can be used in model building, validation, and explanation. Our work, and those of others in the field, creates a baseline for building trustworthy and easy to deploy AI models in biomedicine.
Bio
Dr. Michel Dumontier is the Distinguished Professor of Data Science at Maastricht University, founder and executive director of the Institute of Data Science, and co-founder of the FAIR (Findable, Accessible, Interoperable and Reusable) data principles. His research explores socio-technological approaches for responsible discovery science, which includes collaborative multi-modal knowledge graphs, privacy-preserving distributed data mining, and AI methods for drug discovery and personalized medicine. His work is supported through the Dutch National Research Agenda, the Netherlands Organisation for Scientific Research, Horizon Europe, the European Open Science Cloud, the US National Institutes of Health, and a Marie-Curie Innovative Training Network. He is the editor-in-chief for the journal Data Science and is internationally recognized for his contributions in bioinformatics, biomedical informatics, and semantic technologies including ontologies and linked data.
Richard's entangled aventures in wonderlandRichard Gill
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Seminar of U.V. Spectroscopy by SAMIR PANDASAMIR PANDA
Spectroscopy is a branch of science dealing the study of interaction of electromagnetic radiation with matter.
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Earliest Galaxies in the JADES Origins Field: Luminosity Function and Cosmic ...Sérgio Sacani
We characterize the earliest galaxy population in the JADES Origins Field (JOF), the deepest
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6. Element & atom
• Chemical element: substance that cannot be broken down
into simpler substances & made of only one type of atom
• Chemical element has its own chemical symbol
• Atom: smallest constituent unit of ordinary matter that has
properties of a chemical element.1
6
1 ‘Atom’. Compendium of Chemical Terminology (IUPAC Gold Book) (2nd ed.). IUPAC. Retrieved 2015-04-25
8. Atomic structure
• Atom: nucleus (center) & electron
shells (outside)
• Nucleus: protons & neutrons
• Electron shells consist of electrons
• Most of atom is empty space !!!
8
nucleus
electron
shells
proton neutron
electron
Particle Symbol Mass /kg Relative
mass
Charge /C Relative
charge
Proton p 1.673 × 10–27 1 +1.602 × 10–19 +1
Neutron n 1.675 × 10–27 1 0 0
Electron e 9.109 × 10–31 1/1836 (~0) –1.602 × 10–19 –1
• Most of mass of atom concentrated in nucleus !!!
9. Atomic & mass number
• Proton/atomic number (Z): number of protons
• Each element has a unique Z
• Nucleon/mass number (A): number of protons & neutrons
• Isotopes: atoms with same Z & different A
• Isotopes have same chemical properties
9
10. • Relative atomic mass ( ): weighted mean of mass
numbers of all naturally occurring isotopes
• Example:
Mg: 78.99% 24Mg 10.00% 25Mg 11.01% 26Mg
masses 23.99 24.99 25.98
Isotopes
10
11. Isotopes
11
• Determination of accurate Ar from mass spectrometry
Mass spectrometer
Mass of each isotope
(mass/charge ratio)
How much of each
isotope (relative
abundance)
Isotopic
mass
Abundance
/%
20 90.9
21 0.3
22 8.8
Mass spectrum
12. Electronic structures of atoms
• Schrödinger equation:
• Solution of the equation provides:
Wavefunctions → atomic orbitals (AOs)
Energy E associated with particular
12
Why ?
13. AOs & Quantum Numbers
• AOs defined by 3 quantum numbers:
Principal quantum number ( ): related to size & energy AO
Orbital quantum number ( ): determines shape of orbital
Magnetic quantum number ( ): related to orientation of AO
13
principal
orbital
magnetic
15. Ground State Electronic Configuration
• aufbau (building-up) principle: lowest energy AOs filled first
• Pauli exclusion principle: maximum of 2 electrons in an AO
• Hund’s rule: maximum total spin
Example: ground state electronic
configuration of N (Z = 7), Cu (Z = 29)
15
16. Classic Periodic Table
16
• Most successful classification by Dmitri Mendeleev in 1869
Dmitri Ivanovich Mendeleev
(1834 –1907)
17. Classic Periodic Table
17
• Arrangement of elements in order of atomic weight
• Elements with similar chemical properties in same group
Annalen der Chemie und Pharmacie, VIII, Supplementary Volume for 1872, page 511
18. Modern Periodic Table
18
• Arrangement of elements in order of atomic number
row ~ period
column ~ group
19. The Periodic Table & Electronic Configuration
19
• Period: elements with same electron shells
• Period number: number of electron shells
• Group: elements with same valence electronic
configuration
• Group number: number of valence electrons with ‘1–18’ or
‘I – VIII’ numbering system
21. Basic concepts
• A molecule is an electrically neutral group of two or more
atoms held together by chemical bonds.1
• Classification of chemical bonds:
Ionic bonds: electron transfer between atoms to form ions
Covalent bonds: electron sharing between atoms
Metallic bonds
21
1 ‘Molecule’. Compendium of Chemical Terminology (IUPAC Gold Book) (2nd ed.). IUPAC. Retrieved 2015-04-25
22. Molecular structure & Bonding
1. The Octet Rule
2. Lewis Diagram
3. Molecular Shapes & the VSEPR Model
4. Valence Bond (VB) theory
5. Molecular Orbital (MO) theory
22
23. The octet rule
• Atoms gain, lose or share electrons to give an outer shell
with 8 electrons (an octet)
• Elements in Period 2 strictly obey the octet rule
23
If atoms is not only 8 electron in outer
shell ?
24. The octet rule
• Example: MgO
Mg (Z = 12) 1s2 2s22p6 3s2
Mg – 2e → Mg2+ (1s2 2s22p6)
O (Z = 8) 1s2 2s22p4
O + 2e → O2– (1s2 2s22p6)
→ MgO ionic compound
24
O
x
x
x
x
x
x
x Mg
x
x
x +
x
x
O2–
x
x
x
x
x
x
Mg2+
x
x
x
x
x
x
25. The octet rule
• Example: O2
O (Z = 8) 1s2 2s22p4
→ O2 covalent compound
25
O
x
x
x
x
x
x
O
+ O
x
x
x
x
x
x O
26. Molecular structure & Bonding
1. The Octet Rule
2. Lewis Diagram
3. Molecular Shapes & the VSEPR Model
4. Valence Bond (VB) theory
5. Molecular Orbital (MO) theory
26
27. Lewis Electron-dot Diagrams
• Using dots to describe number & arrangement of valence
electrons in molecules
• Example:
• Bonding pairs: pairs of electrons involved in bonding &
represented by a single line (–)
• Non-bonding electrons: electrons not involved in bonding
Atoms Molecule
28. Lewis Electron-dot Diagrams
• Constructing Lewis structure:
1. Determine overal number of valence electrons
2. Write arrangement of atoms bonded together
3. Distribute electrons in pairs so that each atom has an octet
• Example: CO2, SCN–
28
29. Resonance structures
• More than one valid Lewis structure with a given atomic
arrangement → resonance structures
• Resonance indicated by double-headed arrow
• Actual structure ~ a resonance hybrid of all resonance
structures
• Example: CO3
2–
29
a given chemical formula
resonace structure describe the position of
elctrons in compound.
30. Higher electron counts
• Number of electrons around central atom over 8 →
expanded shell & hypervalent molecules
• Frequently observed for elements of 3rd & higher periods
due to the d orbitals
• Example: PCl5, SF6
30
31. Molecular structure & Bonding
1. The Octet Rule
2. Lewis Diagram
3. Molecular Shapes & the VSEPR Model
4. Valence Bond (VB) theory
5. Molecular Orbital (MO) theory
31
32. VSEPR Model & Molecular Shapes
• Molecular shape: 3D arrangement of atoms that constitute a
molecule
• Experimentally determined by spectroscopic methods
• Valence-Shell Electron-Pair Repulsion (VSEPR) model
• Simple but useful to predict shape of small molecules of
main group elements
32
33. VSEPR Model & Molecular Shapes
• Valence-Shell Electron-Pair Repulsion (VSEPR) model
• Assumption: valence electron pairs adopt arrangements
that minimize repulsion between them
33
Electron
pairs
2 3 4 5 6
Geometry
linear
trigonal
planar
tetrahedral
trigonal
bipyramidal
octahedral
Standard
bond angle
180° 120° 109.5° 90°, 120° 90°
34. VSEPR Model & Molecular Shapes
• Valence-Shell Electron-Pair Repulsion (VSEPR) model
• Electron–electron repulsions decrease in sequence:
lone-lone pair > lone-bonding pair > bonding-bonding pair
triple -single bond > double -single bond > single -single bond
34
VSEPR model describe the 3D shape in space so
of compound
39. Molecular structure & Bonding
1. The Octet Rule
2. Lewis Diagram
3. Molecular Shapes & the VSEPR Model
4. Valence Bond (VB) theory
5. Molecular Orbital (MO) theory
39
40. Valence Bond (VB) theory
• Principle: chemical bonds ~ overlap of valence AOs on two
different nuclei so that both nuclei share pair of electrons
• Localized electron model (localized AOs)
• Homonuclear diatomic molecules
40
x
y
z
↑↓
↑↓ ↑↓
↑↓ ↑↓
F F
σ bond
H H
↑↓
σ bond
H2 F2
F (Z = 9) 1s2 2s2 2p5
41. Valence Bond (VB) theory
• Homonuclear diatomic molecules
41
x
y
z
O (Z = 8) 1s2 2s2 2p4
O2
π bond
↑↓
↑↓ ↑↓
↑ ↓
O O
σ
N (Z = 7) 1s2 2s2 2p3
N2
π
↑↓
↓ ↑
↑ ↓
N N
σ
π
observe empty orbital
in atom and
42. VB theory & Hybridization of AOs
Polyatomic molecules – Hybridization of AOs
• Hybridization of AOs: combinations of AOs to
mathematically obtain new AOs for molecular shapes in
terms of σ-bonds
• Number of hybrid orbitals = number of AOs mixed
• Label of hybrid orbitals reflects contributing AOs
• Type of hybrid orbitals varies with types of AOs mixed
42
46. Hybridization related to d-orbitals
sp3d sp3d2
s
p p p
sp3d
d d d d
d d d d
d
s
p p p
sp3d2
d d d d
d d d d
46
E E
47. Double & Triple bonds in VB theory
C C
↓↑
↓↑
↓↑
↓↑
↓↑
↓↑ ↓↑
↓↑
↑
↑
↓
↓
↑ ↓
47
48. Molecular structure & Bonding
1. The Octet Rule
2. Lewis Diagram
3. Molecular Shapes & the VSEPR Model
4. Valence Bond (VB) theory
5. Molecular Orbital (MO) theory
48
49. Molecular Orbital (MO) theory
• Generalization AO description of atoms to MO description of
molecules
• MOs: space spread whole molecule, in which a single
electron can occupy → delocalized electron model
• Linear Combination of Atomic Orbital (LCAO) approximation
49
when 2 atom bond form bond, and valence atomic
orbital is wave function and the property of
wwave fuction is the same
when 2 atom bond form bond, and valence atomic
orbital is wave function and the property of
wwave fuction is the same
50. Molecular Orbital (MO) theory
• MO arises from interactions between AOs if interactions are:
Allowed if symmetries of AOs are compatible
Efficient if region of overlap is significant
Efficient if AOs are relatively close in energy
• Number of MOs equal number of contributing AOs
50
s s px py px s pz s
51. Molecular Orbital (MO) theory
• Homonuclear diatomic molecules
destructive
interaction
↑↓
↑
H H
H2
51
E
anti-bonding MO
constructive
interaction
bonding MO
↑
53. Orbital energies for main group elements
53
J. B. Mann, T. L. Meek, L. C. Allen, J. Am. Chem. Soc., 2000, 122, 2780
54. MO diagram of X2 (X = O, F, Ne)
X X
X2
2s 2s
2p 2p
54
E
55. MO diagram of X2
X X
X2
2s 2s
2p 2p
2s 2s
2p 2p
X X
X2
55
E
(X = O, F, Ne) (X = Li, Be, B, C, N)
56. Molecular Orbital (MO) theory
• Heteronuclear diatomic molecules
antibonding MOs close to
less electronegative atom
bonding MOs close to
more electronegative atom
56
X Y
XY
less
electronegative
more
electronegative
E
57. MO diagram of CO
2s
2s
2p
2p
C O
CO
(-32.38 eV)
(-15.85 eV)
(-10.66 eV)
(-19.43 eV)
↑↓
↑↓
↑↓
↑ ↑
↑ ↑
↑↓
↑↓
↑↓ ↑↓
↑↓
57
E
58. Molecular Orbital (MO) theory
• Very unequal energies → resulting MOs with energies &
shapes closer to original AOs
58
∗
X Y
XY
E
less
electronegative
more
electronegative
59. MO diagrams of HF & LiF
𝜎
𝜎
𝜎∗
𝜋
2s
1s
2p
H F
HF
(-40.2 eV)
(-18.7 eV)
(-10.7 eV)
↑↓
↑
↑↓ ↑↓ ↑
↑↓
↑↓ ↑↓
↑↓
𝜎
𝜎
𝜎∗
2s
2s
2p
Li F
LiF
(-40.2 eV)
(-18.7 eV)
(-5.4 eV)
↑↓
↑
↑↓ ↑↓ ↑
↑↓
↑↓ ↑↓
↑↓
𝜋
59
E
E