Hydrogen bonding
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1. Hydrogen bonding occurs between hydrogen atoms attached to electronegative atoms like oxygen, fluorine, and nitrogen of one molecule and an electronegative atom of another molecule. 2. Water is able to form extensive hydrogen bonding networks between molecules due to each water molecule having two hydrogen atoms and two lone pairs of electrons on the oxygen atom. 3. The hydrogen bonding network in liquid water is responsible for its unique properties, while hydrogen bonding in ice forms its crystalline lattice structure.
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Hydrogen Bonding Powerpoint
1. Hydrogen bonding occurs when a hydrogen atom bonded to an electronegative atom like nitrogen, oxygen or fluorine interacts with another electronegative atom via electrostatic attraction between the hydrogen and the lone pair of electrons.
2. This leads to higher boiling points and enthalpies of vaporization for compounds that can form multiple hydrogen bonds like water. It also allows for the formation of structures like DNA base pairs and protein secondary structures.
3. Hydrogen bonding plays important roles in determining physical properties of compounds and enabling key biological molecules like DNA and proteins to form their functional three-dimensional structures.
Hydrogen bonding
This document discusses hydrogen bonding, beginning with its discovery in 1920 by Latimer and Rodebush. It defines hydrogen bonding as an electrostatic attraction between a hydrogen atom covalently bonded to an electronegative atom like F, O, or N, and another electronegative atom. Hydrogen bonds are weaker than covalent bonds but stronger than van der Waals attractions. The document outlines the conditions and types of hydrogen bonding, including intermolecular bonding between molecules and intramolecular bonding within molecules. It provides examples and discusses how hydrogen bonding affects various properties like boiling point, solubility, and spectral properties. In closing, it lists several applications of hydrogen bonding.
Covalent bond
This document discusses the octet rule for chemical bonding. It states that the octet rule describes the tendency of atoms to attain noble gas configurations by gaining, losing or sharing electrons to acquire eight electrons in their outer shell. There are exceptions when the octet is not complete, such as with boron trifluoride (BF3), or when it is expanded as in phosphorus pentachloride (PCl5) which has 10 electrons around the phosphorus atom. The limitations of the octet rule are also outlined.
Vander waals forces and its significance
The document discusses van der Waals forces, their characteristics, types (Keesom, Debye, London dispersion forces), and significance. Van der Waals forces are weak intermolecular forces between molecules caused by interactions between permanent or induced molecular dipoles. They are important in protein folding, graphite bonding, polymer formation, animal climbing/walking, and modern technologies that aim to mimic gecko adhesion.
Lattice energy
Lattice energy refers to the energy released when separate ions in the gas phase form an ionic crystal lattice. It can be calculated theoretically using the Born-Landé equation or experimentally using the Born-Haber cycle. The Born-Landé equation considers the electrostatic attraction and repulsive forces between ions, while the Born-Haber cycle uses standard enthalpy data and Hess's law. Lattice energy depends on factors like ion charge and size - higher charge or smaller ions lead to stronger electrostatic forces and higher lattice energy. Lattice energy is an important concept for understanding the properties and stability of ionic compounds.
Hydrogen bonding
1. Hydrogen bonding occurs between hydrogen atoms attached to electronegative atoms like oxygen, fluorine, and nitrogen of one molecule and an electronegative atom of another molecule.
2. Water is well-suited for hydrogen bonding due to its molecular structure, with oxygen atoms accepting hydrogen bonds. This hydrogen bonding network contributes to water's unique properties including its high boiling point.
3. An experiment is proposed to measure the effect of ethylene glycol or salt on the boiling point of water, with the hypothesis that ethylene glycol will increase the boiling point due to additional hydrogen bonding between water and ethylene glycol molecules.
Hyperconjugation - organic chemistry
This presentation describes the concept of Hyperconjugation in simple words, gives definition of hyperconjugation, explains why it is called as 'No bond Resonance' and gives the effects of hyperconjugation on the chemical properties of compounds: alkyl cations and their relative stability, alkyl radicals and their relative stability, alkenes and their relative stability, bond length, anomeric effect and Baker - Nathan effect.
Recommended
Hydrogen Bonding Powerpoint
1. Hydrogen bonding occurs when a hydrogen atom bonded to an electronegative atom like nitrogen, oxygen or fluorine interacts with another electronegative atom via electrostatic attraction between the hydrogen and the lone pair of electrons.
2. This leads to higher boiling points and enthalpies of vaporization for compounds that can form multiple hydrogen bonds like water. It also allows for the formation of structures like DNA base pairs and protein secondary structures.
3. Hydrogen bonding plays important roles in determining physical properties of compounds and enabling key biological molecules like DNA and proteins to form their functional three-dimensional structures.
Hydrogen bonding
This document discusses hydrogen bonding, beginning with its discovery in 1920 by Latimer and Rodebush. It defines hydrogen bonding as an electrostatic attraction between a hydrogen atom covalently bonded to an electronegative atom like F, O, or N, and another electronegative atom. Hydrogen bonds are weaker than covalent bonds but stronger than van der Waals attractions. The document outlines the conditions and types of hydrogen bonding, including intermolecular bonding between molecules and intramolecular bonding within molecules. It provides examples and discusses how hydrogen bonding affects various properties like boiling point, solubility, and spectral properties. In closing, it lists several applications of hydrogen bonding.
Covalent bond
This document discusses the octet rule for chemical bonding. It states that the octet rule describes the tendency of atoms to attain noble gas configurations by gaining, losing or sharing electrons to acquire eight electrons in their outer shell. There are exceptions when the octet is not complete, such as with boron trifluoride (BF3), or when it is expanded as in phosphorus pentachloride (PCl5) which has 10 electrons around the phosphorus atom. The limitations of the octet rule are also outlined.
Vander waals forces and its significance
The document discusses van der Waals forces, their characteristics, types (Keesom, Debye, London dispersion forces), and significance. Van der Waals forces are weak intermolecular forces between molecules caused by interactions between permanent or induced molecular dipoles. They are important in protein folding, graphite bonding, polymer formation, animal climbing/walking, and modern technologies that aim to mimic gecko adhesion.
Lattice energy
Lattice energy refers to the energy released when separate ions in the gas phase form an ionic crystal lattice. It can be calculated theoretically using the Born-Landé equation or experimentally using the Born-Haber cycle. The Born-Landé equation considers the electrostatic attraction and repulsive forces between ions, while the Born-Haber cycle uses standard enthalpy data and Hess's law. Lattice energy depends on factors like ion charge and size - higher charge or smaller ions lead to stronger electrostatic forces and higher lattice energy. Lattice energy is an important concept for understanding the properties and stability of ionic compounds.
Hydrogen bonding
1. Hydrogen bonding occurs between hydrogen atoms attached to electronegative atoms like oxygen, fluorine, and nitrogen of one molecule and an electronegative atom of another molecule.
2. Water is well-suited for hydrogen bonding due to its molecular structure, with oxygen atoms accepting hydrogen bonds. This hydrogen bonding network contributes to water's unique properties including its high boiling point.
3. An experiment is proposed to measure the effect of ethylene glycol or salt on the boiling point of water, with the hypothesis that ethylene glycol will increase the boiling point due to additional hydrogen bonding between water and ethylene glycol molecules.
Hyperconjugation - organic chemistry
This presentation describes the concept of Hyperconjugation in simple words, gives definition of hyperconjugation, explains why it is called as 'No bond Resonance' and gives the effects of hyperconjugation on the chemical properties of compounds: alkyl cations and their relative stability, alkyl radicals and their relative stability, alkenes and their relative stability, bond length, anomeric effect and Baker - Nathan effect.
Chemistry hydrogen bonding
Hydrogen bonding occurs when a hydrogen atom bonded to a highly electronegative atom such as nitrogen, oxygen or fluorine interacts with another electronegative atom. This leads to compounds having higher boiling points than expected due to the need to overcome hydrogen bonding interactions. Hydrogen bonding plays important roles in determining the structure of materials like DNA and proteins and was also utilized in the hydrogen bomb to dramatically increase its explosive yield.
Hydrogen bonding
Hydrogen bonding occurs between polar molecules where a hydrogen atom is bonded to a highly electronegative atom like oxygen or nitrogen. This partial positive charge on the hydrogen allows it to form an attractive force with another partially negatively charged atom. Water exhibits strong hydrogen bonding between its molecules, giving rise to properties like its high boiling point. Ethanol also exhibits hydrogen bonding between its -OH groups, allowing it to have a higher boiling point than similar sized molecules without this bonding, like methoxymethane. This hydrogen bonding also allows ethanol to be soluble in water whereas non-polar molecules cannot disrupt the water's hydrogen bonding network.
Resonance
Resonance effect occurs when a compound can be represented by two or more Lewis structures with the same arrangement of atoms. These contributing structures are called resonance structures.
The resonance structures are hypothetical and do not represent real molecules individually. Together they contribute to a resonance hybrid structure that is more stable than any individual resonance structure. The most stable contributing structure contributes the most to the resonance hybrid.
Conditions for resonance include sp or sp2 hybridized atoms involved, overlapping parallel p-orbitals to form a conjugated system. The sigma bond framework remains the same between structures, atomic positions do not change, and electron counts are equal.
Chemical Bonding
This document discusses the four main types of chemical bonds: ionic bonds, covalent bonds, hydrogen bonds, and metallic bonds. Ionic bonds involve the transfer of electrons between atoms. Covalent bonds involve the sharing of electrons between two atoms. Hydrogen bonds are electrostatic attractions between hydrogen atoms covalently bonded to electronegative atoms and another electronegative atom. Metallic bonds are electrostatic attractions between positively charged metal ions and delocalized electrons in metals. Examples of each type of bond are provided.
Valence bond teory
Valence bond theory describes the formation of covalent bonds between atoms through the overlap of atomic orbitals and the sharing of electron pairs. There are two types of orbital overlaps - sigma bonds form from head-on overlap along the internuclear axis, while pi bonds form from side-by-side overlap perpendicular to the axis. The extent of orbital overlap determines the bond strength. Valence bond theory also explains the directional properties of bonds in molecules and the octet rule by which atoms seek to attain the stable electron configuration of the nearest noble gas.
VSEPR Theory and molecular geometries
VSEPR (Valence Shell Electron Pair Repulsion) theory predicts molecular geometry based on the repulsion of electron pairs in the valence shell of an atom. It states that electron pairs will adopt a geometry that minimizes repulsion by maximizing distance between electron pairs. The number and type of electron pairs (lone pairs or bond pairs) determines the molecular shape. Lone pairs occupy more space than bond pairs due to greater repulsion from the single nucleus. Molecular geometry can be distorted from ideal shapes by the presence of lone pairs.
Hyperconjugation
Dr. Neelam from the Department of Chemistry presented on the topic of hyperconjugation. Hyperconjugation is the delocalization of σ-electrons from a C-H bond into an adjacent unsaturated system. It can occur in alkenes, alkynes, carbocations, and carbon radicals. The number of possible hyperconjugative structures equals the number of alpha hydrogens on sp3 hybridized carbon atoms. Hyperconjugation explains trends in stability and heats of hydrogenation between different alkenes. It is a permanent effect that does not change hybridization and is distance independent.
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Vsepr theory
The document summarizes Valence Shell Electron Pair Repulsion (VSEPR) Theory. [1] VSEPR Theory predicts molecular geometry based on electron pair repulsion around a central atom. [2] The theory states that electron pairs around an atom will position themselves as far apart as possible to minimize repulsion. [3] Molecular geometry can be determined by counting the number of electron pairs and their arrangement around the central atom.
Resonance----(Oganic)
Resonance structures represent different arrangements of electrons in a molecule that have the same positions of nuclei but different bonding patterns. Resonance contributes to the stability of molecules like benzene by delocalizing electrons across multiple equivalent structures. The actual structure of a molecule represented by resonance is a hybrid of the contributing structures, with bond lengths intermediate between single and double bonds. Delocalization of electrons is depicted using curved arrows between resonance structures.
Electronegativity
The electronegativity of an element is a measure of how strongly it attracts electrons in a covalent bond. Electronegativity increases left to right and top to bottom in a period, and metals have the lowest values while nonmetals have the highest. The difference in electronegativity between two bonded atoms indicates bond type - ionic bonds form when difference is >1.7, covalent bonds form when difference is <1.7, and polar covalent bonds form for differences in between. Electronegativity values can be used to predict and investigate bond types.
Lecture 19.1b- Bronsted-Lowry Acids & Bases
Here are the key points about measuring pH:
- pH is a measure of hydrogen ion concentration in a solution on a scale from 0-14.
- A pH below 7 is acidic, above 7 is basic, and 7 is neutral.
- The pH scale is logarithmic, so each unit change in pH represents a 10-fold change in [H+].
- Common indicators like litmus paper, phenolphthalein, and universal indicator can be used to estimate pH based on their color changes.
- pH meters provide a precise numerical measurement of pH by measuring voltage.
So in summary, pH is a quantitative measure of acidity or basicity that can be estimated qualitatively with indicators
Crystal field theory
This document summarizes Crystal Field Theory, which considers the electrostatic interactions between metal ions and ligands. It describes ligands and metal ions as point charges that can have attractive or repulsive forces. This causes the d orbitals of the metal ion to split into two sets depending on if the field created by the ligands is weak or strong. The theory explains color in coordination compounds as being caused by d-d electron transitions under the influence of ligands. However, it has limitations like not accounting for other metal orbitals or the partial covalent nature of metal-ligand bonds.
Molecular orbitals
Molecular orbital theory provides an approach to bonding that considers orbitals encompassing the entire molecule rather than just between atoms. Molecular orbitals form from the combination of atomic orbitals, which can combine constructively to form bonding orbitals or destructively to form antibonding orbitals. Bonding orbitals lower the energy and increase stability of the molecule, while antibonding orbitals have higher energy.
Hard-Soft-Acids-and-Bases-HSAB.pptx
This document discusses Hard and Soft Acids and Bases (HSAB) theory presented by Dr. Satish S. Kola. It defines characteristics of hard vs soft acids and bases, with hard acids/bases being small with high oxidation states and no d-electrons, while soft acids/bases are large with low oxidation states and many d-electrons. Applications of HSAB principles are discussed, including predicting complex formation and metal catalyst poisoning. The theoretical basis involves concepts like pi-bonding, electrostatic interactions, and polarizability. Limitations are noted where inherent acid/base strength may override HSAB predictions.
Molecular orbital theory
Molecular orbital theory (MOT) is an alternative model to valence bond theory that explains how atomic orbitals from different atoms combine to form molecular orbitals. The linear combination of atomic orbitals (LCAO) method considers the probability of finding electrons in atomic orbitals from different atoms. According to the LCAO method, molecular orbitals are formed from constructive and destructive interference of atomic orbitals. MOT can be used to explain bonding in homonuclear diatomic molecules like N2 and O2, heteronuclear diatomic molecules like CO and NO, and polyatomic molecules like CO2 and SF6. It can also describe bonding in octahedral transition metal complexes like hexaaquoferrate(II) ion
Vsepr theory
The document discusses the Valence Shell Electron Pair Repulsion (VSEPR) theory, which proposes that the shape of a molecule is determined by the repulsion between electron pairs in the valence shell of the central atom. The theory was developed in 1940 by Sidgwick and Powell and further refined by Nyholm and Gillespie in 1957. The main postulates of VSEPR theory are that electron pairs repel each other and will adopt an arrangement that minimizes this repulsion, treating multiple bonds as single pairs. The repulsion decreases in the order of lone pair-lone pair, lone pair-bond pair, and bond pair-bond pair interactions. Examples are then given of molecular shapes predicted by VSEPR theory without
Mesomeric Effect
Mesomeric (or resonance) effect refers to the polarity produced in a molecule through delocalization of pi electrons between bonds or a bond and lone pair. This effect can increase or decrease electron density in different parts of a molecule. Negative mesomeric effects are shown by groups like nitro (-NO2) that withdraw electron density, while positive effects are shown by groups like hydroxyl (-OH) that release electron density. Understanding mesomeric effects helps explain a molecule's reactivity toward electrophiles and nucleophiles by determining where electron density is increased or decreased.
Crystal field theory
Crystal Field Theory explains the colors of transition metal complexes based on ligand-metal interactions. The electrostatic interaction between ligands and metal d-orbitals splits the d-orbital energies. For an octahedral complex, the d-orbitals point directly at ligands have higher energy than those that bisect ligands. This splitting pattern determines if the complex is high or low spin, which then dictates its color and magnetic properties. The spectrochemical series orders ligands by their ability to cause crystal field splitting, correlating ligand type with complex color.
Nature of hydrogen bonds in liquids and crystals
The International Institute for Science, Technology and Education (IISTE). Science, Technology and Medicine Journals Call for Academic Manuscripts
Intermolecular Forces in Hydrogen Bonding - Pooja N
This document discusses intermolecular forces, specifically hydrogen bonding. It defines hydrogen bonding as a strong dipole-dipole interaction that occurs between a hydrogen atom bonded to an electronegative atom like oxygen, nitrogen, or fluorine. Examples of hydrogen bonding include water, hydrogen fluoride, alcohols, and carboxylic acids. Hydrogen bonding increases boiling points and melting points by bringing molecules closer together. It also increases the solubility of compounds like alcohols and glucose in water. The strength of hydrogen bonding depends on factors like the electronegativity of the bonded atom.
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Chemistry hydrogen bonding
Hydrogen bonding occurs when a hydrogen atom bonded to a highly electronegative atom such as nitrogen, oxygen or fluorine interacts with another electronegative atom. This leads to compounds having higher boiling points than expected due to the need to overcome hydrogen bonding interactions. Hydrogen bonding plays important roles in determining the structure of materials like DNA and proteins and was also utilized in the hydrogen bomb to dramatically increase its explosive yield.
Hydrogen bonding
Hydrogen bonding occurs between polar molecules where a hydrogen atom is bonded to a highly electronegative atom like oxygen or nitrogen. This partial positive charge on the hydrogen allows it to form an attractive force with another partially negatively charged atom. Water exhibits strong hydrogen bonding between its molecules, giving rise to properties like its high boiling point. Ethanol also exhibits hydrogen bonding between its -OH groups, allowing it to have a higher boiling point than similar sized molecules without this bonding, like methoxymethane. This hydrogen bonding also allows ethanol to be soluble in water whereas non-polar molecules cannot disrupt the water's hydrogen bonding network.
Resonance
Resonance effect occurs when a compound can be represented by two or more Lewis structures with the same arrangement of atoms. These contributing structures are called resonance structures.
The resonance structures are hypothetical and do not represent real molecules individually. Together they contribute to a resonance hybrid structure that is more stable than any individual resonance structure. The most stable contributing structure contributes the most to the resonance hybrid.
Conditions for resonance include sp or sp2 hybridized atoms involved, overlapping parallel p-orbitals to form a conjugated system. The sigma bond framework remains the same between structures, atomic positions do not change, and electron counts are equal.
Chemical Bonding
This document discusses the four main types of chemical bonds: ionic bonds, covalent bonds, hydrogen bonds, and metallic bonds. Ionic bonds involve the transfer of electrons between atoms. Covalent bonds involve the sharing of electrons between two atoms. Hydrogen bonds are electrostatic attractions between hydrogen atoms covalently bonded to electronegative atoms and another electronegative atom. Metallic bonds are electrostatic attractions between positively charged metal ions and delocalized electrons in metals. Examples of each type of bond are provided.
Valence bond teory
Valence bond theory describes the formation of covalent bonds between atoms through the overlap of atomic orbitals and the sharing of electron pairs. There are two types of orbital overlaps - sigma bonds form from head-on overlap along the internuclear axis, while pi bonds form from side-by-side overlap perpendicular to the axis. The extent of orbital overlap determines the bond strength. Valence bond theory also explains the directional properties of bonds in molecules and the octet rule by which atoms seek to attain the stable electron configuration of the nearest noble gas.
VSEPR Theory and molecular geometries
VSEPR (Valence Shell Electron Pair Repulsion) theory predicts molecular geometry based on the repulsion of electron pairs in the valence shell of an atom. It states that electron pairs will adopt a geometry that minimizes repulsion by maximizing distance between electron pairs. The number and type of electron pairs (lone pairs or bond pairs) determines the molecular shape. Lone pairs occupy more space than bond pairs due to greater repulsion from the single nucleus. Molecular geometry can be distorted from ideal shapes by the presence of lone pairs.
Hyperconjugation
Dr. Neelam from the Department of Chemistry presented on the topic of hyperconjugation. Hyperconjugation is the delocalization of σ-electrons from a C-H bond into an adjacent unsaturated system. It can occur in alkenes, alkynes, carbocations, and carbon radicals. The number of possible hyperconjugative structures equals the number of alpha hydrogens on sp3 hybridized carbon atoms. Hyperconjugation explains trends in stability and heats of hydrogenation between different alkenes. It is a permanent effect that does not change hybridization and is distance independent.
Chemical bond
A chemical bond is a lasting attraction between atoms that enables the formation of chemical compounds or substance . The bond may result from the electrostatic force of attraction between atoms with opposite charges, or through the sharing of electrons as in the covalent bonds........
Inductive effect pptx
This document discusses electronic displacement in organic compounds. It describes two types of electronic displacement: permanent displacement including inductive, resonance, and mesomeric effects; and temporary displacement through electromeric effects. Inductive effects are further broken down into +I effects where groups donate electron density and -I effects where groups withdraw electron density. Examples of inductive effects include their impact on acid/base strength, stability of carbocations/carbanions, and dipole moments.
Vsepr theory
The document summarizes Valence Shell Electron Pair Repulsion (VSEPR) Theory. [1] VSEPR Theory predicts molecular geometry based on electron pair repulsion around a central atom. [2] The theory states that electron pairs around an atom will position themselves as far apart as possible to minimize repulsion. [3] Molecular geometry can be determined by counting the number of electron pairs and their arrangement around the central atom.
Resonance----(Oganic)
Resonance structures represent different arrangements of electrons in a molecule that have the same positions of nuclei but different bonding patterns. Resonance contributes to the stability of molecules like benzene by delocalizing electrons across multiple equivalent structures. The actual structure of a molecule represented by resonance is a hybrid of the contributing structures, with bond lengths intermediate between single and double bonds. Delocalization of electrons is depicted using curved arrows between resonance structures.
Electronegativity
The electronegativity of an element is a measure of how strongly it attracts electrons in a covalent bond. Electronegativity increases left to right and top to bottom in a period, and metals have the lowest values while nonmetals have the highest. The difference in electronegativity between two bonded atoms indicates bond type - ionic bonds form when difference is >1.7, covalent bonds form when difference is <1.7, and polar covalent bonds form for differences in between. Electronegativity values can be used to predict and investigate bond types.
Lecture 19.1b- Bronsted-Lowry Acids & Bases
Here are the key points about measuring pH:
- pH is a measure of hydrogen ion concentration in a solution on a scale from 0-14.
- A pH below 7 is acidic, above 7 is basic, and 7 is neutral.
- The pH scale is logarithmic, so each unit change in pH represents a 10-fold change in [H+].
- Common indicators like litmus paper, phenolphthalein, and universal indicator can be used to estimate pH based on their color changes.
- pH meters provide a precise numerical measurement of pH by measuring voltage.
So in summary, pH is a quantitative measure of acidity or basicity that can be estimated qualitatively with indicators
Crystal field theory
This document summarizes Crystal Field Theory, which considers the electrostatic interactions between metal ions and ligands. It describes ligands and metal ions as point charges that can have attractive or repulsive forces. This causes the d orbitals of the metal ion to split into two sets depending on if the field created by the ligands is weak or strong. The theory explains color in coordination compounds as being caused by d-d electron transitions under the influence of ligands. However, it has limitations like not accounting for other metal orbitals or the partial covalent nature of metal-ligand bonds.
Molecular orbitals
Molecular orbital theory provides an approach to bonding that considers orbitals encompassing the entire molecule rather than just between atoms. Molecular orbitals form from the combination of atomic orbitals, which can combine constructively to form bonding orbitals or destructively to form antibonding orbitals. Bonding orbitals lower the energy and increase stability of the molecule, while antibonding orbitals have higher energy.
Hard-Soft-Acids-and-Bases-HSAB.pptx
This document discusses Hard and Soft Acids and Bases (HSAB) theory presented by Dr. Satish S. Kola. It defines characteristics of hard vs soft acids and bases, with hard acids/bases being small with high oxidation states and no d-electrons, while soft acids/bases are large with low oxidation states and many d-electrons. Applications of HSAB principles are discussed, including predicting complex formation and metal catalyst poisoning. The theoretical basis involves concepts like pi-bonding, electrostatic interactions, and polarizability. Limitations are noted where inherent acid/base strength may override HSAB predictions.
Molecular orbital theory
Molecular orbital theory (MOT) is an alternative model to valence bond theory that explains how atomic orbitals from different atoms combine to form molecular orbitals. The linear combination of atomic orbitals (LCAO) method considers the probability of finding electrons in atomic orbitals from different atoms. According to the LCAO method, molecular orbitals are formed from constructive and destructive interference of atomic orbitals. MOT can be used to explain bonding in homonuclear diatomic molecules like N2 and O2, heteronuclear diatomic molecules like CO and NO, and polyatomic molecules like CO2 and SF6. It can also describe bonding in octahedral transition metal complexes like hexaaquoferrate(II) ion
Vsepr theory
The document discusses the Valence Shell Electron Pair Repulsion (VSEPR) theory, which proposes that the shape of a molecule is determined by the repulsion between electron pairs in the valence shell of the central atom. The theory was developed in 1940 by Sidgwick and Powell and further refined by Nyholm and Gillespie in 1957. The main postulates of VSEPR theory are that electron pairs repel each other and will adopt an arrangement that minimizes this repulsion, treating multiple bonds as single pairs. The repulsion decreases in the order of lone pair-lone pair, lone pair-bond pair, and bond pair-bond pair interactions. Examples are then given of molecular shapes predicted by VSEPR theory without
Mesomeric Effect
Mesomeric (or resonance) effect refers to the polarity produced in a molecule through delocalization of pi electrons between bonds or a bond and lone pair. This effect can increase or decrease electron density in different parts of a molecule. Negative mesomeric effects are shown by groups like nitro (-NO2) that withdraw electron density, while positive effects are shown by groups like hydroxyl (-OH) that release electron density. Understanding mesomeric effects helps explain a molecule's reactivity toward electrophiles and nucleophiles by determining where electron density is increased or decreased.
Crystal field theory
Crystal Field Theory explains the colors of transition metal complexes based on ligand-metal interactions. The electrostatic interaction between ligands and metal d-orbitals splits the d-orbital energies. For an octahedral complex, the d-orbitals point directly at ligands have higher energy than those that bisect ligands. This splitting pattern determines if the complex is high or low spin, which then dictates its color and magnetic properties. The spectrochemical series orders ligands by their ability to cause crystal field splitting, correlating ligand type with complex color.
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Similar to Hydrogen bonding
Nature of hydrogen bonds in liquids and crystals
The International Institute for Science, Technology and Education (IISTE). Science, Technology and Medicine Journals Call for Academic Manuscripts
Intermolecular Forces in Hydrogen Bonding - Pooja N
This document discusses intermolecular forces, specifically hydrogen bonding. It defines hydrogen bonding as a strong dipole-dipole interaction that occurs between a hydrogen atom bonded to an electronegative atom like oxygen, nitrogen, or fluorine. Examples of hydrogen bonding include water, hydrogen fluoride, alcohols, and carboxylic acids. Hydrogen bonding increases boiling points and melting points by bringing molecules closer together. It also increases the solubility of compounds like alcohols and glucose in water. The strength of hydrogen bonding depends on factors like the electronegativity of the bonded atom.
Hydrogen bonding pdf.pdf
1.Weak forces of attraction
2.Concepts of Hydrogen bonding
3.Types of hydrogen bonding
4.Properties of hydrogen bond.
5.Methods of detection of hydrogen bond.
6.Importance of Hydrogen bonding.
7.Vander walls forces
a.Ion-dipole
b.Dipole-dipole
c.London forces.
8.Origin of hydrogen bonds.
9.Consequences of hydrogen bonding.
10.Ice has less density than water.
11.Intermolecular forces.
hydrogen bond 2 by KK Sahu sir
Introduction
History
Definition
Types of H bond
Hydrogen bond in water
Bifurcated and over - Coordinated hydrogen bond in water
Hydrogen bonds in DNA and proteins
Hydrogen bonds in polymers
Systematic hydrogen bond
Importance of hydrogen bond
Conclusion
References
Benefits Of Alkaline Water(1)
This document discusses ionized water and its potential health benefits. It begins by explaining how an ionized water unit works, using electrolysis to separate tap water into reduced and oxidized waters. Reduced water has a more negative redox potential, meaning it has excess electrons that can be donated to neutralize reactive oxygen species. The document argues that drinking reduced ionized water may help the body avoid illness by blocking the oxidation of healthy cells caused by an overproduction of free radicals.
Element# of protons# of neutrons# of electronsCarbon (atomic.pdf
Element
# of protons
# of neutrons
# of electrons
Carbon (atomic # 6)
Neutral atom:
Ion:
Isotope:
Neutral atom:
Ion:
Isotope:
Neutral atom:
Ion:
Isotope:
Oxygen (atomic # 8)
Neutral atom:
Ion:
Isotope:
Neutral atom:
Ion:
Isotope:
Neutral atom:
Ion:
Isotope:
Hydrogen (atomic # 1)
Neutral atom:
Ion:
Isotope:
Neutral atom:
Ion:
Isotope:
Neutral atom:
Ion:
Isotope:
Nitrogen (atomic # 7)
Neutral atom:
Ion:
Isotope:
Neutral atom:
Ion:
Isotope:
Neutral atom:
Ion:
Isotope:
How does the number of electrons affect how many bonds are formed by these atoms?
Define the following (see boxes on pp 31, 33)
Polar covalent bond
Non-polar covalent bond
Ionic bond
Hydrogen bond
Hydrophobic
Hydrophilic
Amphipathic
Element
# of protons
# of neutrons
# of electrons
Carbon (atomic # 6)
Neutral atom:
Ion:
Isotope:
Neutral atom:
Ion:
Isotope:
Neutral atom:
Ion:
Isotope:
Oxygen (atomic # 8)
Neutral atom:
Ion:
Isotope:
Neutral atom:
Ion:
Isotope:
Neutral atom:
Ion:
Isotope:
Hydrogen (atomic # 1)
Neutral atom:
Ion:
Isotope:
Neutral atom:
Ion:
Isotope:
Neutral atom:
Ion:
Isotope:
Nitrogen (atomic # 7)
Neutral atom:
Ion:
Isotope:
Neutral atom:
Ion:
Isotope:
Neutral atom:
Ion:
Isotope:
Solution
Element
# of protons
# of electrons
# of neutrons
Ion
Isotopes
Carbon
6
6
6
C+4 or C-4
Carbon-12 which has 6 protons and 6 neutrons and 6 electrons.
Carbon-14 which has 6 protons, 8 neutrons and 6 electrons.
Oxygen
8
8
8
Oxygen generally forms O-2. It has 6 electrons in the outermost orbit.
Three stable isotopes of oxygen are found- O-16,O17 and O-18.
Hydrogen
1
1
1
Hydrogen has a charge of +1.
Hydrogen has 3 stable isotopes – protium (1 proton and 1 electron), deuterium (1 proton, 1
electron and 1 neutron) and tritium (1 proton, 1 electron and 2 neutrons)
Nitrogen
7
7
7
Usually has the charge -3.
N-14 has odd number of protons and neutrons (7 each). N-15 has protons and 8 neutrons.
Carbon does not usually form an ion by loss or gain of 4 electrons. It shares its 4 electrons with 4
others to complete its orbit and forms a covalent bond.
Oxygen has 6 valence lectrons in its outermost shell, therefore it can form 2 bonds
Hydrogen atom has one valence electron in its outermost orbit. The maximum capacity of its
orbital is 2 electrons.
Nonpolar covalent bond
Non-polar covalent bond is formed when the atoms (with similar electron affinity) share
electrons equally. The electronegativity difference between the two bonded atoms is usually less
than 0.5. Examples – diatomic molecules H2 and N2.
Polar covalent bond
Polar covalent bond is formed when the electrons are not equally shared i.e. the atom with higher
electronegativity (affinity) pulls the electrons towards itself. The compound which is formed has
electrostatic potential and has the ability to form weak bonds with other polar molecules. The
electronegativity difference between bonded atoms is 0.5 to 1.9. Examples include water, ozone.
Ionic bond – the valence electrons are completely transferred from one atom to another. The
metal lo.
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weak interaction by KK Sahu sir
Introduction
2. Defination
3. Type of weak interaction
Hydrogen bond
Ionic bond
Vander waals attraction
Hydrophobic interaction
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5. Conclusion
6. Reference
Notes water day day 2
Primarily due to polarity, water has remarkable properties that support life. Water molecules are polar, with slightly positive hydrogen atoms and slightly negative oxygen atoms, allowing hydrogen bonds to form between molecules. These bonds give water high surface tension, heat capacity, and solubility, enabling it to dissolve many substances and regulate temperatures essential for organisms.
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みなさんこんにちはこれ何文字まで入るの？40文字以下不可とか本当に意味わからないけどこれ限界文字数書いてないからマジでやばい文字数いけるんじゃないの？えこ...
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Hydrogen bonding
- 1. Hydrogen Bonding 8.4 Water (Prelim) 8.4.2 The wide distribution and importance of water on Earth is a consequence of its molecular structure and hydrogen bonding
- 2. Syllabus Points Stage 6 preliminary unit on Water. describe hydrogen bonding between molecules plan and perform an investigation to identify and describe the effect of anti-freeze or salt on the boiling point of water
- 4. What is a Hydrogen Bond? Hydrogen bonds are the attractive force between the hydrogen attached to an electronegative atom of one molecule and an electronegative atom of a different molecule. Hydrogen (H) needs to be attached to a highly electronegative Oxygen (O), Fluoride (F), Nitrogen (N) atom.
- 5. Why are O, F & N so attractive?? In each of these molecules note: 1. A hydrogen atom! 2. the hydrogen is attached directly to one of the most electronegative elements, causing the hydrogen to acquire a significant amount of positive charge. 3. Each of the elements to which the hydrogen is attached is not only significantly negative, but also has at least one "active" lone pair. In these interactions, a hydrogen atom is the donor, and the respective electronegative atom is the acceptor of the bond. Electronegativity Values F (4.0) O (3.5) N(3.0)
- 6. Hydrogen Bonding in Water Water could be considered as the "perfect" hydrogen bonded system. Each Oygen (O) has 2 lone pairs of electrons therefore, 4 Hydrogen bonds (H-bonds) can be formed. Hydrogen-bonding network formed in liquid water is responsible for many of the essential and unique properties of water. Ice forms a crystalline lattice, held together by a multitude of H-bonds
- 7. Hydrogen Bonding in molecules the hydrogen is partially positive and attracted to the partially negative charge on the nitrogen. Nitrogen has only one lone pair, thus, only one hydrogen bond can be made to each nitrogen.
- 8. Activity
- 10. Experiment Plan and perform an investigation to identify and describe the effect of anti-freeze (ethylene glycol) or salt on the boiling point of water
- 11. How would you plan this experiment? Plan and perform an investigation to identify and describe the effect of anti-freeze (ethylene glycol) or salt on the boiling point of water
- 12. What safety steps need to be taken?
- 13. What is your hypothesis?
- 14. Collecting data
- 15. What can be concluded?
- 16. Blog Site: hydrogenbonding2y.blogspot.com Questions?