Chemistry hydrogen bonding
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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.
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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.
H bonding
Hydrogen bonding occurs between a hydrogen atom covalently bonded to an electronegative element like fluorine, oxygen, or nitrogen, and another electronegative atom with an unshared pair of electrons. There are two types of hydrogen bonding: intermolecular, which occurs between molecules, and intramolecular, which occurs within a single molecule and can form rings. Properties like boiling point, solubility, and acidity are affected by whether hydrogen bonding is intermolecular or intramolecular.
Intermolecular forces
The document discusses the three main types of intermolecular forces: dipole-dipole forces, London dispersion forces, and hydrogen bonding. It describes each type of force, including that dipole-dipole forces exist between polar molecules, London dispersion forces exist between all molecules, and hydrogen bonding specifically refers to the attraction between hydrogen atoms bonded to electronegative atoms like N, O, or F. The document also discusses how these intermolecular forces can explain the physical properties of liquids and some solids.
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
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 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.
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.
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.
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.
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.
H bonding
Hydrogen bonding occurs between a hydrogen atom covalently bonded to an electronegative element like fluorine, oxygen, or nitrogen, and another electronegative atom with an unshared pair of electrons. There are two types of hydrogen bonding: intermolecular, which occurs between molecules, and intramolecular, which occurs within a single molecule and can form rings. Properties like boiling point, solubility, and acidity are affected by whether hydrogen bonding is intermolecular or intramolecular.
Intermolecular forces
The document discusses the three main types of intermolecular forces: dipole-dipole forces, London dispersion forces, and hydrogen bonding. It describes each type of force, including that dipole-dipole forces exist between polar molecules, London dispersion forces exist between all molecules, and hydrogen bonding specifically refers to the attraction between hydrogen atoms bonded to electronegative atoms like N, O, or F. The document also discusses how these intermolecular forces can explain the physical properties of liquids and some solids.
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
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 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.
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.
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.
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.
2012 Orbital Hybrization, Sigma and Pi Bonds
Carbon forms four equal hybrid orbitals through sp3 hybridization to allow methane to adopt its tetrahedral electron geometry. This involves one s orbital and three p orbitals combining to form four new hybrid orbitals. Sigma bonds are formed by the head-on overlap of hybrid orbitals. Sp2 hybridization with one s and two p orbitals results in trigonal planar geometries like ethene. Pi bonds in double and triple bonds involve overlap of unhybridized p orbitals above and below the sigma bond.
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.
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.
Covalent bond
This document defines and provides examples of covalent bonds. It discusses the history of the term being coined by Irving Langmuir in 1919. A covalent bond is defined as a chemical link where electrons are shared between two atoms. Examples given are the covalent bonds between oxygen and hydrogen in H2O and between hydrogen atoms in H2. The document also describes the different types of covalent bonds - single, double, and triple - based on the number of electron pairs shared between the bonded atoms.
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.
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
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.
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
Valence Bond theory & Hybridization
The document discusses valence bond theory and hybridization. Valence bond theory describes how covalent bonds are formed via the overlapping of atomic orbitals. Hybridization involves mixing atomic orbitals of similar energy to form new hybrid orbitals. Different types of hybridization (sp, sp2, sp3 etc.) result in hybrid orbitals with varying compositions and orientations that can explain the geometry of molecules.
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.
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.
Chapter 3 Alkenes
This document summarizes key information about alkenes (olefins):
1) Alkenes contain carbon-carbon double bonds and are classified as unsaturated hydrocarbons. Common examples include ethylene and propene.
2) Alkenes undergo characteristic reactions such as addition of halogens, hydrogenation to form alkanes, hydration and polymerization. Many of these reactions follow Markovnikov's rule.
3) Alkenes are industrially important as monomers for polymers like polyethylene, polypropylene, PVC and polystyrene. Ethylene and propylene are the largest volume organic chemicals produced.
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.
S & p block elements
The elements in which the valence electron enters the s orbital are called s block elements.
The elements in which the valence electron enters the p orbital are called p block elements.
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.
Intermolecular forces
This document discusses different types of intermolecular forces: permanent dipole-dipole forces, van der Waals forces, and hydrogen bonding. It describes how each type of force arises and provides examples to illustrate the effects of these forces on molecular properties like boiling points. Specifically, it notes that hydrogen bonding results in water having a higher boiling point than other hydrides and ice floating due to its lower density than liquid water.
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.
Resonance ,inductive effect -----(Organic )
Resonance occurs when a molecule can be represented by multiple Lewis structures and the actual structure is a hybrid of all the contributing structures. A resonance hybrid is the average structure between all the canonical or contributing Lewis structures. Delocalization of electrons through resonance leads to decreased potential energy and increased stability compared to any single contributing structure. Electron donating groups cause positive resonance or inductive effects while electron withdrawing groups cause negative resonance or inductive effects. Resonance and inductive effects influence acidity, basicity, and reactivity.
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.
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.
Basics of organic
This document provides an introduction to organic chemistry, focusing on functional groups and how they influence molecular properties and reactivity. It defines functional groups as parts of molecules that determine geometry, physical properties, and reactivity. The document discusses common functional groups containing heteroatoms and pi bonds, and how these features create reactive sites. It also examines how functional groups influence intermolecular forces, physical properties like boiling point and solubility, and chemical reactivity.
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2012 Orbital Hybrization, Sigma and Pi Bonds
Carbon forms four equal hybrid orbitals through sp3 hybridization to allow methane to adopt its tetrahedral electron geometry. This involves one s orbital and three p orbitals combining to form four new hybrid orbitals. Sigma bonds are formed by the head-on overlap of hybrid orbitals. Sp2 hybridization with one s and two p orbitals results in trigonal planar geometries like ethene. Pi bonds in double and triple bonds involve overlap of unhybridized p orbitals above and below the sigma bond.
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.
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.
Covalent bond
This document defines and provides examples of covalent bonds. It discusses the history of the term being coined by Irving Langmuir in 1919. A covalent bond is defined as a chemical link where electrons are shared between two atoms. Examples given are the covalent bonds between oxygen and hydrogen in H2O and between hydrogen atoms in H2. The document also describes the different types of covalent bonds - single, double, and triple - based on the number of electron pairs shared between the bonded atoms.
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.
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
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.
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
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The document discusses valence bond theory and hybridization. Valence bond theory describes how covalent bonds are formed via the overlapping of atomic orbitals. Hybridization involves mixing atomic orbitals of similar energy to form new hybrid orbitals. Different types of hybridization (sp, sp2, sp3 etc.) result in hybrid orbitals with varying compositions and orientations that can explain the geometry of molecules.
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.
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.
Chapter 3 Alkenes
This document summarizes key information about alkenes (olefins):
1) Alkenes contain carbon-carbon double bonds and are classified as unsaturated hydrocarbons. Common examples include ethylene and propene.
2) Alkenes undergo characteristic reactions such as addition of halogens, hydrogenation to form alkanes, hydration and polymerization. Many of these reactions follow Markovnikov's rule.
3) Alkenes are industrially important as monomers for polymers like polyethylene, polypropylene, PVC and polystyrene. Ethylene and propylene are the largest volume organic chemicals produced.
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.
S & p block elements
The elements in which the valence electron enters the s orbital are called s block elements.
The elements in which the valence electron enters the p orbital are called p block elements.
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.
Intermolecular forces
This document discusses different types of intermolecular forces: permanent dipole-dipole forces, van der Waals forces, and hydrogen bonding. It describes how each type of force arises and provides examples to illustrate the effects of these forces on molecular properties like boiling points. Specifically, it notes that hydrogen bonding results in water having a higher boiling point than other hydrides and ice floating due to its lower density than liquid water.
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.
Resonance ,inductive effect -----(Organic )
Resonance occurs when a molecule can be represented by multiple Lewis structures and the actual structure is a hybrid of all the contributing structures. A resonance hybrid is the average structure between all the canonical or contributing Lewis structures. Delocalization of electrons through resonance leads to decreased potential energy and increased stability compared to any single contributing structure. Electron donating groups cause positive resonance or inductive effects while electron withdrawing groups cause negative resonance or inductive effects. Resonance and inductive effects influence acidity, basicity, and reactivity.
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.
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2.Concepts of Hydrogen bonding
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7.Vander walls forces
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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 BOND .pptx
Hydrogen bonding occurs when a hydrogen atom covalently bonded to an electronegative atom like nitrogen, oxygen or fluorine interacts with another electronegative atom. This interaction is stronger than van der Waals forces but weaker than covalent or ionic bonds. Hydrogen bonding plays a key role in the structures of water and DNA, enabling water to have high boiling and melting points and DNA to form its double helix structure. There are two types of hydrogen bonds: intermolecular between molecules and intramolecular within the same molecule.
H bonding
Hydrogen bonding occurs between a hydrogen atom covalently bonded to an electronegative element like fluorine, oxygen, or nitrogen, and another electronegative atom with an unshared pair of electrons. There are two types of hydrogen bonding: intermolecular, which occurs between molecules, and intramolecular, which occurs within a single molecule and can form rings. Properties like boiling point, solubility, and acidity are affected by whether hydrogen bonding is intermolecular or intramolecular.
intermolecular_forces.ppt
This document discusses intermolecular forces and the properties they impart to liquids and solids. It describes three main types of intermolecular forces: dipole-dipole forces between polar molecules, London dispersion forces present in all molecules, and strong hydrogen bonding between a hydrogen atom and electronegative atoms like N, O, F. Hydrogen bonding gives ice its unusual properties and is important in biological molecules like DNA. The strengths of these different intermolecular forces influence physical properties of substances like melting and boiling points. Liquids exhibit surface tension, capillary action, and viscosity due to the attractive intermolecular forces between their molecules.
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Chemistry hydrogen bonding
- 1. 02 02 HYDROGEN BONDING &APPLICATIONS By-Nikhil Sharma Md Azharuddin Yash 24
- 2. HYDROGEN BONDING • Hydrogen bonding is a special case of dipole-dipole interactions. • When hydrogen is bonded to a highly electronegative atom (such as nitrogen, oxygen, fluorine), the bonding electron pair is drawn towards the electronegative atom. 24
- 3. Hydrogen bonding F H F H..... δ- δ+ δ- δ+ Hydrogen has no inner shell electron and is very small in size, the positive charge density developed is high The nucleus of hydrogen atom is exposed to attraction by nearby electron cloud, a lone pair electrons on the electronegative atom 24
- 4. H F . .. . . . H F . .. . . . H F . .. . . . H F . .. . . . Hydrogen Bonding in Hydrogen Fluoride 24
- 5. Hydrogen bonding • More Examples 24
- 6. CONDITION FOR H-BONDING – A hydrogen atom must be directly bonded to a highly electronegative atom (e.g. F, O and N) – An unbonded pair of electrons (lone pair electrons) is presented on the electronegative atom The size of the electronegative atom should be small. The smaller the size ,the greater is the electrostastics attraction. 24
- 7. NOTES • Intermolecular hydrogen bond: Hydrogen bond formed between two molecules Intramolecular hydrogen bond: Hydrogen bond formed between two different atoms in the same molecule Intermolecular hydrogen bond is stronger than van der Waals’ forces N O O O H 24
- 8. Strength of hydrogen bond • Hydrogen bonding is very weak bond.The strength of hydrogen bond is intermediate between the weak van der waal’s forces and the strong covalent bonds.Thus whereas the bond dissociation energy of a covalent bond is 209-418 Kj/mol.That for H-bond is only 12.6-41.8 kj/mol. 24
- 9. Diff. between H-bond &Covalent bond H-bond It involves dipole-dipole attractive interactions. It is formed between a hydrogen and a highly electronegative atom such as F,O,N. The strength of this bond is very small. Covalent bond It involves sharing of electrons. It is formed between two electronegative atoms which may be out of the same element or of different elements . The bond strength of this bond is sufficiently high. 24
- 10. Consequences of H-bond High melting & boiling points: The compound having H-bond show abnormally high melting and boiling point. Due to the containing hydrogen bond to the fact that some extra energy is needed to break these bonds. 24
- 11. Q1: Why do NH3, H2O and HF have abnormally high boiling point? • Explanation: – N, O and F are highly electronegative atoms • Formation of intermolecular hydrogen bonds in their hydrides. – Intermolecular hydrogen bonds are much stronger than van der Waals’ forces – More energy is needed to break the hydrogen bonds in NH3, H2O and HF 24
- 12. Q2: Explain why the order of boiling point is H2O > HF > NH3 Explanation H2O can form 2 hydrogen bond per molecule NH3 and HF can form only 1 hydrogen bond per molecule The boiling point of H2O is higher than NH3 and HF No. of H atoms available for hydrogen bond formation No. of lone pair electrons available No. of hydrogen bonds can form per molecule NH3 3 1 1 H2O 2 2 2 HF 1 3 1 24
- 13. Q2: Explain why the order of boiling point is H2O > HF > NH3 – Comparing N and F, F is more electronegative – The intermolecular hydrogen bond formed between HF molecules is stronger than NH3 – HF has a higher boiling point than NH3 24
- 14. Explanation of lower density of ice than water • In case of solid ice,the hydrogen bonding gives rise to a cage like structure of water molecules as shown.As a matter of fact ,each water molecule is linked tetrahedrally to four water molecules tetrahedrally. 24
- 15. This section of water is frozen This section of water is liquid 24
- 16. APPLICATIONS 24
- 17. Importance of Hydrogen Bonding in Physical Phenomena • Boiling Point and Solubility of Alcohols – The boiling points of alcohols are higher than the thiols because there are intermolecular hydrogen bonds between alcohol molecules, but only dipole-dipole interactions between molecules of thiols Alcohols: Thiols: 24
- 18. • Alcohols of low molecular masses are soluble in water because they can form intermolecular hydrogen bonds with water molecules 24
- 19. Importance of Hydrogen Bonding in Biochemistry • Secondary structure Of Protein 24
- 20. Hydrogen Bonding in DNA DNA stands for Deoxyribose Nucleic Acid DNA contains genetic information Consists of 2 macromolecular strands spiralling around each other in the form of a double helix 24
- 21. H-BOMB • The hydrogen bomb is a nuclear weapon that used a mixture of fission and fusion (Hydrogen fuel) to produce a massive explosion • It produces 700-times the energy released by the bombs exploded in Hiroshima and Nagasaki. 24
- 22. Nagasaki MushroomNagasaki Mushroom cloud.cloud. 13 km high!13 km high! 24
- 23. • It is made up of essentially, 3 things:- 1.Implosion Type Fission Bomb:- It contains a core of U-235 and a tamper of U-238 2.Fusion Bomb:- A cylinder containing a core rod of Plutonium-239 surrounded by lithium-duterate fuel with a tamper of U-238. 3.Polystyrene foam:- The empty spaces are filled by it. 24
- 24. 24