1) Chemical bonds form between atoms as a way for atoms to decrease their potential energy and become more stable.
2) There are three main types of chemical bonds: ionic bonds, which form between ions and are electrostatic in nature; covalent bonds, which form through the sharing of electron pairs between atoms; and metallic bonds, which form between positive metal ions and delocalized electrons in metals.
3) Ionic compounds consist of positive and negative ions that are attracted to one another via ionic bonding. The ions arrange themselves in repeating three-dimensional patterns called crystal lattices to minimize potential energy.
Chemical bonds form between atoms to achieve more stable arrangements with lower potential energy. The main types of bonds are ionic, covalent, and metallic. Ionic bonds form between oppositely charged ions, covalent bonds form when atoms share electrons, and metallic bonds result from delocalized electrons being shared among many atoms in a metal. The degree of ionic or covalent character in a bond depends on the electronegativity difference between the atoms. Molecular compounds are made of molecules held together by covalent bonds, while ionic compounds are made of ionic bonds between cations and anions.
Chapter 6.2 : Covalent Bonding and Molecular CompoundsChris Foltz
Covalent bonding occurs when atoms share valence electrons to achieve stable electron configurations. Molecules are formed when atoms bond covalently, with molecular formulas indicating the types and numbers of atoms. Lewis structures represent molecules by showing bonding pairs and unshared electron pairs. Exceptions to the octet rule include hydrogen and boron. Resonance structures are used to depict molecules that cannot be represented by a single Lewis structure.
This document provides notes on covalent bonding concepts including:
- The octet rule and how atoms form bonds to gain or lose electrons
- Ionic, covalent, and polar covalent bonds defined by electron transfer and sharing
- Bond polarity determined by electronegativity differences
- Molecular polarity based on bond polarity and molecular geometry
- Intermolecular forces of dispersion, dipole, and hydrogen bonding explained
Electron-dot notation represents valence electrons of atoms using dots placed around the element's symbol. It can show molecules by combining the notations of atoms joined by covalent bonds, represented as shared pairs of dots or dashes. Lewis structures use this notation to represent covalent bonding, showing shared pairs between atoms but not unshared electron pairs. Atoms can form multiple bonds by sharing more than one electron pair, shown as double or triple bonds. Some molecules like ozone cannot be fully represented by a single Lewis structure and instead use resonance structures that average the bonding.
Here are the key points about polar and nonpolar covalent bonds:
- Nonpolar covalent bonds form between atoms with similar electronegativity. Electrons are shared equally. Examples include H2, Cl2, etc.
- Polar covalent bonds form between atoms with different electronegativity. Electrons are pulled slightly closer to the more electronegative atom. This creates partial positive and negative charges called dipoles.
- Molecules with polar bonds are polar molecules. They are attracted to electric fields due to the separation of charge in the dipoles. Water is a common example.
- Molecules with nonpolar covalent bonds have equal charge distribution and are nonpolar molecules. They
1. The document provides an introductory overview of basic chemical terms and concepts including electron configuration, chemical bonds, substances, mixtures, compounds, elements, and properties of single particles.
2. Key concepts discussed include the atomic mass unit, relative atomic mass, amount of substance, mole, molar mass, Avogadro's law, electron configuration, types of chemical bonds, and hybridization.
3. The document serves as a recapitulation of lectures covering fundamental chemistry topics at an introductory level.
This document summarizes key concepts from Chapter 8 of a chemistry textbook, including:
- The three main types of chemical bonds are covalent, ionic, and metallic bonds. Covalent bonds result from shared electron pairs, ionic bonds from electron transfers, and metallic bonds hold pure metals together.
- Lewis structures represent electron arrangements and bonding using dots around atom symbols. The octet rule states atoms seek full outer shells of 8 electrons.
- Ionic bonds form when metals donate electrons to nonmetals, giving stable noble gas configurations. They form crystalline lattices with strong attractions.
- Covalent bonds share electron pairs so atoms both attain full outer shells. Polarity arises from differing electronegativities.
1. The document discusses the historical development of atomic theory from Democritus' idea of indivisible atoms to the modern atomic model.
2. Key contributors included Dalton who proposed atoms of different elements have different properties, Thomson who discovered the electron, and Rutherford whose gold foil experiment showed the atom's small, dense nucleus.
3. The modern atomic model consists of a small, positively charged nucleus surrounded by electrons in regions of probable location called electron clouds.
Chemical bonds form between atoms to achieve more stable arrangements with lower potential energy. The main types of bonds are ionic, covalent, and metallic. Ionic bonds form between oppositely charged ions, covalent bonds form when atoms share electrons, and metallic bonds result from delocalized electrons being shared among many atoms in a metal. The degree of ionic or covalent character in a bond depends on the electronegativity difference between the atoms. Molecular compounds are made of molecules held together by covalent bonds, while ionic compounds are made of ionic bonds between cations and anions.
Chapter 6.2 : Covalent Bonding and Molecular CompoundsChris Foltz
Covalent bonding occurs when atoms share valence electrons to achieve stable electron configurations. Molecules are formed when atoms bond covalently, with molecular formulas indicating the types and numbers of atoms. Lewis structures represent molecules by showing bonding pairs and unshared electron pairs. Exceptions to the octet rule include hydrogen and boron. Resonance structures are used to depict molecules that cannot be represented by a single Lewis structure.
This document provides notes on covalent bonding concepts including:
- The octet rule and how atoms form bonds to gain or lose electrons
- Ionic, covalent, and polar covalent bonds defined by electron transfer and sharing
- Bond polarity determined by electronegativity differences
- Molecular polarity based on bond polarity and molecular geometry
- Intermolecular forces of dispersion, dipole, and hydrogen bonding explained
Electron-dot notation represents valence electrons of atoms using dots placed around the element's symbol. It can show molecules by combining the notations of atoms joined by covalent bonds, represented as shared pairs of dots or dashes. Lewis structures use this notation to represent covalent bonding, showing shared pairs between atoms but not unshared electron pairs. Atoms can form multiple bonds by sharing more than one electron pair, shown as double or triple bonds. Some molecules like ozone cannot be fully represented by a single Lewis structure and instead use resonance structures that average the bonding.
Here are the key points about polar and nonpolar covalent bonds:
- Nonpolar covalent bonds form between atoms with similar electronegativity. Electrons are shared equally. Examples include H2, Cl2, etc.
- Polar covalent bonds form between atoms with different electronegativity. Electrons are pulled slightly closer to the more electronegative atom. This creates partial positive and negative charges called dipoles.
- Molecules with polar bonds are polar molecules. They are attracted to electric fields due to the separation of charge in the dipoles. Water is a common example.
- Molecules with nonpolar covalent bonds have equal charge distribution and are nonpolar molecules. They
1. The document provides an introductory overview of basic chemical terms and concepts including electron configuration, chemical bonds, substances, mixtures, compounds, elements, and properties of single particles.
2. Key concepts discussed include the atomic mass unit, relative atomic mass, amount of substance, mole, molar mass, Avogadro's law, electron configuration, types of chemical bonds, and hybridization.
3. The document serves as a recapitulation of lectures covering fundamental chemistry topics at an introductory level.
This document summarizes key concepts from Chapter 8 of a chemistry textbook, including:
- The three main types of chemical bonds are covalent, ionic, and metallic bonds. Covalent bonds result from shared electron pairs, ionic bonds from electron transfers, and metallic bonds hold pure metals together.
- Lewis structures represent electron arrangements and bonding using dots around atom symbols. The octet rule states atoms seek full outer shells of 8 electrons.
- Ionic bonds form when metals donate electrons to nonmetals, giving stable noble gas configurations. They form crystalline lattices with strong attractions.
- Covalent bonds share electron pairs so atoms both attain full outer shells. Polarity arises from differing electronegativities.
1. The document discusses the historical development of atomic theory from Democritus' idea of indivisible atoms to the modern atomic model.
2. Key contributors included Dalton who proposed atoms of different elements have different properties, Thomson who discovered the electron, and Rutherford whose gold foil experiment showed the atom's small, dense nucleus.
3. The modern atomic model consists of a small, positively charged nucleus surrounded by electrons in regions of probable location called electron clouds.
Covalent bonds form when two nonmetal atoms share one or more pairs of valence electrons to achieve a noble gas configuration. This sharing of electrons results in the formation of molecules with lower potential energy than the individual atoms. Covalent bonds can be either polar or nonpolar depending on whether the electrons are shared equally or unequally between the bonded atoms. Multiple bonds are also possible where two or more pairs of electrons are shared, such as double and triple bonds.
The document discusses chemical bonding and Lewis structures. It begins by defining a chemical bond as the force that holds atoms together, and discusses how atoms combine or share electrons to form ionic or covalent bonds. It then explains Lewis structures, showing how to draw the Lewis dot symbols and determine the hybridization of atoms. Examples are provided of writing Lewis structures for different molecules like CCl4 and NH4+. The document also discusses exceptions to the octet rule and theories of covalent bonding like valence bond theory and hybridization theory.
This document discusses the electronic configuration of carbon and how it forms bonds. It explains that carbon normally forms four single bonds by undergoing sp3 hybridization, where one 2s orbital and three 2p orbitals combine to form four new hybrid orbitals oriented toward the corners of a tetrahedron. It also discusses sp2 and sp hybridization which allow carbon to form multiple and triple bonds. The document contrasts primary covalent, ionic, and coordinate covalent bonds from secondary bonds formed by hydrogen bonding and van der Waals forces.
For Chem 1:
Significanceof the ELectron in Bonding
The Octet Rule
Lewis Symbol/Structures
Formal Charge
Polyatomic Ions
Types of Bonds (Ionic, Covalent, Coordinate Covalent, Metallic Bonds, Multiple Bonds)
Exceptions to the Octet Rules
Oxidation Number is not included in the class discussion and exam. ;D
The document discusses chemical bonding and basic concepts related to bonding. It defines valence electrons as the outer shell electrons that participate in chemical bonding. It describes the octet rule where atoms gain, lose, or share electrons to acquire a noble gas configuration with eight electrons in their outer shell. The document also defines and distinguishes between ionic bonds, which involve electron transfer, and covalent bonds, where electrons are shared between atoms.
This document discusses different types of chemical bonds including ionic bonds, covalent bonds, and the factors that determine bond strength. It defines ionic bonds as electrostatic attractions between oppositely charged ions, and covalent bonds as the sharing of valence electrons between nonmetals. Bond length and bond energy are inversely related, with shorter, stronger bonds forming from increased orbital overlap and multiple bonds. Lewis structures are used to represent electron arrangements in molecules using dots or lines to indicate bonding and non-bonding electron pairs.
This document discusses different types of chemical bonds including ionic bonds, covalent bonds, and polar covalent bonds. It explains how ionic bonds form between a metal and nonmetal when electrons are transferred, covalent bonds form through shared electron pairs, and polar covalent bonds result from unequal electron sharing. The document also covers bond energies, dipole moments, electronegativity, and Lewis structures.
1. The document discusses chemical bonding, specifically ionic and covalent bonds.
2. Ionic bonds form when electrons are transferred from one atom to another, giving positively charged cations and negatively charged anions. Covalent bonds form through the mutual sharing of electron pairs between atoms.
3. Factors that favor ionic bond formation include differences in electronegativity between atoms, ionization energies, electron affinities, and lattice energies of resulting ionic compounds. Ionic compounds have properties like being crystalline solids, poor electrical conductors, high melting/boiling points, and solubility in polar solvents.
Here are the key points about polar and nonpolar covalent bonds:
- Nonpolar covalent bonds form between atoms with similar electronegativity. Electrons are equally shared. Examples include H2, Cl2, etc.
- Polar covalent bonds form between atoms with different electronegativity. Electrons are unequally shared, resulting in partial positive and negative charges on the atoms. Examples include HCl, H2O.
- Polar molecules are attracted to each other due to their partial charges. They are soluble in polar solvents like water. Nonpolar molecules are not attracted and are soluble in nonpolar solvents like hexane.
- Most bonds have some ionic character based on electrone
This document discusses covalent bonding between non-metal atoms. It defines covalent bonding as the sharing of electron pairs between non-metal atoms. Valence electrons are electrons in the outer shell that can be gained or lost. Many elements are stable when they have eight valence electrons through sharing. Examples of covalent bonding include hydrogen and chlorine atoms sharing electrons to achieve full outer shells.
The document summarizes key concepts about covalent bonding from a chemistry textbook chapter:
1) Covalent bonds form when two nonmetal atoms share one or more pairs of electrons to achieve a noble gas configuration, forming molecules like H2, O2, and CO2.
2) Molecular compounds formed by covalent bonds tend to have lower melting and boiling points than ionic compounds due to the weaker nature of the covalent bond.
3) Electron dot structures and Lewis diagrams are used to represent how atoms share electrons to form single, double or triple covalent bonds in molecules like H2O and NH3.
A chemical bond is a lasting attraction between atoms that enables the formation of chemical compounds. 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
The document summarizes different types of chemical bonding:
1. Ionic bonding results from the attraction between oppositely charged ions
2. Covalent bonding results from the sharing of electron pairs between atoms
3. Metallic bonding allows for electron delocalization and mobility in metal solids due to overlapping vacant orbitals, contributing to metals' electrical and thermal conductivity properties.
Nature of Bonding in Organic Molecules - Sahana KamathBebeto G
The document discusses the nature of bonding in organic molecules. It describes the different types of hybridization that carbon undergoes, including sp3, sp2, and sp hybridization, and how this allows carbon to form single, double, and triple bonds. It also discusses sigma and pi bonds, bond length, bond angle, bond energy, localized and delocalized bonds, and hydrogen bonding. In summary:
1) Carbon can undergo sp3, sp2, and sp hybridization to form tetrahedral, trigonal planar, and linear geometries respectively and allow carbon to form single, double, and triple bonds.
2) Sigma bonds are formed by head-on overlap of orbitals while pi bonds involve side
chemical bonding and molecular structure class 11sarunkumar31
hybridisation, bonding and antiboding, dipole moment, VSPER theory, Molecular orbital diagram, Phosphorous pentachloride, ionic bond, bond order, bond enthalpy, bond dissociation, sp and sp2hybridisation, hydrogen bonding,electron pair,lone pair repulsion, resonance structure of ozone, how to find electron pair and lone pair, sp3 hybridization of methane.
Chapter 6.1 : Introduction to Chemical BondingChris Foltz
This document discusses chemical bonding. It defines chemical bonds as how most atoms are joined together in nature. It describes the two main types of chemical bonds: ionic bonding which results from the transfer of electrons between ions, and covalent bonding which results from the sharing of electron pairs between atoms. Atoms form chemical bonds to decrease their potential energy and become more stable. Bonds are rarely purely ionic or covalent, but instead exist on a spectrum depending on the electronegativity difference between the atoms.
The document discusses chemical bonding and molecular structure. It begins by introducing different types of chemical bonds including ionic and covalent bonds. It then discusses how valence electrons are distributed in molecules as either bond pairs or lone pairs according to the octet rule. Examples of drawing Lewis structures for different molecules like NH3, CO2, and SO32- are provided. The document also discusses exceptions to the octet rule and concepts like formal charge, resonance structures, and bond order. Finally, it introduces VSEPR theory for predicting molecular geometry and discusses how bond properties like length and energy depend on factors like bond order.
Chemical bonds form between atoms to achieve more stable arrangements with lower potential energy. The type of bonding depends on differences in electronegativity between atoms. Ionic bonds form between ions, covalent bonds involve shared electron pairs, and metallic bonds result from delocalized electrons shared among many atoms in a lattice. Molecular geometry and intermolecular forces also influence molecular properties.
This document provides information about covalent bonding including:
- Covalent bonds result from the sharing of valence electrons between nonmetal atoms.
- Molecules form when two or more atoms are bonded covalently. Diatomic molecules like O2, N2, and F2 contain two atoms of the same element bonded together.
- Single covalent bonds involve the sharing of one pair of electrons, double bonds two pairs, and triple bonds three pairs. Lewis structures are used to represent electron arrangements in molecules.
This document discusses different types of chemical bonds: ionic bonds form between metals and nonmetals and result from the transfer of electrons. Covalent bonds form between nonmetals and result from the sharing of valence electrons. Molecules are formed when two or more atoms are bonded covalently. Covalent bonds can be single, double or triple bonds depending on how many electron pairs are shared. Bond length and bond strength are related, with shorter bonds generally being stronger. Lewis structures are used to represent how atoms are arranged and bonded in a molecule using electron dots.
This document provides information about molecular and ionic compounds, including:
- Molecular compounds are formed by covalent bonds between nonmetal atoms, while ionic compounds involve metal and nonmetal atoms bonded by ionic bonds.
- Molecular formulas show the actual number and type of atoms in a molecule, while ionic formulas use the lowest whole number ratio.
- Covalent bonds are represented by electron dot structures that show how atoms share electrons to achieve stable configurations. Multiple and coordinate covalent bonds are also discussed.
- Polarity arises in polar covalent bonds due to unequal electron sharing. Polar molecules have dipole moments while intermolecular forces include hydrogen bonding, dipole-dipole interactions, and
Covalent bonds form when two nonmetal atoms share one or more pairs of valence electrons to achieve a noble gas configuration. This sharing of electrons results in the formation of molecules with lower potential energy than the individual atoms. Covalent bonds can be either polar or nonpolar depending on whether the electrons are shared equally or unequally between the bonded atoms. Multiple bonds are also possible where two or more pairs of electrons are shared, such as double and triple bonds.
The document discusses chemical bonding and Lewis structures. It begins by defining a chemical bond as the force that holds atoms together, and discusses how atoms combine or share electrons to form ionic or covalent bonds. It then explains Lewis structures, showing how to draw the Lewis dot symbols and determine the hybridization of atoms. Examples are provided of writing Lewis structures for different molecules like CCl4 and NH4+. The document also discusses exceptions to the octet rule and theories of covalent bonding like valence bond theory and hybridization theory.
This document discusses the electronic configuration of carbon and how it forms bonds. It explains that carbon normally forms four single bonds by undergoing sp3 hybridization, where one 2s orbital and three 2p orbitals combine to form four new hybrid orbitals oriented toward the corners of a tetrahedron. It also discusses sp2 and sp hybridization which allow carbon to form multiple and triple bonds. The document contrasts primary covalent, ionic, and coordinate covalent bonds from secondary bonds formed by hydrogen bonding and van der Waals forces.
For Chem 1:
Significanceof the ELectron in Bonding
The Octet Rule
Lewis Symbol/Structures
Formal Charge
Polyatomic Ions
Types of Bonds (Ionic, Covalent, Coordinate Covalent, Metallic Bonds, Multiple Bonds)
Exceptions to the Octet Rules
Oxidation Number is not included in the class discussion and exam. ;D
The document discusses chemical bonding and basic concepts related to bonding. It defines valence electrons as the outer shell electrons that participate in chemical bonding. It describes the octet rule where atoms gain, lose, or share electrons to acquire a noble gas configuration with eight electrons in their outer shell. The document also defines and distinguishes between ionic bonds, which involve electron transfer, and covalent bonds, where electrons are shared between atoms.
This document discusses different types of chemical bonds including ionic bonds, covalent bonds, and the factors that determine bond strength. It defines ionic bonds as electrostatic attractions between oppositely charged ions, and covalent bonds as the sharing of valence electrons between nonmetals. Bond length and bond energy are inversely related, with shorter, stronger bonds forming from increased orbital overlap and multiple bonds. Lewis structures are used to represent electron arrangements in molecules using dots or lines to indicate bonding and non-bonding electron pairs.
This document discusses different types of chemical bonds including ionic bonds, covalent bonds, and polar covalent bonds. It explains how ionic bonds form between a metal and nonmetal when electrons are transferred, covalent bonds form through shared electron pairs, and polar covalent bonds result from unequal electron sharing. The document also covers bond energies, dipole moments, electronegativity, and Lewis structures.
1. The document discusses chemical bonding, specifically ionic and covalent bonds.
2. Ionic bonds form when electrons are transferred from one atom to another, giving positively charged cations and negatively charged anions. Covalent bonds form through the mutual sharing of electron pairs between atoms.
3. Factors that favor ionic bond formation include differences in electronegativity between atoms, ionization energies, electron affinities, and lattice energies of resulting ionic compounds. Ionic compounds have properties like being crystalline solids, poor electrical conductors, high melting/boiling points, and solubility in polar solvents.
Here are the key points about polar and nonpolar covalent bonds:
- Nonpolar covalent bonds form between atoms with similar electronegativity. Electrons are equally shared. Examples include H2, Cl2, etc.
- Polar covalent bonds form between atoms with different electronegativity. Electrons are unequally shared, resulting in partial positive and negative charges on the atoms. Examples include HCl, H2O.
- Polar molecules are attracted to each other due to their partial charges. They are soluble in polar solvents like water. Nonpolar molecules are not attracted and are soluble in nonpolar solvents like hexane.
- Most bonds have some ionic character based on electrone
This document discusses covalent bonding between non-metal atoms. It defines covalent bonding as the sharing of electron pairs between non-metal atoms. Valence electrons are electrons in the outer shell that can be gained or lost. Many elements are stable when they have eight valence electrons through sharing. Examples of covalent bonding include hydrogen and chlorine atoms sharing electrons to achieve full outer shells.
The document summarizes key concepts about covalent bonding from a chemistry textbook chapter:
1) Covalent bonds form when two nonmetal atoms share one or more pairs of electrons to achieve a noble gas configuration, forming molecules like H2, O2, and CO2.
2) Molecular compounds formed by covalent bonds tend to have lower melting and boiling points than ionic compounds due to the weaker nature of the covalent bond.
3) Electron dot structures and Lewis diagrams are used to represent how atoms share electrons to form single, double or triple covalent bonds in molecules like H2O and NH3.
A chemical bond is a lasting attraction between atoms that enables the formation of chemical compounds. 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
The document summarizes different types of chemical bonding:
1. Ionic bonding results from the attraction between oppositely charged ions
2. Covalent bonding results from the sharing of electron pairs between atoms
3. Metallic bonding allows for electron delocalization and mobility in metal solids due to overlapping vacant orbitals, contributing to metals' electrical and thermal conductivity properties.
Nature of Bonding in Organic Molecules - Sahana KamathBebeto G
The document discusses the nature of bonding in organic molecules. It describes the different types of hybridization that carbon undergoes, including sp3, sp2, and sp hybridization, and how this allows carbon to form single, double, and triple bonds. It also discusses sigma and pi bonds, bond length, bond angle, bond energy, localized and delocalized bonds, and hydrogen bonding. In summary:
1) Carbon can undergo sp3, sp2, and sp hybridization to form tetrahedral, trigonal planar, and linear geometries respectively and allow carbon to form single, double, and triple bonds.
2) Sigma bonds are formed by head-on overlap of orbitals while pi bonds involve side
chemical bonding and molecular structure class 11sarunkumar31
hybridisation, bonding and antiboding, dipole moment, VSPER theory, Molecular orbital diagram, Phosphorous pentachloride, ionic bond, bond order, bond enthalpy, bond dissociation, sp and sp2hybridisation, hydrogen bonding,electron pair,lone pair repulsion, resonance structure of ozone, how to find electron pair and lone pair, sp3 hybridization of methane.
Chapter 6.1 : Introduction to Chemical BondingChris Foltz
This document discusses chemical bonding. It defines chemical bonds as how most atoms are joined together in nature. It describes the two main types of chemical bonds: ionic bonding which results from the transfer of electrons between ions, and covalent bonding which results from the sharing of electron pairs between atoms. Atoms form chemical bonds to decrease their potential energy and become more stable. Bonds are rarely purely ionic or covalent, but instead exist on a spectrum depending on the electronegativity difference between the atoms.
The document discusses chemical bonding and molecular structure. It begins by introducing different types of chemical bonds including ionic and covalent bonds. It then discusses how valence electrons are distributed in molecules as either bond pairs or lone pairs according to the octet rule. Examples of drawing Lewis structures for different molecules like NH3, CO2, and SO32- are provided. The document also discusses exceptions to the octet rule and concepts like formal charge, resonance structures, and bond order. Finally, it introduces VSEPR theory for predicting molecular geometry and discusses how bond properties like length and energy depend on factors like bond order.
Chemical bonds form between atoms to achieve more stable arrangements with lower potential energy. The type of bonding depends on differences in electronegativity between atoms. Ionic bonds form between ions, covalent bonds involve shared electron pairs, and metallic bonds result from delocalized electrons shared among many atoms in a lattice. Molecular geometry and intermolecular forces also influence molecular properties.
This document provides information about covalent bonding including:
- Covalent bonds result from the sharing of valence electrons between nonmetal atoms.
- Molecules form when two or more atoms are bonded covalently. Diatomic molecules like O2, N2, and F2 contain two atoms of the same element bonded together.
- Single covalent bonds involve the sharing of one pair of electrons, double bonds two pairs, and triple bonds three pairs. Lewis structures are used to represent electron arrangements in molecules.
This document discusses different types of chemical bonds: ionic bonds form between metals and nonmetals and result from the transfer of electrons. Covalent bonds form between nonmetals and result from the sharing of valence electrons. Molecules are formed when two or more atoms are bonded covalently. Covalent bonds can be single, double or triple bonds depending on how many electron pairs are shared. Bond length and bond strength are related, with shorter bonds generally being stronger. Lewis structures are used to represent how atoms are arranged and bonded in a molecule using electron dots.
This document provides information about molecular and ionic compounds, including:
- Molecular compounds are formed by covalent bonds between nonmetal atoms, while ionic compounds involve metal and nonmetal atoms bonded by ionic bonds.
- Molecular formulas show the actual number and type of atoms in a molecule, while ionic formulas use the lowest whole number ratio.
- Covalent bonds are represented by electron dot structures that show how atoms share electrons to achieve stable configurations. Multiple and coordinate covalent bonds are also discussed.
- Polarity arises in polar covalent bonds due to unequal electron sharing. Polar molecules have dipole moments while intermolecular forces include hydrogen bonding, dipole-dipole interactions, and
This document provides an overview of chemical bonding. It defines a chemical bond as a force of attraction between atoms or ions that holds atoms together in molecules or compounds. Atoms form bonds to achieve stable electron configurations. There are three main types of bonds: ionic, covalent, and metallic. Ionic bonds form through the transfer of electrons between metals and nonmetals. Covalent bonds form through the sharing of electrons, usually between nonmetals. Metallic bonds involve the pooling of electrons between metal atoms. The document further explores bond formation and properties.
Lecture 8.2- Lewis Dot Structures for MoleculesMary Beth Smith
The document discusses ionic and covalent bonding. It explains how to draw Lewis dot structures to show electron sharing between atoms to form single, double or triple covalent bonds. Examples are given of molecules like H2O, NH3, CH4, CO2, and O3 that form different types of covalent bonds through electron sharing.
1. Covalent bonds form when two atoms share one or more pairs of valence electrons in order to achieve a stable octet of electrons.
2. Molecules are formed when atoms are bonded together by covalent bonds, and molecular compounds are composed of molecules.
3. Molecular compounds tend to have lower melting and boiling points than ionic compounds and many are gases or liquids at room temperature.
The document provides an overview of chemical bonding principles including ionic bonding, covalent bonding, and molecular geometry. It discusses how atoms interact to achieve stable electronic configurations through ionic bonding by exchanging electrons or covalent bonding by sharing electrons. Lewis structures are introduced as a way to represent valence electrons in molecules and determine molecular geometry and polarity based on electron pair arrangements around central atoms. Key concepts covered include octet rule, electronegativity, bond polarity, and using Lewis structures to systematically determine molecular structure characteristics.
CHEMICAL BONDING AND MOLECULAR STRUCTUREniralipatil
Chemical bonding can occur via ionic bonds or covalent bonds. Ionic bonds form when electrons are transferred from one atom to another, leaving cation and anion. Covalent bonds form when atoms share electrons via overlapping orbitals. The octet rule states that atoms seek to obtain eight electrons in their valence shell. Hybridization is the mixing of atomic orbitals to form new hybrid orbitals for bonding. Common hybridizations include sp, sp2, and sp3 which determine molecular geometry. Bond properties like order, length, energy are influenced by hybridization.
The document discusses chemical bonding and molecular structures. It explains that chemical bonding occurs through ionic bonding via the transfer of electrons between atoms, or covalent bonding via the sharing of electron pairs between atoms. It also describes molecular geometry models including VSEPR theory, which predicts the three-dimensional arrangements of atoms in molecules based on electron pair repulsion. Common molecular shapes such as linear, trigonal planar, tetrahedral and octahedral are defined.
This document provides a summary of a lecture on inorganic chemistry. It discusses chemical bonds, including the definition of a chemical bond as a force that holds atoms together in a stable molecule. It describes four primary types of chemical bonds: ionic, covalent, coordinate covalent, and metallic. Covalent bonds are formed by the sharing of electron pairs between atoms. Coordinate covalent or dative bonds occur when both electrons in the shared pair come from one atom. Examples of compounds containing different bond types are also provided.
Covalent bonds form when atoms share valence electrons in order to achieve stable electron configurations like the noble gases. Atoms form single, double or triple covalent bonds by sharing one, two or three pairs of electrons. Molecular compounds are held together by covalent bonds and have lower melting and boiling points than ionic compounds. Molecular formulas show the types and numbers of atoms in a molecule but not their arrangement, which can be represented using structural formulas.
Lewis symbols and the octet rule are used to represent valence electrons and chemical bonding. Lewis symbols show valence electrons as dots around the element symbol. The octet rule states that atoms bond to gain, lose, or share electrons to achieve an octet of 8 valence electrons. Exceptions include odd total electron molecules, molecules where atoms have less than an octet, and hypervalent molecules where the central atom has more than 8 electrons.
The document provides information about chemical bonds including ionic bonds, covalent bonds, and bond energies. It defines ionic and covalent bonding, discusses factors that determine lattice energy of ionic compounds, introduces electronegativity and bond polarity. It also covers Lewis structures, resonance structures, and exceptions to the octet rule. Bond enthalpies, which measure bond strength, are discussed along with average bond enthalpies from bond dissociation data.
Chemical bonds are formed by the sharing or transfer of valence electrons between atoms. Valence electrons play an important role in bond formation as atoms seek to achieve stable electronic configurations like noble gases. There are two main types of bonds:
1) Ionic bonds are formed by the transfer of electrons from metals to nonmetals, resulting in positively charged cations and negatively charged anions that are attracted to each other.
2) Covalent bonds are formed by the sharing of electron pairs between nonmetals. Atoms share electrons to achieve stable octet configurations. Single, double, and triple covalent bonds are distinguished by the number of electron pairs shared.
Lewis structures use dots or crosses to represent valence
CH 4 CHEMICAL BONDING AND MOLECULAR STRUCTURE.pdfLUXMIKANTGIRI
English chapter we will discuss about bonding how the molecules and the ions are in texting as a molecule make the structure there energy their transmission and other
The document discusses chemical bonding and molecular structures. It begins by defining a chemical bond as the force that binds two atoms together within a molecule. It then discusses the different types of bonds ranked by decreasing bond strength - ionic, covalent, coordinate, hydrogen, and Van der Waals. Ionic bonds form through the transfer of electrons from metals to nonmetals. Covalent bonds form through the sharing of electron pairs between atoms. The document also discusses bond parameters such as bond length, bond order, bond energy, bond angle, and dipole moment. It introduces concepts such as Lewis structures, formal charge, resonance structures, and hybridization. It concludes with an overview of valence bond theory and molecular orbital theory.
There are three main types of chemical bonds: ionic, covalent, and metallic. Ionic bonds involve the electrostatic attraction between oppositely charged ions. Covalent bonds involve the sharing of electrons between atoms. Metallic bonds involve metal atoms bonded to several other metal atoms. The strength of bonds can be estimated from bond enthalpy values, which provide the energy required to break chemical bonds. Bond strength increases as bond length decreases.
There are three main types of chemical bonds: ionic, covalent, and metallic. Ionic bonds involve the electrostatic attraction between oppositely charged ions. Covalent bonds involve the sharing of electrons between atoms. Metallic bonds involve metal atoms bonded to several other metal atoms. The strength of bonds can be estimated from bond enthalpies, which measure the energy required to break chemical bonds. Bond strength increases as bond length decreases.
Similar to Chapter6 chemicalbonding-100707021031-phpapp01 (20)
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Covalent bonding results from the sharing of electron pairs between
two atoms
In purely covalent bond, electrons shared equally between two atoms
7. Ionic or Covalent?
Bonding is rarely purely one or the other
Depending on how strongly the atoms attract electrons,
falls somewhere between
Electronegativity (EN) atom’s ability to attract
electrons
Degree of bonding between atoms of 2 elements being
ionic or covalent estimated by calculating difference in
elements’ ENs
8. Example
Fluorine’s EN = 4.0, Cesium’s
EN = 0.7
4.0-0.7 = 3.3
According to table, F-Cs is ionic
The greater the difference, the
more ionic the bond
9.
Bonding between atoms with EN difference of less than or equal to (≤) 1.7
has ionic character less than or equal to (≤) 50%
Classified as covalent
Bonding between atoms of same element is completely covalent
10. Nonpolar-covalent Bonds
H-H bond has 0% ionic character
Nonpolar-covalent bond a covalent bond in which the bonding
electrons are shared equally by the bonded atoms, resulting in a
balanced distribution of electrical charge
0-5% ionic character (0-0.3 EN difference) is nonpolar-covalent bond
11. Polar-covalent Bonds
Bonds that have significantly different Ens, electrons more strongly
attracted by more-EN atom
These bonds are polar they have an uneven distribution of charge
12.
Covalent bonds with 5-50% ionic character (0.3-1.7 EN difference) are
polar
δ+
δ−
Polar-covalent bond covalent bond in which the bonded atoms have
an unequal attraction for the shared electrons
13. Sample Problem
Use electronegativity differences to classify bonding
between sulfur, S, and the following elements:
hydrogen, H; cesium, Cs; and chlorine, Cl. In each
pair, which atom will be more negative?
15. Practice Problem
Use electronegativity differences to classify bonding
between chlorine, Cl, and the following elements:
calcium, Ca; oxygen, O; and bromine, Br. Indicate the
more-negative atom in each pair.
17. Section 2 – Covalent Bonding and
Molecular Compounds
18.
Many chemical compounds are molecules
Molecule neutral group of atoms that are held
together by covalent bonds
Single molecule of compound is individual unit
Capable of existing on its own
May consist of 2 or more atoms of same element or two
or more different atoms
Molecular compound chemical compound whose
simplest units are molecules
19.
20. Formation of Covalent Bond
Bonded atoms have lower potential energy than unbonded atoms
At large distance atoms don’t influence each other
Potential energy set at 0
21.
Each H has (+) proton
Nucleus surrounded by (-) electron
As atoms near each other, charged particles start to interact
22. Approaching nuclei and
electrons are attracted
to each other
Decrease in total
potential energy
At the same time, two nuclei and
two electrons repel each other
Increase in potential energy
23.
The amount of attraction/repulsion depends on how close the atoms are
to each other
When atoms first “see” each other, electron-proton attraction stronger
than e-e or p-p repulsions
So atoms drawn to each other and potential energy lowered
24.
Attractive force dominates until a distance is reached where repulsion
equals attraction
Valley of the curve
25.
Closer the atoms get, potential energy rises sharply
Repulsion becomes greater than attraction
26. Characteristics of the Covalent Bond
Bonded atoms vibrate a bit
As long as energy stays close to minimum they stay
covalently bonded
Bond length the distance between two bonded atoms
at their minimum potential energy (average distance
between two bonded atoms)
27.
To form covalent bond, hydrogen atoms need to release energy
Amount of energy equals difference between potential energy at zero
level (separated atoms) and at bottom of valley (bonded atoms)
28.
Same amount of energy must be added to separate bonded atoms
Bond energy energy required to break a chemical bond and form
neutral isolated atoms
29.
Units of bond energy usually kJ/mol
Indicates energy required to break one mole of bonds in isolated
molecules
Ex. 436 kJ/mol is energy needed to break H-H bonds in 1 mol hydrogen
molecules and form 2 mol of separated H atoms
Bond lengths and bond energies vary with the types of atoms that have
combined
30.
31.
All individual H atoms contain single, unpaired e- in 1s orbital
Sharing allows electrons to experience effect of stable electron
configuration of helium, 1s2
32. The Octet Rule
Noble-gas atoms have minimum energy existing on their own b/c of
electron configurations
Outer orbitals completely full
Other atoms full orbitals by sharing electrons
Bond formation follows octet rule chemical compounds tend to form
so that each atom, by gaining, losing, or sharing electrons, has an octet
of electrons in its highest occupied energy level
33. Example: Bonding of Fluorine
2 F atoms bond to form F2
7 e- in highest energy level
35. Exceptions to Octet Rule
Most main-group elements form covalent bonds
according to octet rule
Ex. H-H only 2 electrons
Boron, B, has 3 valence electrons ([He]2s22p1)
Boron tends to form bonds where it is surrounded by 6
e- (e- pairs)
Others can be surrounded by more than 8 when bonding
to highly electronegative elements
36. Electron Dot Notation
To keep track of valence electrons, it is helpful to use electron dot
notation electron-configuration notation in which only the valence
electrons of an atom of a particular element are shown, indicated by
dots placed around the element’s symbol
Inner-shell electrons NOT shown
37.
38. Sample Problem 1
Write the electron-dot notation for hydrogen.
A hydrogen atom has only one occupied energy level, the n=1 level, which
contains a single electron. So, e-dot notation is written as
H
39. Sample Problem 2
Write the e-dot notation for nitrogen.
Group notation for nitrogen’s family is ns2np3 which means nitrogen has 5
valence electrons. E-dot notation is written as
40. Lewis Structures
E-dot notation can also be used to represent molecules
Ex. H2 represented by combining notations of 2
individual H atoms
Pair of dots represents e- being shared
41. F2
Each F atom surrounded by 3 pairs e- that are not shared in bonds
Unshared (lone) pair pair of e- that is not involved in bonding and
that belongs completely to one atom
42. Lewis Structures
Pair of dots representing shared pair in covalent bond often replaced by
long dash
H H
H H
Lewis structures formulas in which atomic symbols represent nuclei
and inner-shell electrons, dot-pairs or dashes between two atomic
symbols represent electron pairs in covalent bonds, and dots next to only
one atomic symbol represent unshared electrons
43.
Common to write Lewis structures that show only
shared e- using dashes
Structural formula indicates the kind, number,
arrangement, and bonds but not the unshared pairs of
atoms in a molecule
F-F
H - Cl
44.
Lewis structures and structural formulas for many
molecules can be drawn if you know the composition of
the molecule and which atoms are bonded to each other
Single bond covalent bond made by sharing of one
pair of e- between 2 atoms
45. Sample Problem
Draw the Lewis structure of iodomethane, CH3I.
1. Determine type and number of atoms in molecule.
1 C, 1 I, 3 H
2. Write the e-dot notation for each type of atom in the molecule.
46. 3. Determine the total number of valence e- in the atoms
to be combined.
C
1 x 4e- = 4e-
I 1 x 7e- = 7eH
3 x 1e- = 3e14e-
47. 4. Arrange the atoms to form a skeleton structure for the molecule
If carbon is present, it is the central atom
Otherwise, the least-electronegative atom is central (except for hydrogen
which is NEVER central)
Then connect the atoms by electron-pair bonds.
48. 5. Add unshared pairs of electrons so that each hydrogen atom shares a pair
of electrons and each other nonmetal is surrounded by 8 electrons.
49. 6. Count the electrons in the structure to be sure that the
number of valence e- used equals the number
available. Be sure the central atom and other atoms
besides H have an octect.
There are eight e- in the four covalent bonds and six e- in
the three unshared pairs, giving the correct total of 14
valence electrons
52. Multiple Covalent Bonds
Atoms of same elements (especially C, N and O) can share more than one
e- pair
Double bond covalent bond made by the sharing of two pairs of ebetween two atoms
Shown by two side-by-side pairs of dots or two parallel dashes
53.
Triple bond covalent bond made by sharing of 3 pairs of e- between 2
atoms
Ex. N2
Each N has 5 valence
Each N shares 3 e- with other
54.
55.
Multiple bonds double and triple bonds
Double bonds have higher bond energies and are shorter
than single bonds
Triple bonds have higher bond energies and are shorter
than double bonds
60. Practice Problem
Draw the Lewis structure for hydrogen cyanide, which contains one
hydrogen atom, one carbon atom, and one nitrogen atom.
61. Resonance Structures
Some molecules/ions cannot be represented correctly by single Lewis
structure
Ex. Ozone (O3)
Each structure has one single and one double bond
62.
Chemists used to think ozone spends time alternating or
“resonating” between two structures
Now know that actual structure is something like an
average between the two
Resonance bonding in molecules or ions that cannot
be correctly represented by a single Lewis structure
65. Ionic Bonding
Ionic compound composed of positive and negative
ions that are combined so that the numbers of positive
and negative charges are equal
Most exist as crystalline solids, a 3-D network of (+) and
(-) ions mutually attracted to one another
66.
Different from molecular compound b/c ionic compound
not made of independent, neutral units
Chemical formula represents simplest ratio of
compound’s combined ions that give electrical
neutrality
67.
Chemical formula of ionic compound shows ratio of ions
present in ANY sample of ANY size
Formula unit simplest collection of atoms from
which an ionic compound’s formula can be recognized
Ex. NaCl is formula unit for sodium chloride
One sodium cation and one chlorine anion
68.
Ratio of ions in formula depends on charges of ions combined
Ex. Calcium and fluorine
Ca2+
F1- = total +1
So need 2 F1- to equal +2+(-2) = 0
Formula unit is CaF2
69. Formation of Ionic
Compoundsbe used to demonstrate changes that take place in
E-dot notation can
ionic bonding
Do not usually form by combination of isolated ions
71. Characteristics of Ionic
Bonding ions minimize potential energy by combining in orderly
In ionic crystals,
arrangement called a crystal lattice
72.
Attractive forces: between oppositely charged ions
(cations and anions) and between nuclei and electrons
Repulsive forces: between like-charged ions and
between electrons
Crystal lattice structure represents balance between
these two forces
73.
Within arrangement, each Na+ is surrounded by 6 Cl-
At the same time, each Cl- is surrounded by 6 Na+
74.
3-D arrangements of ions and strengths of attraction are
different with sizes and charges of ions and number of
ions of different charges
Ex. CaF2, there are 2 anions for each cation
Each Ca2+ is surrounded by 8 F-
Each F- is surrounded by 4 Ca2+
75. Lattice Energy
To compare bond strengths in ionic compounds,
chemists compare amounts of energy released when
separated ions in gas form crystalline solid
Lattice energy energy released when one mole of an
ionic compound is formed from gaseous ions
76. Comparison of Ionic and Molecular
Compounds
Force that holds ions together in ionic compounds is
very strong overall between opposite charges
Molecular compound – bonds making up each molecule
also strong, but forces between molecules not strong
77.
Because of bond strength difference, molecular
compounds melt at lower temperatures
Ionic compounds have higher melting and boiling points
78.
Ionic compounds are hard but brittle
Slight shift of one row of ions causes large buildup of
repulsive forces
Repulsive forces make layers split completely
79.
80.
In solid state ions cannot move – compounds are not electrical conductors
Molten state – ions can move freely and can carry electric current
Many ionic compounds dissolve in water
Attraction to water molecules overcomes attraction to each other
81. Polyatomic Ions
Certain atoms bond covalently to each other to form
group of atoms that has molecular AND ionic
characteristics
Polyatomic ion a charged group of covalently bonded
atoms
82. Lewis Structures of Polyatomic Ions
Polyatomic ions combine with ions of opposite charge to form ionic
compounds
To find Lewis structure, follow previous instructions except
If ion is negative, add to the total number of valence electrons a number
of e- same as the ions negative charge
If ion positive, subtract same number of e- as the positive charge
85. Metallic Bonding is Different
Metals have unique property of highly movable electrons (why they
conduct electricity so well)
In molecular compounds e- cannot move, held in shared bond
In ionic compounds, e- cannot move, held to individual ions
86. Metallic-Bond Model
Highest energy levels of most metal atoms only
occupied by few e-
Ex. s-block metals have one or two valence e- where all
3 p orbitals are empty
d-block metals have many empty d orbitals just below
highest energy level
87. Overlapping Orbitals
Within metal, empty orbitals in outer energy levels
overlap
Allows outer e- to move freely
e- are delocalized do not belong to any one atom
Metallic bonding chemical bonding that results from
attraction between metal atoms and surrounding sea of
electrons
88. Metallic Properties
Freedom of e- to move around causes high electrical and
thermal conductivity
b/c many orbitals separated by very small energy
differences, metals can absorb wide range of light
frequencies
Absorption of light excites e- to higher energy levels
e- immediately fall back down to lower levels, giving off
light (why metals are shiny)
89.
Malleability ability of a substance to be hammered
or beaten into thin sheets
Ductility ability of a substance to be pulled into
wires
Both possible because of structure, one line of metal
atoms can slide without breaking bonds
Not possible with ionic
structures
crystal
90. Metallic Bond Strength
Bond strength varies with nuclear charge of metal
atoms and number of e- in metal’s e- sea
Both factors reflected as heat of vaporization
When metal vaporized, bonded atoms in solid state
converted to individual atoms in gas state
Higher heat of vaporization, higher bond strength
92. Molecular Geometry
Properties of molecules depend on bonding of atoms
and the 3-Dimensional arrangement of molecule’s atoms
in space
Polarity of each bond, along with geometry of molecule,
determines molecular polarity uneven distribution
of molecular charge
Strongly influences forces that act BETWEEN molecules
93. VSEPR Theory
Diatomic molecules must be linear (only two atoms)
To predict geometries of more complex molecules,
consider locations of all e- pairs surrounding bonded
atoms
This is basis of VSEPR
94. “Valence-shell, electron-pair repulsion”
VSEPR theory repulsion between the sets of valencelevel e- surrounding an atom causes these set to be
oriented as far apart as possible
How does this account for molecular shape?
Let’s consider only molecules with no unshared valence
e- on central atom
95. Ex. BeF2
Be doesn’t follow octect rule
Be forms covalent bond with each F atom
Surrounded by only two electron pairs it shares with F
atoms
According to VSEPR, shared pairs oriented as far away
from each other as possible
96.
Distance between e- pairs maximized if bonds to F are
on opposite sides of Be, 180˚ apart
So, all 3 atoms lie in straight line – molecule is linear
97.
If we represent central atom in molecule by “A” and
atoms bonded to “A” are represented by “B” then BeF2
is an example of an AB2 molecule
AB2 is linear
What would AB3 look like?
98.
The 3 A-B bonds stay farthest apart by pointing to
corners of equilateral triangle, giving 120˚ between
bonds
= trigonal-planar geometry
99.
AB4 molecules following octect rule by sharing 4 e- pairs
with B atoms
Distance between e- pairs maximized if each A-B bond
points to one of 4 corners of tetrahedron (tetrahedral
geometry)
Angle is 109.5˚
101.
This molecule is an exception to the octet rule because
in this case Al forms only three bonds
Aluminum trichloride is an AB3 type of molecule
Therefore, according to VSEPR theory, it should have
trigonal-planar geometry
102. Practice Problem
Use VSEPR theory to predict the molecular geometry of the following
molecules:
a. HI
linear
b. CBr4
tetrahedral
c. AlBr3
Trigonal-planar
d. CH2Cl2
tetrahedral
103. VSEPR and Unshared e- Pairs
Ammonia, NH3, and water, H2O, are examples of
molecules where central atom has both shared and
unshared e- pairs
How does VSEPR account for the geometries?
104.
Lewis structure of ammonia shows in addition to 3 epairs it shares with 3 H atoms, the central N has one
unshared pair of e-
VSEPR theory says that lone pair occupies space around
N atom just as bonding pairs do
So, as an AB4 molecule, e- pairs maximize separation by
assuming 4 corners of tetrahedron
105.
Lone pairs occupy space but description of shape of
molecule refers to positions of atoms only
So, molecular geometry of ammonia molecule is
pyramid with triangular base
General formula is AB3E
E is unshared e- pair
106.
Water molecule has 2 unshared e- pairs
It is AB2E2 molecule
A (O) is at center of tetrahedron
2 corners occupied by B (H)
Other 2 corners occupied by E (unshared e-)
108. Molecular
Shape
Atoms bonded
to central
atom
Lone pairs of
electrons
Bond angle
Trigonalpyramidal
3
1
Less than
109.5˚
Bent or Angular
2
2
Less than
109.5˚
Trigonalbipyramidal
5
0
90, 120, and
80˚
Octahedral
6
0
90 and 180˚
110. Intermolecular Forces
Intermolecular forces forces of attraction between
molecules
Vary in strength
Generally weaker than bonds that hold molecules
together (covalent, ionic)
111. Molecular Polarity and
Dipole-Dipole Forces
Strongest intermolecular forces exist between polar
molecules
Polar molecules act as tiny dipoles b/c of uneven charge
distribution
Dipole created by equal but opposite charges that
are separated by a short distance
112.
Direction of dipole is from dipole’s positive pole to its negative pole
Represented by arrow with head pointing toward negative pole and
crossed tail pointing toward positive pole
H – Cl
Cl more electronegative, and so is negative end