Covalent bonding occurs when two nonmetal atoms share pairs of electrons to achieve stable octets. Atoms form covalent bonds by sharing electrons, depicted using Lewis dot structures. There are three main types of bonding: ionic, metallic, and covalent. Covalent bonding is the focus of this chapter and forms molecules by electron sharing between nonmetals.
The document summarizes the key differences between ionic and covalent bonding. Ionic bonds form when a metal transfers electrons to a nonmetal, creating oppositely charged ions. Covalent bonds form when nonmetals share electrons to obtain a full outer shell. Ionic compounds have high melting points, are brittle solids, and dissolve well in water, while covalent compounds have lower melting points, are soft and pliable, and are generally insoluble in water.
The document discusses electron configurations, which describe how electrons are distributed in atomic orbitals. It explains the Aufbau principle, which states that electrons fill lower energy orbitals first. The Pauli exclusion principle is described, stating that no more than two electrons can occupy any single orbital. Hund's rule is also covered, regarding the filling of degenerate orbitals. Examples are provided to illustrate these principles.
1. The document provides an overview of writing formulas and naming ionic and covalent compounds. It reviews the periodic table and properties of metals, nonmetals and metalloids.
2. Key concepts covered include ion formation, the octet rule, polyatomic ions, oxidation numbers, naming conventions for ionic compounds containing metals or transition metals, and prefixes used in naming covalent compounds.
3. The document distinguishes between ionic and covalent bonding, lattice structures, and molecular structures of compounds.
Chemical bonds form through different types of attractions between atoms. Ionic bonds form when electrons are transferred from one atom to another, creating oppositely charged ions that are attracted to each other. Covalent bonds form when atoms share electrons equally. Ionic bonds are generally stronger than covalent bonds because more energy is required to overcome the electrostatic forces between ions.
This document discusses molecular geometry and polarity using valence shell electron pair repulsion (VSEPR) theory and hybridization theory. It begins by explaining the objectives and defining molecular polarity in terms of bond polarity and molecular geometry. It then applies VSEPR theory to predict the shapes of different molecules and polyatomic ions. The document continues by introducing hybridization theory to explain bonding orbitals and different hybridizations like sp3, sp2, and sp. It concludes by describing various intermolecular forces like dipole-dipole forces, hydrogen bonding, London dispersion forces, and how they relate to boiling points.
This document provides information about electron configuration. It begins by defining electron configuration as the arrangement of electrons in an atom's orbitals, which is described using quantum numbers. It then discusses the three main rules for writing electron configurations: 1) Aufbau principle, which states that electrons fill the lowest available energy levels first, 2) Pauli exclusion principle, which limits each orbital to two electrons of opposite spin, and 3) Hund's rule, which states that degenerate orbitals will fill with one electron each before pairing. The document provides examples of writing full and condensed electron configurations and drawing orbital diagrams for various elements. It includes an activity for students to practice these skills.
1. The document discusses the formation of chemical bonds through ionic bonds and covalent bonds. Ionic bonds involve the transfer of electrons between metal and non-metal atoms to form ions, while covalent bonds involve the sharing of electron pairs between non-metal atoms.
2. Key differences between ionic and covalent bonds are outlined, such as ionic bonds forming between metal and non-metal atoms that become oppositely charged ions, while covalent bonds form between non-metal atoms that share electrons.
3. Examples of ionic compounds such as sodium chloride and covalent compounds such as methane are given to illustrate the different types of bonds.
Isotopes are atoms of the same element that have different numbers of neutrons, while ions are atoms or groups of atoms that have gained or lost electrons, giving them a positive or negative charge. Ions are formed by gaining or losing electrons, not protons. Cations are positively charged ions formed when atoms lose electrons, such as metals, while anions are negatively charged ions formed when nonmetals gain electrons. The periodic table can be used to predict which elements will form cations or anions based on their group.
The document summarizes the key differences between ionic and covalent bonding. Ionic bonds form when a metal transfers electrons to a nonmetal, creating oppositely charged ions. Covalent bonds form when nonmetals share electrons to obtain a full outer shell. Ionic compounds have high melting points, are brittle solids, and dissolve well in water, while covalent compounds have lower melting points, are soft and pliable, and are generally insoluble in water.
The document discusses electron configurations, which describe how electrons are distributed in atomic orbitals. It explains the Aufbau principle, which states that electrons fill lower energy orbitals first. The Pauli exclusion principle is described, stating that no more than two electrons can occupy any single orbital. Hund's rule is also covered, regarding the filling of degenerate orbitals. Examples are provided to illustrate these principles.
1. The document provides an overview of writing formulas and naming ionic and covalent compounds. It reviews the periodic table and properties of metals, nonmetals and metalloids.
2. Key concepts covered include ion formation, the octet rule, polyatomic ions, oxidation numbers, naming conventions for ionic compounds containing metals or transition metals, and prefixes used in naming covalent compounds.
3. The document distinguishes between ionic and covalent bonding, lattice structures, and molecular structures of compounds.
Chemical bonds form through different types of attractions between atoms. Ionic bonds form when electrons are transferred from one atom to another, creating oppositely charged ions that are attracted to each other. Covalent bonds form when atoms share electrons equally. Ionic bonds are generally stronger than covalent bonds because more energy is required to overcome the electrostatic forces between ions.
This document discusses molecular geometry and polarity using valence shell electron pair repulsion (VSEPR) theory and hybridization theory. It begins by explaining the objectives and defining molecular polarity in terms of bond polarity and molecular geometry. It then applies VSEPR theory to predict the shapes of different molecules and polyatomic ions. The document continues by introducing hybridization theory to explain bonding orbitals and different hybridizations like sp3, sp2, and sp. It concludes by describing various intermolecular forces like dipole-dipole forces, hydrogen bonding, London dispersion forces, and how they relate to boiling points.
This document provides information about electron configuration. It begins by defining electron configuration as the arrangement of electrons in an atom's orbitals, which is described using quantum numbers. It then discusses the three main rules for writing electron configurations: 1) Aufbau principle, which states that electrons fill the lowest available energy levels first, 2) Pauli exclusion principle, which limits each orbital to two electrons of opposite spin, and 3) Hund's rule, which states that degenerate orbitals will fill with one electron each before pairing. The document provides examples of writing full and condensed electron configurations and drawing orbital diagrams for various elements. It includes an activity for students to practice these skills.
1. The document discusses the formation of chemical bonds through ionic bonds and covalent bonds. Ionic bonds involve the transfer of electrons between metal and non-metal atoms to form ions, while covalent bonds involve the sharing of electron pairs between non-metal atoms.
2. Key differences between ionic and covalent bonds are outlined, such as ionic bonds forming between metal and non-metal atoms that become oppositely charged ions, while covalent bonds form between non-metal atoms that share electrons.
3. Examples of ionic compounds such as sodium chloride and covalent compounds such as methane are given to illustrate the different types of bonds.
Isotopes are atoms of the same element that have different numbers of neutrons, while ions are atoms or groups of atoms that have gained or lost electrons, giving them a positive or negative charge. Ions are formed by gaining or losing electrons, not protons. Cations are positively charged ions formed when atoms lose electrons, such as metals, while anions are negatively charged ions formed when nonmetals gain electrons. The periodic table can be used to predict which elements will form cations or anions based on their group.
This document discusses ionic and metallic bonding. It explains that ions are formed when atoms gain or lose electrons to achieve stable noble gas electron configurations. Metals form cations by losing electrons while nonmetals form anions by gaining electrons. Ionic compounds contain cations and anions in ratios represented by chemical formulas. Metallic bonding occurs via delocalized valence electrons that are shared between metal atoms.
The document discusses the electronic configuration of atoms, which is the arrangement of electrons in an atom's orbitals. It defines the key terms of energy levels and sublevels, which are the orbitals where electrons are arranged. Examples of electronic configurations are given for several elements, such as iodine and silicon. Rules for determining electronic configuration, such as Aufbau's principle, Pauli's exclusion principle, and Hund's rule are also outlined.
This document discusses electron configuration and the rules that define how electrons are arranged in an atom's orbitals. It explains:
1) There are three main rules that define electron configuration: the Aufbau principle, Pauli exclusion principle, and Hund's rule.
2) Higher energy levels can hold more electrons than lower energy levels because they are associated with larger volumes that can contain more orbitals.
3) Electron configuration can be represented using orbital diagrams with arrows or electron configuration notation using symbols and superscripts.
This document outlines the development of atomic theory from ancient Greek philosophers to modern physics. It describes key contributors such as Democritus, Dalton, Thomson, Rutherford, Bohr, and Schrodinger who proposed models of the atom based on experiments. The modern atomic theory is that atoms contain a nucleus of protons and neutrons surrounded by electrons in electron clouds, and that elements are defined by their atomic number while isotopes differ in neutron number. The four fundamental forces that act on atoms are also summarized.
The document discusses the determination of lattice energy of ionic compounds using the Born-Haber cycle. It explains that the lattice energy of sodium chloride can be determined experimentally by considering its formation through two different methods. Method 1 is the direct combination of solid sodium and gaseous chlorine to form solid sodium chloride. Method 2 involves 5 steps including sublimation, dissociation, ionization, and combining to form ions and the ionic solid. Using Hess's law, the lattice energy is calculated by equating the enthalpy change between the two methods. For sodium chloride, the calculated lattice energy is -773.95 kJ/mol.
The document provides information about ionic and covalent (molecular) bonding:
- Ionic bonds occur between metals and non-metals and involve the transfer of electrons. Covalent bonds occur between two non-metals and involve the sharing of electrons.
- Ionic compounds have high melting and boiling points and conduct electricity when melted or dissolved. Molecular compounds have lower melting and boiling points and do not conduct electricity.
- Ionic compounds exist as crystal lattices of ions, while molecular compounds exist as distinct molecules made of two or more nonmetal atoms bonded together.
The octet rule states that atoms are most stable when their outer electron shells contain 8 electrons. Atoms attain this stable electron configuration by gaining, losing, or sharing electrons with other atoms so that their outer shells match the noble gas configuration. The rule applies mainly to nonmetals like carbon, nitrogen, oxygen, and halogens, and also some metals such as sodium and magnesium.
The document discusses empirical and molecular formulas. An empirical formula shows the simplest whole number ratio of atoms in a compound, while a molecular formula shows the exact number of each atom. To calculate molecular formula from empirical formula: 1) make a table with elements, percentages, atomic masses, moles, and simplest ratios; 2) the empirical formula mass is the sum of the atoms' masses; 3) divide the molar mass by the empirical formula mass to get the common factor for the molecular formula. Two examples are given to calculate empirical and molecular formulas from percentage compositions and molar masses.
Ionic bonds form when electrons are transferred between metals and nonmetals, resulting in oppositely charged ions. Covalent bonds form when nonmetals share electrons in order to achieve a stable electron configuration. Ionic compounds have high melting points and conduct electricity well when melted but not solid, while covalent compounds have low melting points and generally do not conduct electricity or dissolve in water well.
The document summarizes key information about atomic structure:
- The nucleus is positively charged and contains nearly all an atom's mass, while electrons are much smaller and negatively charged, orbiting in shells outside the nucleus.
- Electrons are arranged in shells (also called energy levels) around the nucleus, with the first shell holding up to 2 electrons and subsequent shells holding up to 8 electrons each.
- Atoms can be represented using Bohr models that show the nucleus and electrons arranged in shells, with the number of protons and neutrons indicated in the nucleus.
This document provides information on naming hydrocarbon compounds according to IUPAC rules. It discusses the key aspects of hydrocarbon names including:
1) The suffixes that indicate type of hydrocarbon (-ane, -ene, -yne).
2) Using prefixes to indicate the number of carbons.
3) Numbering carbons and determining position of multiple bonds.
4) Naming side chains as alkyl groups and attaching them to the parent chain name alphabetically based on position.
An example demonstrates applying the rules to name a hydrocarbon with multiple side chains.
Electron configurations provide information about the location of electrons in an atom's orbitals. There are three main rules for determining electron configurations: 1) electrons fill orbitals starting with the lowest energy levels and moving upwards, 2) the Pauli exclusion principle states that no two electrons can occupy the same orbital with the same spin, and 3) Hund's rule states that electrons will occupy orbitals singly before pairing up. Writing out electron configurations involves determining the element's number of electrons and following the Aufbau principle of filling orbitals in order of increasing energy.
Ions are atoms that have gained or lost electrons, giving them an electrical charge. Atoms that lose electrons become positively charged cations, while atoms that gain electrons become negatively charged anions. The number of protons and neutrons remain the same in an ion, only the number of electrons changes. Ionic bonding occurs through the transfer of electrons from one atom to another, while covalent bonding involves the sharing of electron pairs between two atoms. Sodium chloride is an example of ionic bonding through electron transfer, while an oxygen molecule is held together by covalent bonding from shared electron pairs.
Metallic bonding occurs when metal atoms lose valence electrons to become positively charged ions embedded in a "sea" of delocalized electrons. This electron sea model explains several properties of metals, including their ability to conduct electricity and heat, as well as their malleability, ductility, lustrous appearance, and high melting and boiling points. The mobile electrons allow for heat and charge conduction, while the metallic lattice structure enables atoms to slide past one another under pressure.
This document discusses the four main types of chemical bonds: ionic bonds, covalent bonds, hydrogen bonds, and metallic bonds. Ionic bonds involve the transfer of electrons between atoms. Covalent bonds involve the sharing of electrons between two atoms. Hydrogen bonds are electrostatic attractions between hydrogen atoms covalently bonded to electronegative atoms and another electronegative atom. Metallic bonds are electrostatic attractions between positively charged metal ions and delocalized electrons in metals. Examples of each type of bond are provided.
The document discusses the mole concept in chemistry. It defines a mole as an amount of substance containing Avogadro's number of particles, which is 6.022x10^23 particles. It notes that a mole allows chemists to easily calculate quantities at the molecular scale rather than having to use actual numbers of atoms and molecules. The document provides examples of how to calculate moles using molar mass and mass quantities. It explains that the molar mass listed on the periodic table gives the mass in grams of one mole of that element or compound.
The document discusses Lewis dot structures, which use dots to represent valence electrons around an atomic symbol. It explains that ions have Lewis dot diagrams with fewer (for cations) or more (for anions) dots than the corresponding atom due to gaining or losing electrons. The document provides examples of Lewis dot diagrams for various ions, such as Ca2+ and O2-. It also includes practice problems asking students to draw Lewis dot diagrams for additional ions.
Covalent bonding occurs when two nonmetal atoms share pairs of electrons to acquire stable electron configurations. Atoms form covalent bonds by sharing electrons, drawing Lewis dot structures to show these electron pairs. There are three main types of bonding: ionic, metallic, and covalent. Covalent bonding is the focus of this chapter and forms molecules by electron sharing between nonmetals. Molecular compounds are named using prefixes to indicate the number of each element present. Bond polarity is determined using electronegativity values and can classify bonds and molecules as ionic, polar covalent, or nonpolar covalent.
This document discusses ionic and metallic bonding. It explains that ions are formed when atoms gain or lose electrons to achieve stable noble gas electron configurations. Metals form cations by losing electrons while nonmetals form anions by gaining electrons. Ionic compounds contain cations and anions in ratios represented by chemical formulas. Metallic bonding occurs via delocalized valence electrons that are shared between metal atoms.
The document discusses the electronic configuration of atoms, which is the arrangement of electrons in an atom's orbitals. It defines the key terms of energy levels and sublevels, which are the orbitals where electrons are arranged. Examples of electronic configurations are given for several elements, such as iodine and silicon. Rules for determining electronic configuration, such as Aufbau's principle, Pauli's exclusion principle, and Hund's rule are also outlined.
This document discusses electron configuration and the rules that define how electrons are arranged in an atom's orbitals. It explains:
1) There are three main rules that define electron configuration: the Aufbau principle, Pauli exclusion principle, and Hund's rule.
2) Higher energy levels can hold more electrons than lower energy levels because they are associated with larger volumes that can contain more orbitals.
3) Electron configuration can be represented using orbital diagrams with arrows or electron configuration notation using symbols and superscripts.
This document outlines the development of atomic theory from ancient Greek philosophers to modern physics. It describes key contributors such as Democritus, Dalton, Thomson, Rutherford, Bohr, and Schrodinger who proposed models of the atom based on experiments. The modern atomic theory is that atoms contain a nucleus of protons and neutrons surrounded by electrons in electron clouds, and that elements are defined by their atomic number while isotopes differ in neutron number. The four fundamental forces that act on atoms are also summarized.
The document discusses the determination of lattice energy of ionic compounds using the Born-Haber cycle. It explains that the lattice energy of sodium chloride can be determined experimentally by considering its formation through two different methods. Method 1 is the direct combination of solid sodium and gaseous chlorine to form solid sodium chloride. Method 2 involves 5 steps including sublimation, dissociation, ionization, and combining to form ions and the ionic solid. Using Hess's law, the lattice energy is calculated by equating the enthalpy change between the two methods. For sodium chloride, the calculated lattice energy is -773.95 kJ/mol.
The document provides information about ionic and covalent (molecular) bonding:
- Ionic bonds occur between metals and non-metals and involve the transfer of electrons. Covalent bonds occur between two non-metals and involve the sharing of electrons.
- Ionic compounds have high melting and boiling points and conduct electricity when melted or dissolved. Molecular compounds have lower melting and boiling points and do not conduct electricity.
- Ionic compounds exist as crystal lattices of ions, while molecular compounds exist as distinct molecules made of two or more nonmetal atoms bonded together.
The octet rule states that atoms are most stable when their outer electron shells contain 8 electrons. Atoms attain this stable electron configuration by gaining, losing, or sharing electrons with other atoms so that their outer shells match the noble gas configuration. The rule applies mainly to nonmetals like carbon, nitrogen, oxygen, and halogens, and also some metals such as sodium and magnesium.
The document discusses empirical and molecular formulas. An empirical formula shows the simplest whole number ratio of atoms in a compound, while a molecular formula shows the exact number of each atom. To calculate molecular formula from empirical formula: 1) make a table with elements, percentages, atomic masses, moles, and simplest ratios; 2) the empirical formula mass is the sum of the atoms' masses; 3) divide the molar mass by the empirical formula mass to get the common factor for the molecular formula. Two examples are given to calculate empirical and molecular formulas from percentage compositions and molar masses.
Ionic bonds form when electrons are transferred between metals and nonmetals, resulting in oppositely charged ions. Covalent bonds form when nonmetals share electrons in order to achieve a stable electron configuration. Ionic compounds have high melting points and conduct electricity well when melted but not solid, while covalent compounds have low melting points and generally do not conduct electricity or dissolve in water well.
The document summarizes key information about atomic structure:
- The nucleus is positively charged and contains nearly all an atom's mass, while electrons are much smaller and negatively charged, orbiting in shells outside the nucleus.
- Electrons are arranged in shells (also called energy levels) around the nucleus, with the first shell holding up to 2 electrons and subsequent shells holding up to 8 electrons each.
- Atoms can be represented using Bohr models that show the nucleus and electrons arranged in shells, with the number of protons and neutrons indicated in the nucleus.
This document provides information on naming hydrocarbon compounds according to IUPAC rules. It discusses the key aspects of hydrocarbon names including:
1) The suffixes that indicate type of hydrocarbon (-ane, -ene, -yne).
2) Using prefixes to indicate the number of carbons.
3) Numbering carbons and determining position of multiple bonds.
4) Naming side chains as alkyl groups and attaching them to the parent chain name alphabetically based on position.
An example demonstrates applying the rules to name a hydrocarbon with multiple side chains.
Electron configurations provide information about the location of electrons in an atom's orbitals. There are three main rules for determining electron configurations: 1) electrons fill orbitals starting with the lowest energy levels and moving upwards, 2) the Pauli exclusion principle states that no two electrons can occupy the same orbital with the same spin, and 3) Hund's rule states that electrons will occupy orbitals singly before pairing up. Writing out electron configurations involves determining the element's number of electrons and following the Aufbau principle of filling orbitals in order of increasing energy.
Ions are atoms that have gained or lost electrons, giving them an electrical charge. Atoms that lose electrons become positively charged cations, while atoms that gain electrons become negatively charged anions. The number of protons and neutrons remain the same in an ion, only the number of electrons changes. Ionic bonding occurs through the transfer of electrons from one atom to another, while covalent bonding involves the sharing of electron pairs between two atoms. Sodium chloride is an example of ionic bonding through electron transfer, while an oxygen molecule is held together by covalent bonding from shared electron pairs.
Metallic bonding occurs when metal atoms lose valence electrons to become positively charged ions embedded in a "sea" of delocalized electrons. This electron sea model explains several properties of metals, including their ability to conduct electricity and heat, as well as their malleability, ductility, lustrous appearance, and high melting and boiling points. The mobile electrons allow for heat and charge conduction, while the metallic lattice structure enables atoms to slide past one another under pressure.
This document discusses the four main types of chemical bonds: ionic bonds, covalent bonds, hydrogen bonds, and metallic bonds. Ionic bonds involve the transfer of electrons between atoms. Covalent bonds involve the sharing of electrons between two atoms. Hydrogen bonds are electrostatic attractions between hydrogen atoms covalently bonded to electronegative atoms and another electronegative atom. Metallic bonds are electrostatic attractions between positively charged metal ions and delocalized electrons in metals. Examples of each type of bond are provided.
The document discusses the mole concept in chemistry. It defines a mole as an amount of substance containing Avogadro's number of particles, which is 6.022x10^23 particles. It notes that a mole allows chemists to easily calculate quantities at the molecular scale rather than having to use actual numbers of atoms and molecules. The document provides examples of how to calculate moles using molar mass and mass quantities. It explains that the molar mass listed on the periodic table gives the mass in grams of one mole of that element or compound.
The document discusses Lewis dot structures, which use dots to represent valence electrons around an atomic symbol. It explains that ions have Lewis dot diagrams with fewer (for cations) or more (for anions) dots than the corresponding atom due to gaining or losing electrons. The document provides examples of Lewis dot diagrams for various ions, such as Ca2+ and O2-. It also includes practice problems asking students to draw Lewis dot diagrams for additional ions.
Covalent bonding occurs when two nonmetal atoms share pairs of electrons to acquire stable electron configurations. Atoms form covalent bonds by sharing electrons, drawing Lewis dot structures to show these electron pairs. There are three main types of bonding: ionic, metallic, and covalent. Covalent bonding is the focus of this chapter and forms molecules by electron sharing between nonmetals. Molecular compounds are named using prefixes to indicate the number of each element present. Bond polarity is determined using electronegativity values and can classify bonds and molecules as ionic, polar covalent, or nonpolar covalent.
Chemical bonds result from the attraction between positively charged nuclei and negatively charged electrons. There are three main types of bonding: ionic, metallic, and covalent. Ionic bonding involves the transfer of electrons between atoms to form ions with opposite charges that are attracted to each other. Covalent bonding involves the sharing of electrons between nonmetal atoms. Molecules form through covalent bonds and have molecular formulas that describe the number and type of atoms present. The shape and polarity of molecules can be predicted from their Lewis structures.
Intro chem ch 12 chemical bonding sp08khalidmohmed
Chemical bonds are attractive forces that hold groups of atoms together. There are two main types of bonds: ionic and covalent. Ionic bonds form between metals and nonmetals and involve the transfer of electrons. Covalent bonds form when atoms share electrons. The octet rule states that atoms seek to obtain a noble gas configuration by gaining, losing, or sharing electrons. Lewis structures can represent ionic and covalent bonding using dots and lines to show valence electrons. Molecular shape is determined by valence shell electron pair repulsion theory.
Chemical bonds are attractive forces that hold groups of atoms together. There are two main types of bonds: ionic and covalent. Ionic bonds form between metals and nonmetals via electrostatic attraction as electrons are transferred. Covalent bonds form when atoms share electrons. Lewis structures use dots and lines to represent valence electrons and show how atoms bond to achieve stable electron configurations like the octet rule.
Lewis structures show the bonding between atoms in a molecule using dots to represent valence electrons. The octet rule states that atoms are most stable when their valence shells are filled with eight electrons. Valence shell electron pair repulsion (VSEPR) theory predicts molecular geometry based on minimizing electron pair repulsion around a central atom. VSEPR identifies five basic molecular geometries (linear, trigonal planar, tetrahedral, trigonal bipyramidal, and octahedral) depending on the number of electron pairs around the central atom. Lone pairs influence molecular geometry differently than bonding pairs.
This document discusses the structure of matter and different types of chemical bonds. It begins by defining important terms like elements, atoms, and compounds. It then explains the three main types of atomic bonds: covalent, ionic, and metallic. Covalent bonds form when atoms share electrons, ionic bonds form when electrons are transferred between atoms to form ions, and metallic bonds result from the attraction between positively charged metal ions and delocalized electrons. The document provides examples of each bond type and discusses naming conventions for compounds. It also covers polyatomic ions and transition metals.
1. The document discusses covalent bonding and molecular compounds. It defines covalent bonds as the sharing of electrons between nonmetal atoms.
2. Molecular compounds are formed from covalent bonds between atoms. They have lower melting and boiling points than ionic compounds.
3. Molecular formulas show the number and type of atoms in a molecule, but not their arrangement. Water's molecular formula is H2O.
Covalent bonds form between nonmetal atoms by sharing valence electrons. Atoms share electrons to attain stable electron configurations like noble gases. Lewis structures show how valence electrons are arranged between bonded atoms. To draw Lewis structures, count the total valence electrons and distribute them to form single or double bonds between atoms until each atom has an octet of electrons. Examples of molecules held by covalent bonds are hydrogen, oxygen, and chlorine.
1) The document discusses molecular geometry and polarity. It defines molecular geometry as the three-dimensional arrangement of atoms in a molecule.
2) Lewis electron dot structures are used to represent valence electrons around atoms. The octet rule states that atoms gain, share, or transfer electrons to achieve a stable 8 electron configuration.
3) Valence shell electron pair repulsion theory predicts molecular shapes based on the number of bonding pairs and lone pairs around a central atom. Symmetric shapes are nonpolar while asymmetric shapes are polar due to unequal electron sharing.
This document provides an overview of covalent bonding and molecular compounds. It begins by defining covalent bonds as the sharing of electrons between nonmetals to form molecules. Molecular compounds are groups of atoms joined by covalent bonds. They typically have lower melting and boiling points than ionic compounds. The document then discusses how atoms form single, double and triple covalent bonds to achieve stable electron configurations through electron sharing. Examples are provided to illustrate how covalent bonds form in molecules like H2, NH3, HCN and CO2. The nature of coordinate covalent bonds is also explained. Finally, molecular orbital theory and VSEPR theory are introduced as models for describing covalent bonding at the molecular level.
This document provides an overview of chemical bonding concepts including:
- The octet rule which states that main group elements form ions to achieve 8 valence electrons.
- Ionic and covalent bonds are formed through the transfer or sharing of electrons respectively.
- Lewis structures are used to represent electron pairing in molecules and predict molecular geometry based on electron pair repulsion.
1. The document discusses different types of chemical formulas including molecular, empirical, and structural formulas. It provides examples of each type like H2O for water and C6H12 for hexene.
2. It also discusses ionic and covalent bonding. Ionic bonding involves the complete transfer of electrons from one atom to another, like from sodium to chlorine in NaCl. Covalent bonding involves the sharing of electron pairs between atoms.
3. The document describes electronegativity and how it relates to the polarity of covalent bonds. Polar covalent bonds form between atoms with an electronegativity difference of 0.5-1.6, while ionic bonds form between atoms with a difference above
This document discusses covalent compounds and their formation through shared electron pairs between nonmetals. It covers the octet rule for achieving stable electron configurations, different types of covalent bonds, and how to draw Lewis structures by arranging electrons around atoms. Exceptions to the octet rule are presented. Guidelines for naming covalent compounds from their formulas and writing formulas from names are also provided, along with examples.
The document discusses Lewis structures and covalent bonding. It provides steps for writing Lewis structures, including determining the molecular formula and connectivity, counting valence electrons, connecting atoms with bonds, adding electron pairs, checking for octets, and calculating formal charges. Constitutional isomers are described as isomers that differ in the order atoms are connected. Resonance structures are also discussed, where multiple Lewis structures can be written that differ in electron positions but have the same atomic positions.
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.
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 a summary of covalent bonding and molecular compounds in 3 paragraphs:
Covalent bonds result from the sharing of valence electrons between nonmetallic elements. Atoms joined by covalent bonds form molecules, the smallest units of a molecular substance. Molecules have a molecular formula showing the number and type of atoms, and may be represented by Lewis structures or structural formulas.
Multiple bonds can form when atoms share more than one pair of valence electrons. The octet rule describes how atoms bond to acquire a full outer shell of 8 electrons. Molecular shape is determined by VSEPR theory based on electron pair repulsion. Hybridization involves the mixing of atomic orbitals to form new hybrid orbitals for bonding
This document discusses different types of chemical bonds including ionic, covalent, and metallic bonds. It describes the formation of ionic bonds between metals and nonmetals and how ionization energy, electron affinity, and lattice energy contribute to the energetics of ionic bonding. Covalent bonding is explained as the sharing of electrons between nonmetals. Factors that determine bond polarity like electronegativity are also covered. The document provides details on writing Lewis structures, accounting for valence electrons and formal charges. Exceptions to the octet rule for molecules with odd numbers of electrons, incomplete octets, and expanded octets are explained.
Chemical bonding results from the attraction between nuclei and electrons. There are three main types of bonding: ionic, covalent, and metallic. Ionic bonding involves the transfer of electrons between atoms to form ions. Covalent bonding involves the sharing of electron pairs between atoms to form molecules. Metallic bonding occurs between metal atoms through delocalized valence electrons. The type of bonding determines the physical properties of the substance such as melting/boiling points and conductivity.
This document discusses the polarity of molecules. It defines electronegativity and the VSEPR theory for determining molecular geometry. A molecule's polarity depends on the polarity of its bonds, determined by electronegativity differences, and its geometry. Polar bonds result from unequal sharing of electron pairs between different elements. Molecular geometry is predicted using VSEPR theory to minimize electron pair repulsion. Both factors, bond polarity and geometry, must be considered to classify a molecule as polar or nonpolar. Examples are provided to demonstrate how to determine a molecule's polarity.
This document discusses atomic mass and related concepts. It defines atomic mass unit and discusses how to calculate average atomic mass of elements based on the natural abundances of isotopes. It also defines average molecular mass as the sum of average atomic masses of atoms in a molecule. Additionally, it defines average formula mass as the sum of atomic masses of atoms in an ionic compound. The document provides examples of calculations and practice problems related to these concepts.
There are several methods used to determine the age of rocks:
1. Radiometric dating examines the proportion of isotopes in rocks to determine their age based on known decay rates.
2. Relative dating studies the order of rock layers and fossils to determine the relative ages of rocks.
3. Absolute dating analyzes the chemical composition of rocks, including the ratio of uranium to lead, to calculate the precise numerical age based on known half-lives of radioactive isotopes.
Stratified rock layers are formed over long periods of time as sediments like sand, silt and clay settle in layers on the bottoms of bodies of water. These layers harden into rock as they are compressed by additional layers of sediment deposited over thousands or millions of years, forming identifiable rock strata. The document asks how stratified rock layers are formed and what stratification is.
Hazard maps show areas that are affected by or vulnerable to natural hazards such as earthquakes, volcanic eruptions, and landslides. These maps include a title to identify the area and hazard, legends and symbols to convey important information using lines and colors, a north arrow to indicate direction, and a scale to show the ratio between the map and the actual area depicted. Hazard maps are useful for understanding risks from natural disasters.
Earth is composed of many different types of rocks and minerals. Rocks are made up of one or more minerals, and together the minerals and rocks make up the solid part of our planet. Understanding the composition and properties of rocks and minerals can provide insight into Earth's formation and history.
Chapter wise All Notes of First year Basic Civil Engineering.pptxDenish Jangid
Chapter wise All Notes of First year Basic Civil Engineering
Syllabus
Chapter-1
Introduction to objective, scope and outcome the subject
Chapter 2
Introduction: Scope and Specialization of Civil Engineering, Role of civil Engineer in Society, Impact of infrastructural development on economy of country.
Chapter 3
Surveying: Object Principles & Types of Surveying; Site Plans, Plans & Maps; Scales & Unit of different Measurements.
Linear Measurements: Instruments used. Linear Measurement by Tape, Ranging out Survey Lines and overcoming Obstructions; Measurements on sloping ground; Tape corrections, conventional symbols. Angular Measurements: Instruments used; Introduction to Compass Surveying, Bearings and Longitude & Latitude of a Line, Introduction to total station.
Levelling: Instrument used Object of levelling, Methods of levelling in brief, and Contour maps.
Chapter 4
Buildings: Selection of site for Buildings, Layout of Building Plan, Types of buildings, Plinth area, carpet area, floor space index, Introduction to building byelaws, concept of sun light & ventilation. Components of Buildings & their functions, Basic concept of R.C.C., Introduction to types of foundation
Chapter 5
Transportation: Introduction to Transportation Engineering; Traffic and Road Safety: Types and Characteristics of Various Modes of Transportation; Various Road Traffic Signs, Causes of Accidents and Road Safety Measures.
Chapter 6
Environmental Engineering: Environmental Pollution, Environmental Acts and Regulations, Functional Concepts of Ecology, Basics of Species, Biodiversity, Ecosystem, Hydrological Cycle; Chemical Cycles: Carbon, Nitrogen & Phosphorus; Energy Flow in Ecosystems.
Water Pollution: Water Quality standards, Introduction to Treatment & Disposal of Waste Water. Reuse and Saving of Water, Rain Water Harvesting. Solid Waste Management: Classification of Solid Waste, Collection, Transportation and Disposal of Solid. Recycling of Solid Waste: Energy Recovery, Sanitary Landfill, On-Site Sanitation. Air & Noise Pollution: Primary and Secondary air pollutants, Harmful effects of Air Pollution, Control of Air Pollution. . Noise Pollution Harmful Effects of noise pollution, control of noise pollution, Global warming & Climate Change, Ozone depletion, Greenhouse effect
Text Books:
1. Palancharmy, Basic Civil Engineering, McGraw Hill publishers.
2. Satheesh Gopi, Basic Civil Engineering, Pearson Publishers.
3. Ketki Rangwala Dalal, Essentials of Civil Engineering, Charotar Publishing House.
4. BCP, Surveying volume 1
Communicating effectively and consistently with students can help them feel at ease during their learning experience and provide the instructor with a communication trail to track the course's progress. This workshop will take you through constructing an engaging course container to facilitate effective communication.
LAND USE LAND COVER AND NDVI OF MIRZAPUR DISTRICT, UPRAHUL
This Dissertation explores the particular circumstances of Mirzapur, a region located in the
core of India. Mirzapur, with its varied terrains and abundant biodiversity, offers an optimal
environment for investigating the changes in vegetation cover dynamics. Our study utilizes
advanced technologies such as GIS (Geographic Information Systems) and Remote sensing to
analyze the transformations that have taken place over the course of a decade.
The complex relationship between human activities and the environment has been the focus
of extensive research and worry. As the global community grapples with swift urbanization,
population expansion, and economic progress, the effects on natural ecosystems are becoming
more evident. A crucial element of this impact is the alteration of vegetation cover, which plays a
significant role in maintaining the ecological equilibrium of our planet.Land serves as the foundation for all human activities and provides the necessary materials for
these activities. As the most crucial natural resource, its utilization by humans results in different
'Land uses,' which are determined by both human activities and the physical characteristics of the
land.
The utilization of land is impacted by human needs and environmental factors. In countries
like India, rapid population growth and the emphasis on extensive resource exploitation can lead
to significant land degradation, adversely affecting the region's land cover.
Therefore, human intervention has significantly influenced land use patterns over many
centuries, evolving its structure over time and space. In the present era, these changes have
accelerated due to factors such as agriculture and urbanization. Information regarding land use and
cover is essential for various planning and management tasks related to the Earth's surface,
providing crucial environmental data for scientific, resource management, policy purposes, and
diverse human activities.
Accurate understanding of land use and cover is imperative for the development planning
of any area. Consequently, a wide range of professionals, including earth system scientists, land
and water managers, and urban planners, are interested in obtaining data on land use and cover
changes, conversion trends, and other related patterns. The spatial dimensions of land use and
cover support policymakers and scientists in making well-informed decisions, as alterations in
these patterns indicate shifts in economic and social conditions. Monitoring such changes with the
help of Advanced technologies like Remote Sensing and Geographic Information Systems is
crucial for coordinated efforts across different administrative levels. Advanced technologies like
Remote Sensing and Geographic Information Systems
9
Changes in vegetation cover refer to variations in the distribution, composition, and overall
structure of plant communities across different temporal and spatial scales. These changes can
occur natural.
How to Make a Field Mandatory in Odoo 17Celine George
In Odoo, making a field required can be done through both Python code and XML views. When you set the required attribute to True in Python code, it makes the field required across all views where it's used. Conversely, when you set the required attribute in XML views, it makes the field required only in the context of that particular view.
2. Chemical Bond
A Quick Review….
• A bond results from the attraction of nuclei
for electrons
– All atoms are trying to achieve a stable octet
• IN OTHER WORDS
– the protons (+) in one nucleus are attracted to
the electrons (-) of another atom
• This is Electronegativity !!
2
4. Three Major Types of Bonding
• Ionic Bonding
– forms ionic compounds
– transfer of valence e-
• Metallic Bonding
• Covalent Bonding
– forms molecules
– sharing of valence e-
– This is our focus this chapter
4
5. Ionic Bonding
• Always formed between metal cations
and non-metals anions
• The oppositely charged ions stick like
magnets
[METALS ]+ [NON-METALS ]-
Lost e-
Gained e-
5
7. Covalent Bonding
• Pairs of e- are
shared
between 2 non-
metal atoms to
acquire the electron
configuration of a
noble gas.
molecules
7
8. Covalent Bonding
• Occurs between nonmetal atoms which need to gain
electrons to get a stable octet of electrons or a filled
outer shell.
9. Drawing molecules (covalent)
using Lewis Dot Structures
• Symbol represents the KERNEL of the atom (nucleus and inner
electrons)
• dots represent valence electrons
• The ones place of the group number indicates the number of
valence electrons on an atom.
• Draw a valence electron on each side (top, right, bottom, left)
before pairing them.
9
10. Always remember atoms are trying
to complete their valence shell!
“2 will do but 8 is great!”
The number of electrons the atoms needs is the
total number of bonds they can make.
Ex. … H? O? F? N? Cl? C?
one two one three one four
10
11. Draw Lewis Dot Structures
You may represent valence electrons
from different atoms with the
following symbols x, ,
H or H or H
x
11
12. Covalent bonding
• The atoms form a covalent bond by
sharing their valence electrons to get a
stable octet of electrons.(filled valence
shell of 8 electrons)
• Electron-Dot Diagrams of the atoms are
combined to show the covalent bonds
• Covalently bonded atoms form
MOLECULES
13. Methane CH4
• This is the finished Lewis dot structure
• Every atom has a filled valence shell
How did we get here?
OR
13
14. General Rules for Drawing Lewis Structures
• All valence electrons of the atoms in Lewis structures must
be shown.
• Generally each atom needs eight electrons in its valence
shell (except Hydrogen needs only two electrons and
Boron needs only 6).
• Multiple bonds (double and triple bonds) can be formed by
C, N, O, P, and S.
• Central atoms have the most unpaired electrons.
• Terminal atoms have the fewest unpaired electrons.
14
15. • When carbon is one of you atoms, it will
always be in the center
• Sometimes you only have two atoms, so
there is no central atom
Cl2 HBr H2 O2 N2 HCl
• We will use a method called ANS
(Available, Needed, Shared) to help us draw
our Lewis dot structures for molecules
15
16. EXAMPLE 1: Write the Lewis structure for H2O where oxygen is the central atom.
Step 1: Determine the total number of electrons available for bonding. Because only valence
electrons are involved in bonding we need to determine the total number of valence electrons.
AVAILABLE valenceelectrons:
Electrons available
2 H Group 1 2(1) = 2
O Group 6 6
8
There are 8 electrons available for bonding.
Step 2: Determine the number of electrons needed by
each atom to fill its valence shell.
NEEDED valence electrons
Electrons needed
2 H each H needs 2 2(2) = 4
O needs 8 8
12
There are 12 electrons needed.
16
17. Step 3: More electrons are needed then there are available. Atoms therefore make bonds by sharing
electrons. Two electrons are shared per bond.
SHARED (two electrons per bond)
# of bonds = (# of electrons needed – # of electrons available) = (N-A) = (12 – 8) = 2 bonds.
2 2 2
Draw Oxygen as the central atom. Draw the Hydrogen atoms on either side of the oxygen atom.
Draw the 2 bonds that can be formed to connect the atoms.
OR
Step 4: Use remaining available electrons to fill valence shells for each atom. All atoms need 8 electrons
to fill their valence shell (except hydrogen needs only 2 electrons to fill its valence shell, and
boron only needs 6). For H2O there are 2 bonds, and 2 electrons per bond.
# available electrons remaining = # electrons available – # electrons shared = A-S = 8 – 2(2) = 4 extra e-
s
17
18. Sometimes multiple bonds must be formed to get
the numbers of electrons to work out
• DOUBLE bond
– atoms that share two e- pairs (4 e-)
O O
• TRIPLE bond
– atoms that share three e- pairs (6 e-)
N N 18
20. Step 3: SHARED (two electrons per bond)
# of bonds = (N – A) = (20 – 12) = 4 bonds.
2 2
Draw carbon as the central atom (Hint: carbon is always the center when it is present!). Draw the
Hydrogen atoms and oxygen atom around the carbon atom. Draw 2 bonds of the 4 bonds that can
be formed to connect the H atoms. Draw the remaining 2 bonds to connect the O atom (oxygen
can form double bonds)
Step 4: Use remaining available electrons to fill valence shell for each atom.
# electrons remaining = Available – Shared = A – S = 12 – 4(2) = 4 extra e-
s
20
21. Let’s Practice
H2
A = 1 x 2 = 2
N = 2 x 2 = 4
S = 4 - 2= 2 ÷ 2 = 1 bond
Remaining = A – S = 2 – 2 = 0
DRAW
21
22. Let’s Practice
CH4
A = C 4x1 = 4 H 1x4 = 4 4 + 4 = 8
N = C 8x1 = 8 H 2x4 = 8 8 + 8 = 16
S = (A-N)16 – 8 = 8 ÷2 = 4 bonds
Remaining = A-S = 8 – 8 = 0
DRAW
22
23. Let’s Practice
NH3
A = N 5x1 = 5 H 1x3 = 3 = 8
N = N 8x1 = 8 H 2x3 = 6 = 14
S = 14-8 = 6 ÷2 = 3 bonds
Remaining = (A-S) 8 – 6 = 2
DRAW
23
24. Let’s Practice
CO2
A = C 4x1 = 4 O 6x2 = 12 = 16
N = C 8x1 = 8 O 8x2 = 16 = 24
S = 24-16 = 8 ÷ 2 = 4 bonds
Remaining = (A-S) 16 – 8 = 8 not bonding
DRAW – carbon is the central atom
24
25. Let’s Practice
BCl3 boron only needs 6 valence electrons, it is an exception.
A = B 3 x 1 = 3 Cl 7 x 3 = 21 = 24
N = B(6) x 1 = 6 Cl 8 x 3 = 24 = 30
S = 30-24 = 6 ÷ 2 = 3 bonds
Remaining = 24 – 6 = 18 e- not bonding
DRAW
25
32. Rules for Naming
Molecular compounds
• The most “metallic” nonmetal
element is written first (the one that
is furthest left)
• The most nonmetallic of the two
nonmetals is written last in the
formula
• NO2 not O2N
• All binary molecular compounds end
in -ide
33. • Ionic compounds use charges to determine the
chemical formula
• The molecular compound‘s name tells you the
number of each element in the chemical
formula.
• Uses prefixes to tell you the quantity of each
element.
• You need to memorize the prefixes !
Molecular compounds
35. • If there is only one of the first element do
not put (prefix) mono
• Example: carbon monoxide (not monocarbon monoxide)
• If the nonmetal starts with a vowel, drop
the vowel ending from all prefixes except
di and tri
• monoxide not monooxide
• tetroxide not tetraoxide
More Molecular Compound Rules
60. Use Electronegativity Values to
Determine Bond Types
• Ionic bonds
– Electronegativity (EN) difference > 2.0
• Polar Covalent bonds
– EN difference is between .21 and 1.99
• Non-Polar Covalent bonds
– EN difference is < .20
– Electrons shared evenly in the bond
60
61. Ionic Character
“Ionic Character” refers to a bond’s
polarity
–In a polar covalent bond,
•the closer the EN difference is to 2.0,
the more POLAR its character
•The closer the EN difference is to .20,
the more NON-POLAR its character
61
62. Place these molecules in order of increasing
bond polarity using the electronegativity
values on your periodic table
• HCl
• CH4
• CO2
• NH3
• N2
• HF
a.k.a.
“ionic character”
62
1 EN difference = 0
2 EN difference = 0.4
3 EN difference = 0.9
4 EN difference = 1.0
3 EN difference = 0.9
5 EN difference = 1.9
63. Polar vs. Nonpolar
MOLECULES
• Sometimes the bonds within a
molecule are polar and yet the
molecule itself is non-polar
63
64. Nonpolar Molecules
• Molecule is Equal on all sides
–Symmetrical shape of molecule
(atoms surrounding central atom are
the same on all sides)
H
H
H
H C
Draw Lewis dot first and
see if equal on all sides
64
65. Polar Molecules
• Molecule is Not Equal on all sides
–Not a symmetrical shape of molecule
(atoms surrounding central atom are
not the same on all sides)
Cl
H
H
H C
65
70. H
H
O
Water is a POLAR molecule
ANY time there are unshared pairs
of electrons on the central atom, the
molecule is POLAR
70
71. Making sense of the polar
non-polar thing
BONDS
Non-polar Polar
EN difference EN difference
0 - .2 .21 – 1.99
MOLECULES
Non-polar Polar
Symmetrical Asymmetrical
OR
Unshared e-s on
Central Atom
71
72. 5 Shapes of Molecules
you must know!
(memorize)
72
73. Copy this slide
• VSEPR – Valence Shell Electron Pair
Repulsion Theory
– Covalent molecules assume geometry
that minimizes repulsion among electrons
in valence shell of atom
– Shape of a molecule can be predicted
from its Lewis Structure
73
74. 1. Linear (straight line)
Ball and stick
model
Molecule geometry X A X
OR
A X
Shared Pairs = 2 Unshared Pairs = 0
OR
74
75. 2. Trigonal Planar
Ball and stick
model
Molecule geometry X
A
X X
Shared Pairs = 3 Unshared Pairs = 0 75
79. • I can describe the 3 intermolecular
forces of covalent compounds and
explain the effects of each force.
79
80. • Attractions
within or inside
molecules, also
known as bonds.
– Ionic
– Covalent
– metallic
Intramolecular attractions
80
Roads within a state
81. • Attractions between
molecules
– Hydrogen “bonding”
• Strong attraction
between special polar
molecules (F, O, N, P)
– Dipole-Dipole
• Result of polar covalent
Bonds
– Induced Dipole
(Dispersion Forces)
• Result of non-polar
covalent bonds
Intermolecular attractions
81
82. More on intermolecular forces
Hydrogen “Bonding”
• STRONG
intermolecular force
– Like magnets
• Occurs ONLY
between H of one
molecule and N, O,
F of another
molecule
Hydrogen
“bond”
-
+
+
-
+ +
+
+
-
82
Hydrogen bonding
1 min
83. Why does Hydrogen
“bonding” occur?
• Nitrogen, Oxygen and Fluorine
– are small atoms with strong nuclear
charges
• powerful atoms
– Have very high electronegativities,
these atoms hog the electrons in a bond
– Create very POLAR molecules
83
84. Dipole-Dipole Interactions
– WEAK intermolecular force
– Bonds have high EN differences
forming polar covalent molecules,
but not as high as those that result
in hydrogen bonding.
.21<EN<1.99
– Partial negative and partial
positive charges slightly attracted
to each other.
– Only occur between polar
covalent molecules
84
86. Induced Dipole Attractions
– VERY WEAK intermolecular force
– Bonds have low EN differences EN < .20
– Temporary partial negative or positive charge
results from a nearby polar covalent molecule.
– Only occur between NON-POLAR & POLAR
molecules
86
Induced dipole video
30 sec
88. Intermolecular Forces
affect chemical properties
• For example, strong intermolecular
forces cause high Boiling Point
– Water has a high boiling point compared
to many other liquids
88
90. Which substance has the
highest boiling point?
• HF
• NH3
• CO2
• WHY?
The H-F bond has the highest
electronegativity difference
SO
HF has the most polar bond
resulting in the strongest H
bonding (and therefore needs the
most energy to overcome the
intermolecular forces and boil)
90