The document discusses different types of chemical bonding including ionic bonding, covalent bonding, and dative covalent bonding. Ionic bonding involves electron transfer between metals and nonmetals to form ionic compounds with electrostatic attraction between cations and anions. Covalent bonding involves sharing of electron pairs between atoms. Sigma and pi bonds are discussed as well as dot-and-cross diagrams and Lewis structures. Dative covalent bonding occurs when both electrons in a covalent bond come from the same atom, such as in a coordinate bond where a donor atom provides both electrons.
This document provides information about variants of question papers used for Cambridge International Examinations (CIE) assessments. It explains that for some popular exams, CIE uses two closely related variants of the question paper to maintain best assessment practices. Both variants assess the same content and skills to the same standard. Accompanying materials like mark schemes and examiner reports are also available in two variants. Schools will use only one variant for a given exam session. This provides students and teachers access to more past paper material for preparation. The document outlines the relationship between the question paper variants and corresponding marking and examiner materials. It also provides contact information for any questions about these changes.
1. The document discusses moles, molar mass, and calculations involving moles. It provides examples of calculating the number of particles, ions, or molecules in 1 mole of various substances like iron, carbon dioxide, sodium chloride, and magnesium chloride.
2. It also explains how to calculate molar mass using relative atomic mass and gives examples for iron, carbon dioxide, sodium chloride, and magnesium chloride.
3. Key concepts discussed include the definition of 1 mole as 6.02 x 1023 particles and its relationship to the Avogadro constant, and how molar mass is used to express the mass of 1 mole of a substance.
The document summarizes key concepts about transition elements including:
1. Transition elements are defined as elements that can form at least one stable ion with a partially filled d subshell. Common errors in defining transition elements are also discussed.
2. The electronic configurations of transition elements and their ions are covered, noting exceptions like Cr and Cu.
3. Physical properties of transition elements discussed include atomic/ionic radii, ionization energy, melting points and how they vary within the period and group.
The document discusses various theories of chemical bonding including valence bond theory, hybridization of atomic orbitals, molecular orbital theory, and bonding in different types of molecules. It explains how hybrid orbitals such as sp, sp2, sp3, sp3d, and sp3d2 are formed by mixing atomic orbitals. It also describes how sigma and pi bonds are formed through the overlap of hybrid and atomic orbitals. Molecular orbital theory is introduced as an alternative approach that considers the orbitals of whole molecules rather than individual atoms.
1) Ozone (O3) has a bond order of 1.5, with bond lengths intermediate between single and double bonds. Its structure is a resonance hybrid of two contributing resonance structures.
2) Ozone in the stratosphere absorbs harmful UV-B and UV-C radiation from the sun, protecting life on Earth. However, ozone in the troposphere is a air pollutant.
3) Chlorofluorocarbons (CFCs) and other ozone-depleting substances (ODS) can reach the stratosphere where they are broken down by UV radiation, releasing chlorine radicals that catalytically destroy ozone via chain reactions.
Haloalkanes and haloarenes notes by rawat sirRawat DA Greatt
This document contains notes on haloalkanes and haloarenes. It discusses the nature of the C-X bond in alkyl halides and how haloalkanes and haloarenes are prepared through various reactions like Sandmeyer's reaction, Wurtz reaction, and electrophilic aromatic substitution. The physical properties of haloalkanes and haloarenes are described, including their solubility, density, polarity, and boiling points. The document also outlines several chemical properties and reactions of haloalkanes like nucleophilic substitution, elimination, and reactions with metals. Important reactions for conversions of functional groups are also listed.
This document provides information about variants of question papers used for Cambridge International Examinations (CIE) assessments. It explains that for some popular exams, CIE uses two closely related variants of the question paper to maintain best assessment practices. Both variants assess the same content and skills to the same standard. Accompanying materials like mark schemes and examiner reports are also available in two variants. Schools will use only one variant for a given exam session. This provides students and teachers access to more past paper material for preparation. The document outlines the relationship between the question paper variants and corresponding marking and examiner materials. It also provides contact information for any questions about these changes.
1. The document discusses moles, molar mass, and calculations involving moles. It provides examples of calculating the number of particles, ions, or molecules in 1 mole of various substances like iron, carbon dioxide, sodium chloride, and magnesium chloride.
2. It also explains how to calculate molar mass using relative atomic mass and gives examples for iron, carbon dioxide, sodium chloride, and magnesium chloride.
3. Key concepts discussed include the definition of 1 mole as 6.02 x 1023 particles and its relationship to the Avogadro constant, and how molar mass is used to express the mass of 1 mole of a substance.
The document summarizes key concepts about transition elements including:
1. Transition elements are defined as elements that can form at least one stable ion with a partially filled d subshell. Common errors in defining transition elements are also discussed.
2. The electronic configurations of transition elements and their ions are covered, noting exceptions like Cr and Cu.
3. Physical properties of transition elements discussed include atomic/ionic radii, ionization energy, melting points and how they vary within the period and group.
The document discusses various theories of chemical bonding including valence bond theory, hybridization of atomic orbitals, molecular orbital theory, and bonding in different types of molecules. It explains how hybrid orbitals such as sp, sp2, sp3, sp3d, and sp3d2 are formed by mixing atomic orbitals. It also describes how sigma and pi bonds are formed through the overlap of hybrid and atomic orbitals. Molecular orbital theory is introduced as an alternative approach that considers the orbitals of whole molecules rather than individual atoms.
1) Ozone (O3) has a bond order of 1.5, with bond lengths intermediate between single and double bonds. Its structure is a resonance hybrid of two contributing resonance structures.
2) Ozone in the stratosphere absorbs harmful UV-B and UV-C radiation from the sun, protecting life on Earth. However, ozone in the troposphere is a air pollutant.
3) Chlorofluorocarbons (CFCs) and other ozone-depleting substances (ODS) can reach the stratosphere where they are broken down by UV radiation, releasing chlorine radicals that catalytically destroy ozone via chain reactions.
Haloalkanes and haloarenes notes by rawat sirRawat DA Greatt
This document contains notes on haloalkanes and haloarenes. It discusses the nature of the C-X bond in alkyl halides and how haloalkanes and haloarenes are prepared through various reactions like Sandmeyer's reaction, Wurtz reaction, and electrophilic aromatic substitution. The physical properties of haloalkanes and haloarenes are described, including their solubility, density, polarity, and boiling points. The document also outlines several chemical properties and reactions of haloalkanes like nucleophilic substitution, elimination, and reactions with metals. Important reactions for conversions of functional groups are also listed.
The document discusses the formation of ionic and covalent bonds, including how to determine the charge of ions, represent ionic compounds using dot-cross diagrams, and represent covalent molecules using dot-cross diagrams, structural formulas, and Lewis structures to show how atoms share electrons to achieve stable configurations. It also discusses exceptions where central atoms in some covalent compounds have fewer than 8 valence electrons.
This document presents information about chemical bonding from the chemistry project of students Akarshik Banerjee, Pratyush Dey, and Sayantan Biswas. It discusses various types of chemical bonds including covalent bonds, which form when atoms share electron pairs, and ionic bonds, which form through complete electron transfer. It also describes concepts like the octet rule, Lewis dot structures, formal charge, resonance structures, and molecular geometry based on valence shell electron pair repulsion theory. Hybridization and molecular orbitals are explained as well. In summary, the document provides an overview of key concepts in chemical bonding from the perspective of high school chemistry students.
This document contains a 50-question chemistry semester review covering various topics including:
1) Properties of ionic compounds including their high melting points.
2) Bond types and strengths in molecules like NH3.
3) Types of bonds that would form between different sets of elements.
4) Factors that determine states of matter for different compounds at various temperatures including intermolecular forces.
The review covers concepts of bonding, molecular shapes, properties of states of matter, acid/base chemistry and chemical reactions.
This document provides an introduction to organic chemistry, covering topics such as:
- The definition of organic chemistry as the study of carbon compounds.
- Electronic structure of atoms and how they bond through ionic and covalent bonding.
- Resonance structures and how they are used to represent molecules.
- Factors that influence acidity such as electronegativity, size, and resonance.
- The definitions of nucleophiles and electrophiles and their roles in bond formation.
This document summarizes several theories and concepts related to stereochemistry in main group compounds, including:
1) VSEPR theory which describes the geometry of molecules based on electron pairs around the central atom. It explains linear, trigonal, tetrahedral, trigonal bipyramidal, and octahedral geometries.
2) Bent's rule which describes how atomic s-character concentrates in orbitals directed toward electropositive substituents.
3) Walsh diagrams which use molecular orbital energies to determine molecular geometry.
4) Berry pseudorotation and fluxionality concepts which explain rapid ligand exchange in molecules like PF5.
This document discusses atomic structure and properties. It begins by reviewing the basic atomic model including protons, neutrons, and electrons. It then discusses atomic number, mass number, isotopes, ions, and electron configuration. The document also covers average atomic mass and mass spectrometry. It concludes by discussing the Bohr model of the atom and how this helped explain emission spectra of elements. In summary, the key topics covered are the fundamental particles of atoms, atomic notation, isotopes, ions, electron configuration, and early atomic structure models.
Dr. B. R. Thorat teaches chemistry at Ismail Yusuf College in Maharashtra, India. The document discusses orbital overlapping and conditions required for covalent bond formation. It also explains concepts of atomic orbitals, hybridization, resonance structures, and formal charges through examples of diatomic and polyatomic molecules. Valence bond theory alone cannot explain bonding in molecules like methane, requiring the additional concept of orbital hybridization.
1. Hydrogen bonding occurs between hydrogen atoms attached to electronegative atoms like oxygen, fluorine, and nitrogen of one molecule and an electronegative atom of another molecule.
2. Water is able to form extensive hydrogen bonding networks between molecules due to each water molecule having two hydrogen atoms and two lone pairs of electrons on the oxygen atom.
3. The hydrogen bonding network in liquid water is responsible for its unique properties, while hydrogen bonding in ice forms its crystalline lattice structure.
IB Chemistry on Redox, Oxidation states and Oxidation numberLawrence kok
This document provides information on oxidation states and oxidation numbers of elements. It lists common oxidation states of metals in groups 1-3 and transition metals. It also lists common oxidation states of nonmetals in groups 5-7. Rules are provided for assigning oxidation numbers based on the element and its bonds. Oxidation numbers are assigned by assuming ionic bonding and imagining electron flow to the more electronegative element. Exceptions and examples are given. Polyatomic ions and molecules are discussed.
This document discusses plane waves. Plane waves are waves whose wavefronts are infinite parallel planes. They have equal amplitude in two dimensions. Some key points:
- Plane waves transfer energy from one point to another without transferring matter. The transfer of energy is called wave motion.
- Examples of plane waves include electromagnetic waves in the far field region of an antenna and water waves that appear planar from a distance.
- The basic equations for an electromagnetic plane wave show that the electric and magnetic fields oscillate sinusoidally as they propagate in the direction of the wave vector.
- Waves can be one-dimensional, two-dimensional, or three-dimensional depending on whether the object has length only,
IB Chemistry on Born Haber Cycle and Lattice EnthalpyLawrence kok
The document describes the Born-Haber cycle process for determining the lattice enthalpy of various ionic compounds such as LiCl, NaCl, KCl, NaBr, NaF, and NaH. The Born-Haber cycle involves breaking down the ionic compound into gaseous ions and then measuring the standard enthalpy changes associated with steps such as atomization, ionization, and electron affinity which are used to calculate the unknown lattice enthalpy value. Diagrams of the Born-Haber cycle are provided for each example compound.
1. The document discusses experimental methods to verify the validity of the Debye-Hückel and Onsager equations, which describe electrolyte behavior at high dilutions.
2. Using liquid membrane cells, studies were able to examine dilutions as high as 10-4 M and found that the Debye-Hückel equation holds for 1:1 electrolytes but shows negative deviations for 2:2 or 3:2 electrolytes.
3. Conductance measurements of various electrolytes at different temperatures found close agreement with the Onsager equation at low concentrations, validating it as accurately describing conductance behavior at high dilutions.
Includes a discussion of Voltaic and electrolytic cells, the Nernst equation and the relationship between electrochemical processes, chemical equilibrium and free energy.
**More good stuff available at:
www.wsautter.com
and
http://www.youtube.com/results?search_query=wnsautter&aq=f
This document describes different types of covalent bonding, including:
1) Covalent bonds form when electron pairs are shared between two atoms, attracting their nuclei together.
2) Covalent compounds are usually composed of nonmetals and include many familiar substances like water and methane.
3) Dot and cross diagrams can represent covalent bonds, showing shared electron pairs between atoms.
This document is the first page of a physics exam from the University of Cambridge. It provides instructions to candidates regarding writing their information on the exam, using appropriate materials, and answering all questions. It includes a list of relevant physics formulas, constants, and units that may be used for reference. The exam consists of 15 printed pages and 1 blank page.
T21 IB Chemistry- Spectroscopy continued Robert Hughes
This document outlines lessons for a spectroscopy course, including:
1. A review of infrared spectroscopy, mass spectroscopy, and proton NMR spectroscopy.
2. Practice using these techniques to answer exam questions and deduce organic compound structures.
3. Details on high resolution proton NMR, including spin-spin coupling and interpreting splitting patterns.
4. Interpreting NMR spectra to determine functional groups and hydrogen environments.
5. A brief overview of X-ray crystallography and some practice questions reviewing all three techniques.
The course aims to build students' skills in applying spectroscopy methods to organic compound analysis through lessons, examples, and practice questions.
The document discusses the use of variant question papers by CIE (Cambridge International Examinations) for some of their popular assessments. It explains that while the content and standard of assessment remains the same, there are now two variants of the question papers, mark schemes, and principal examiner reports available for one component. This provides centers with more past examination material than usual to access. It includes both variants in the document. It also provides contact information for any questions about these changes and instructions for identifying the relevant parts of the document.
IB Chemistry on Redox, Oxidizing, Reducing Agents and writing half redox equa...Lawrence kok
The document discusses oxidation numbers (also called oxidation states), which are used to keep track of electrons in chemical reactions. Some key points:
- Oxidation numbers are assigned to each atom in a chemical species by assuming ionic bonding and counting electrons.
- Common rules are outlined for assigning oxidation numbers to elements, such as metals in Group 1 have a +1 oxidation state and nonmetals in Group 7 have a -1 oxidation state.
- Oxidation numbers can be used to determine if a reaction is a redox reaction by looking for changes in oxidation numbers between reactants and products.
- Transition metals can have multiple common oxidation states. Roman numerals are used to distinguish, such as
This document is a chemistry exam paper that consists of 7 questions testing knowledge of various chemistry concepts. It provides instructions for candidates on how to answer the questions, including writing in boxes provided and not writing in barred code or grey areas. It also lists the number of marks allocated for each question or part of a question. Finally, it states that the exam paper is made up of 14 printed pages and 2 blank pages.
csonn t1 atoms, molecules and stoichiometrycheeshengonn
Here are the key steps to solve this problem:
1) Isotope 35Cl has a relative abundance of 75.8%
2) Isotope 37Cl has a relative abundance of 24.2%
3) The relative atomic mass of 35Cl is 35 amu
4) The relative atomic mass of 37Cl is 37 amu
5) Use the formula: Relative Atomic Mass = Σ (Relative Abundance of isotope x Atomic Mass of isotope)
6) For 35Cl: Relative Abundance = 75.8%, Atomic Mass = 35 amu. So, contribution is 75.8% of 35 = 26.43
7) For 37Cl: Relative Abundance = 24
The document summarizes Dalton's atomic theory and provides information about atomic structure and subatomic particles. It discusses Dalton's four main postulates, including that atoms are indivisible and atoms of different elements combine in whole number ratios. The document also outlines the discoveries of key subatomic particles like electrons, protons, and neutrons by scientists such as Thomson, Rutherford, and Chadwick. It describes Bohr's model of the atom and introduces concepts like orbitals, electron configuration, and quantum numbers.
The document discusses the formation of ionic and covalent bonds, including how to determine the charge of ions, represent ionic compounds using dot-cross diagrams, and represent covalent molecules using dot-cross diagrams, structural formulas, and Lewis structures to show how atoms share electrons to achieve stable configurations. It also discusses exceptions where central atoms in some covalent compounds have fewer than 8 valence electrons.
This document presents information about chemical bonding from the chemistry project of students Akarshik Banerjee, Pratyush Dey, and Sayantan Biswas. It discusses various types of chemical bonds including covalent bonds, which form when atoms share electron pairs, and ionic bonds, which form through complete electron transfer. It also describes concepts like the octet rule, Lewis dot structures, formal charge, resonance structures, and molecular geometry based on valence shell electron pair repulsion theory. Hybridization and molecular orbitals are explained as well. In summary, the document provides an overview of key concepts in chemical bonding from the perspective of high school chemistry students.
This document contains a 50-question chemistry semester review covering various topics including:
1) Properties of ionic compounds including their high melting points.
2) Bond types and strengths in molecules like NH3.
3) Types of bonds that would form between different sets of elements.
4) Factors that determine states of matter for different compounds at various temperatures including intermolecular forces.
The review covers concepts of bonding, molecular shapes, properties of states of matter, acid/base chemistry and chemical reactions.
This document provides an introduction to organic chemistry, covering topics such as:
- The definition of organic chemistry as the study of carbon compounds.
- Electronic structure of atoms and how they bond through ionic and covalent bonding.
- Resonance structures and how they are used to represent molecules.
- Factors that influence acidity such as electronegativity, size, and resonance.
- The definitions of nucleophiles and electrophiles and their roles in bond formation.
This document summarizes several theories and concepts related to stereochemistry in main group compounds, including:
1) VSEPR theory which describes the geometry of molecules based on electron pairs around the central atom. It explains linear, trigonal, tetrahedral, trigonal bipyramidal, and octahedral geometries.
2) Bent's rule which describes how atomic s-character concentrates in orbitals directed toward electropositive substituents.
3) Walsh diagrams which use molecular orbital energies to determine molecular geometry.
4) Berry pseudorotation and fluxionality concepts which explain rapid ligand exchange in molecules like PF5.
This document discusses atomic structure and properties. It begins by reviewing the basic atomic model including protons, neutrons, and electrons. It then discusses atomic number, mass number, isotopes, ions, and electron configuration. The document also covers average atomic mass and mass spectrometry. It concludes by discussing the Bohr model of the atom and how this helped explain emission spectra of elements. In summary, the key topics covered are the fundamental particles of atoms, atomic notation, isotopes, ions, electron configuration, and early atomic structure models.
Dr. B. R. Thorat teaches chemistry at Ismail Yusuf College in Maharashtra, India. The document discusses orbital overlapping and conditions required for covalent bond formation. It also explains concepts of atomic orbitals, hybridization, resonance structures, and formal charges through examples of diatomic and polyatomic molecules. Valence bond theory alone cannot explain bonding in molecules like methane, requiring the additional concept of orbital hybridization.
1. Hydrogen bonding occurs between hydrogen atoms attached to electronegative atoms like oxygen, fluorine, and nitrogen of one molecule and an electronegative atom of another molecule.
2. Water is able to form extensive hydrogen bonding networks between molecules due to each water molecule having two hydrogen atoms and two lone pairs of electrons on the oxygen atom.
3. The hydrogen bonding network in liquid water is responsible for its unique properties, while hydrogen bonding in ice forms its crystalline lattice structure.
IB Chemistry on Redox, Oxidation states and Oxidation numberLawrence kok
This document provides information on oxidation states and oxidation numbers of elements. It lists common oxidation states of metals in groups 1-3 and transition metals. It also lists common oxidation states of nonmetals in groups 5-7. Rules are provided for assigning oxidation numbers based on the element and its bonds. Oxidation numbers are assigned by assuming ionic bonding and imagining electron flow to the more electronegative element. Exceptions and examples are given. Polyatomic ions and molecules are discussed.
This document discusses plane waves. Plane waves are waves whose wavefronts are infinite parallel planes. They have equal amplitude in two dimensions. Some key points:
- Plane waves transfer energy from one point to another without transferring matter. The transfer of energy is called wave motion.
- Examples of plane waves include electromagnetic waves in the far field region of an antenna and water waves that appear planar from a distance.
- The basic equations for an electromagnetic plane wave show that the electric and magnetic fields oscillate sinusoidally as they propagate in the direction of the wave vector.
- Waves can be one-dimensional, two-dimensional, or three-dimensional depending on whether the object has length only,
IB Chemistry on Born Haber Cycle and Lattice EnthalpyLawrence kok
The document describes the Born-Haber cycle process for determining the lattice enthalpy of various ionic compounds such as LiCl, NaCl, KCl, NaBr, NaF, and NaH. The Born-Haber cycle involves breaking down the ionic compound into gaseous ions and then measuring the standard enthalpy changes associated with steps such as atomization, ionization, and electron affinity which are used to calculate the unknown lattice enthalpy value. Diagrams of the Born-Haber cycle are provided for each example compound.
1. The document discusses experimental methods to verify the validity of the Debye-Hückel and Onsager equations, which describe electrolyte behavior at high dilutions.
2. Using liquid membrane cells, studies were able to examine dilutions as high as 10-4 M and found that the Debye-Hückel equation holds for 1:1 electrolytes but shows negative deviations for 2:2 or 3:2 electrolytes.
3. Conductance measurements of various electrolytes at different temperatures found close agreement with the Onsager equation at low concentrations, validating it as accurately describing conductance behavior at high dilutions.
Includes a discussion of Voltaic and electrolytic cells, the Nernst equation and the relationship between electrochemical processes, chemical equilibrium and free energy.
**More good stuff available at:
www.wsautter.com
and
http://www.youtube.com/results?search_query=wnsautter&aq=f
This document describes different types of covalent bonding, including:
1) Covalent bonds form when electron pairs are shared between two atoms, attracting their nuclei together.
2) Covalent compounds are usually composed of nonmetals and include many familiar substances like water and methane.
3) Dot and cross diagrams can represent covalent bonds, showing shared electron pairs between atoms.
This document is the first page of a physics exam from the University of Cambridge. It provides instructions to candidates regarding writing their information on the exam, using appropriate materials, and answering all questions. It includes a list of relevant physics formulas, constants, and units that may be used for reference. The exam consists of 15 printed pages and 1 blank page.
T21 IB Chemistry- Spectroscopy continued Robert Hughes
This document outlines lessons for a spectroscopy course, including:
1. A review of infrared spectroscopy, mass spectroscopy, and proton NMR spectroscopy.
2. Practice using these techniques to answer exam questions and deduce organic compound structures.
3. Details on high resolution proton NMR, including spin-spin coupling and interpreting splitting patterns.
4. Interpreting NMR spectra to determine functional groups and hydrogen environments.
5. A brief overview of X-ray crystallography and some practice questions reviewing all three techniques.
The course aims to build students' skills in applying spectroscopy methods to organic compound analysis through lessons, examples, and practice questions.
The document discusses the use of variant question papers by CIE (Cambridge International Examinations) for some of their popular assessments. It explains that while the content and standard of assessment remains the same, there are now two variants of the question papers, mark schemes, and principal examiner reports available for one component. This provides centers with more past examination material than usual to access. It includes both variants in the document. It also provides contact information for any questions about these changes and instructions for identifying the relevant parts of the document.
IB Chemistry on Redox, Oxidizing, Reducing Agents and writing half redox equa...Lawrence kok
The document discusses oxidation numbers (also called oxidation states), which are used to keep track of electrons in chemical reactions. Some key points:
- Oxidation numbers are assigned to each atom in a chemical species by assuming ionic bonding and counting electrons.
- Common rules are outlined for assigning oxidation numbers to elements, such as metals in Group 1 have a +1 oxidation state and nonmetals in Group 7 have a -1 oxidation state.
- Oxidation numbers can be used to determine if a reaction is a redox reaction by looking for changes in oxidation numbers between reactants and products.
- Transition metals can have multiple common oxidation states. Roman numerals are used to distinguish, such as
This document is a chemistry exam paper that consists of 7 questions testing knowledge of various chemistry concepts. It provides instructions for candidates on how to answer the questions, including writing in boxes provided and not writing in barred code or grey areas. It also lists the number of marks allocated for each question or part of a question. Finally, it states that the exam paper is made up of 14 printed pages and 2 blank pages.
csonn t1 atoms, molecules and stoichiometrycheeshengonn
Here are the key steps to solve this problem:
1) Isotope 35Cl has a relative abundance of 75.8%
2) Isotope 37Cl has a relative abundance of 24.2%
3) The relative atomic mass of 35Cl is 35 amu
4) The relative atomic mass of 37Cl is 37 amu
5) Use the formula: Relative Atomic Mass = Σ (Relative Abundance of isotope x Atomic Mass of isotope)
6) For 35Cl: Relative Abundance = 75.8%, Atomic Mass = 35 amu. So, contribution is 75.8% of 35 = 26.43
7) For 37Cl: Relative Abundance = 24
The document summarizes Dalton's atomic theory and provides information about atomic structure and subatomic particles. It discusses Dalton's four main postulates, including that atoms are indivisible and atoms of different elements combine in whole number ratios. The document also outlines the discoveries of key subatomic particles like electrons, protons, and neutrons by scientists such as Thomson, Rutherford, and Chadwick. It describes Bohr's model of the atom and introduces concepts like orbitals, electron configuration, and quantum numbers.
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 chemical bonding, including:
1. Defining ionic and covalent bonding, and explaining how different types of bonds are formed through electron sharing or transfer.
2. Describing the properties of ionic and covalent compounds, such as high melting points for ionic solids and variable states of matter for covalent substances.
3. Illustrating examples of single, double, and triple covalent bonds through Lewis dot structures of molecules like H2, O2, and N2.
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 structures for various molecules by placing the atoms and distributing electrons to achieve full octets. Exceptions to the octet rule are also noted. Hybridization and theories of covalent bonding such as valence bond theory are introduced.
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.
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 chemical bonding, specifically ionic and covalent bonding. Ionic bonding involves the transfer of electrons between metals and non-metals to form ions that achieve stable noble gas electron configurations. Covalent bonding involves the sharing of electrons between non-metals. Both ionic and covalent bonds form when atoms attain noble gas electron configurations. Ionic compounds have characteristics like high melting points and conductivity when molten or dissolved, while covalent compounds have low melting points and are non-conductive.
Chemical bonds form when atoms attract each other and bind together. There are three main types of bonds: ionic bonds form when a metal transfers electrons to a non-metal, metallic bonds involve delocalized electrons that move freely between metal atoms, and covalent bonds occur when two non-metals share pairs of electrons. Ionic bonds are strong but brittle, metallic bonds allow metals to conduct heat and electricity, and covalent bonds can be single, double or triple depending on how many electron pairs are shared.
Ionic bonding occurs through the transfer of electrons between atoms to form ions, resulting in electrostatic attraction. Covalent bonding involves the sharing of electrons between atoms to form molecules. The shape of covalently bonded molecules can be predicted using VSEPR theory. Hydrogen bonding and van der Waals forces are weaker intermolecular forces that influence properties like boiling points. Bonding types exist on a continuum between purely ionic and purely covalent.
The document discusses valence bond theory and molecular orbital theory. It explains that valence bond theory describes covalent bond formation through the overlapping of atomic orbitals between atoms. There are two types of covalent bonds: sigma bonds formed by head-on overlapping and pi bonds formed by side-by-side overlapping. Hybridization involves mixing atomic orbitals to form new hybrid orbitals that can accommodate molecular geometries. Common hybridizations include sp, sp2, and sp3 which result in linear, trigonal planar, and tetrahedral geometries respectively.
A day in the life of a Carbon atom - St Michael's Catholic Primary School bookGlobal CCS Institute
This student's self-illustrated storybook was a huge hit with the expert panel from the CarbonKids CCS communication challenge for Australia's National CCS Week.
The path of a carbon atom in a glucose molecule from the time entering your body in a dry biscuit to moving to a body cell and being burnt in respiration that occurs there and then in carbon dioxide that is then breathed out.
Carbon compounds can be divided into organic and inorganic compounds. Organic compounds contain carbon and are obtained from living things, having low boiling points. Inorganic compounds do not come from living things and have higher boiling points. Hydrocarbons are organic compounds made of only carbon and hydrogen. They can be saturated, containing only single bonds, or unsaturated, containing double or triple bonds. The molecular and structural formulas provide information on the atoms and bonds in a molecule. Naming carbon compounds according to IUPAC guidelines involves a stem/root indicating the number of carbons and an ending denoting the compound class.
This document discusses carbon bonding and the formation of carbon compounds. It explains that carbon can form strong covalent bonds with other carbon atoms through a process called catenation, allowing it to form straight chains, branches, and rings. This bonding ability arises because carbon is tetravalent and can hybridize its orbitals, taking on different hybridization states like sp, sp2, and sp3. Some carbon compounds exhibit resonance, where electrons are delocalized over multiple carbon atoms. This results in more stable structures that are hybrids of different resonant forms. Overall, carbon's unique bonding properties allow it to form a diverse array of stable organic compounds.
The document discusses the combustion stoichiometry of benzene (C6H6) as a fuel. It defines the complete combustion reaction of benzene with air and calculates the stoichiometric air to fuel ratio. It also discusses fuel lean and fuel rich mixtures, showing how the composition of combustion products varies with excess or insufficient air. Plots are included showing how nitrogen content, carbon dioxide content, and dew point temperature change with air to fuel ratio.
The document discusses different types of atomic bonds and properties of compounds formed by these bonds. It defines valence number as the number of electrons in an atom's outer energy level, and oxidation number as the number of electrons an atom will gain, lose, or share during bonding. There are three main types of bonds: ionic bonds form ionic compounds with charged particles; covalent bonds form covalent compounds with weaker forces; and metallic bonds form metals with properties like luster and conductivity. Lewis dot diagrams are used to represent valence electrons and how elements combine.
This document discusses chemical bonding and macromolecular structures. It begins by explaining the different types of bonds - ionic bonds formed between metals and non-metals by electron transfer, and covalent bonds formed between non-metals by electron sharing. It describes the properties of ionic and covalent compounds. It then discusses macromolecular structures found in substances like diamond, graphite and metals. It explains metallic bonding and compares the structures and properties of diamond and graphite. In the end, it discusses the different uses of diamond and graphite based on their properties.
This document discusses chemical bonding and molecular structure. It begins by describing ionic and covalent bonding, including how molecular orbitals form through the overlap of atomic orbitals. It then discusses how valence electron Lewis dot structures are used to represent electron distribution in molecules as bond pairs and lone pairs. Rules for constructing Lewis structures, such as the octet rule, are covered. Exceptions to the octet rule for certain elements are also explained. Finally, the concept of resonance structures and using formal charges to determine the most important Lewis structure are introduced.
The document summarizes key concepts from Chapter 2 of an organic chemistry textbook, including:
1) Molecular orbital theory and how atomic orbitals combine to form bonding and antibonding molecular orbitals through sigma and pi bonding.
2) Hybridization of atomic orbitals (sp, sp2, sp3) and how this determines molecular geometry and bond angles.
3) Intermolecular forces such as hydrogen bonding, dipole-dipole interactions, and London dispersion forces and how they influence properties like boiling points.
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2. Chemical bond
• Chemical bonds are electrostatic forces (attraction between
positive charge and negative charge) that bind particles together to
form matter.
• Different types of chemical bonds:
• a) Metallic bond
• b) Ionic bond
• c) Covalent bond
• d) Intermolecular forces (NEED TO SPECIFY WHICH ONE - LATER!)
3. Ionic bonding
• Ionic compound has lattice of cations + anions.
• Ionic bond is the electrostatic force of attraction
between oppositely charged ions formed by
electron transfer.
• Octet rule - general but not 100% always!
4. Ionic bonding - ‘dot-and-
cross’
Draw dot-and-cross diagrams for the following ionic
compounds.
a) MgO
b) CaCl2
c) Na2O
5. Ionic bonding
• NOT electron transfer
• INVOLVES electron transfer
• Also known as electrovalent bonding (CIE syllabus)
• Metals lose electrons and form cations => think about
ionisation energies (topic 2)
• Across a period, Zeff, hence first IE?
• Down a group, Zeff, hence first IE?
6. Ionic bonding
• Non-metals gain electrons
• (A2 syllabus: electron affinity - 1st electron affinity -
enthalpy change when one mole of electrons is added
to one mole of gaseous atoms to form one mole of
gaseous singly charged anions)
• O (g) + e- —> O- (g) 1st EA = -140 kJmol-1
• O+ (g) + e- —> O2- (g) 2nd EA = + 798 kJmol-1
• WHY difference? - think of electronic configuration?
7. Ionic bonding - Lattice
(formation) enthalpy
• the enthalpy change
• when 1 mole of ionic solid crystal
• is formed from its scattered gaseous ions.
• under standard conditions (298 K and 1 atm pressure)
• Lattice formation enthalpies are always negative (ΔH <
0)
• Eqn example?
8. Ionic bonding - Lattice
(formation) enthalpy
• Greater charge densities of ions, the more they
attract each other, the larger the lattice enthalpy
• The more exothermic the lattice enthalpy, for ionic
compounds, the higher the m.p.
9. Ionic bonding
• Most important factor is LATTICE ENERGY
• The lattice (formation) enthalpy is the enthalpy change
when 1 mole of solid crystal is formed from its
scattered gaseous ions.
• E.g. Na+ (g) + Cl- (g) —> NaCl (s)
• E.g. Write an equation to show lattice formation
enthalpy change for formation of MgO and Al2O3.
• Lattice formation enthalpy is always EXOTHERMIC
10. Ionic bonding
• Two factors affect lattice energy:
• a) charges on the ion
• b) size of the ion
• X-ray diffraction studies - absolute proof
• For simple ions,
• a) charge determines the balance between numbers of cations
and anions
• b) radii determine the way ions pack in lattice
11. Ionic bonding - NaCl
• Called the “rock salt structure”
• NaCl and MgO has this ionic lattice structure
• 6:6 coordination number - meaning?
• Can sketch?
12. Ionic bonding - Properties
• High melting and boiling points
• Good electrical conductivity only when molten
• Generally soluble in polar solvents
13. Ionic bonding - Q1
Which of the following statements are correct for the
sequence of compounds below considered from left to right?
NaF MgO AlN SiC
(1) The electronegativity difference between the elements in
each compound increases.
(2) The formulae-units of these compounds are isoelectronic
(have the same number of electrons).
(3) The bonding becomes increasingly covalent.
15. Covalent bonding
• The covalent bond is the electrostatic attraction
between the localised shared electrons and the two
positively charged nuclei.
• Covalent bonds are directed in space.
• NOT sharing electrons
• INVOLVES sharing of electrons
• Simple molecular or giant molecular? Single, double or
triple bonds? Sigma or pi bond?
18. Covalent bonding - sigma
and pi bonds
• Extent of bonding depends on orbital overlap
• similar in energy (i.e. similar sized orbitals)
• similar in symmetry
• Can you form pi bond without sigma bond?
• Absolutely not; if given the options of forming a sigma or a pi
bond, the two atoms “prefer” to form a sigma bond.
• A sigma bond is formed from more effective overlap of the
atomic orbitals compared to pi bond.
• A sigma bond is thus more stable and stronger than a pi bond.
19. Covalent bonding - sigma
and pi bonds
• Can there be two sigma bonds formed between two
atoms?
• Absolutely not; formation of two sigma bonds will result
in too much accumulation of electron density within
the inter-nuclei region.
• The inter-electronic repulsion will be too great.
• The different ways of forming sigma and pi bonds
spread out the inter-nuclei electron density to
minimise inter-electronic repulsion but maximising
bond strength.
20. Covalent bonding - sigma
and pi bonds
Sigma (σ) bond Pi (π) bond
Formed due to the axial overlap of two
orbitals (‘s-s’, ‘s-p’or’p-p’).
Formed by the lateral (sideways) overlap of
two ‘p’ orbitals.
Only one sigma bond exists between two
atoms.
There can be more than one pi bonds between
the two atoms.
The electron density is maximum and
cylindrically symmetrical about the bond axis.
The electron density is high along the
direction at right angles to the bond axis.
Free rotation about the sigma bond is possible. Free rotation about the pi bond is not possible.
This bond can be independently formed, i.e.,
without the formation of a pi bond.
The pi bond is formed after the sigma bond
has been formed,
Sigma bond is relatively strong. Pi bond is a relatively weaker bond.
21. Covalent bonding - sigma
and pi bonds - Q1
Which diagram describes the formation of a pi bond
from the overlap of its orbitals?5
5 Which diagram describes the formation of a π bond from the overlap of its orbitals?
D
A
B
C
6 For an ideal gas, the plot of pV against p is a straight line. For a real gas, such a plot shows a
22. Covalent bonding - sigma
and pi bonds - Q2
Which statements about covalent bonds are correct?
(1) A triple bond consists of one pi bond and two
sigma bonds.
(2) The electron density in a sigma bond is highest
along the axis between the two bonded atoms.
(3) A pi bond restricts rotation about the sigma bond
axis.
24. Covalent bonding - sigma
and pi bonds - Q3
What is always involved in a carbon-carbon pi bond?
(1) a shared pair of electrons
(2) a sideways overlap of p orbitals
(3) delocalised electrons
26. Covalent bonding - sigma
and pi bonds - Q4
Carvone is found in spearmint.
How many sigma and pi bonds are present in this
molecule?
r
proton number
r
proton number
r
proton number
r
proton number
20 Carvone is found in spearmint.
C
C
C
C
CH2
H3C
H2C
CH2
CH3
C
H
HO
carvone
How many σ and π bonds are present in this molecule?
σ π
A 13 3
28. Covalent bonding - sigma
and pi bonds - Q5
The diagram shows a molecule that has sigma bonds
and pi bonds.
How many sigma and pi bonds are present in this
molecule?
at is compound X?
CH3CO2C2H5
CH3CH2COCH3
(CH3)3COH
CH3CH2CHOHCH3
e diagram shows a molecule that has σ bonds and π bonds.
C
O
OCH2CH2 CH2 CH2CHCH
w many σ bonds are present in this molecule?
15 B 17 C 18 D 21
30. Covalent bonding - sigma
and pi bonds - Q6
4
8 A covalent molecule contains
● 14 electrons,
● one lone pair of electrons,
● two π bonds.
What is the molecule?
A C2H4 B HCN C H2O2 D N2
9 Which value is essential to calculate the lattice energy of the compound NaH?
32. Covalent bonding - sigma
and pi bonds - Q7
Which statement about bond formation is not correct?
A. A triple bond consists of one σ bond and two π bonds.
B. A π bond restricts rotation about the σ bond axis.
C. Bonds formed from atomic s orbitals are always σ
bonds.
D. End-to-end orbital overlap results in a bond with
electron density above and below the bond axis.
38. Covalent bonding - Dot-and-cross
diagrams & Lewis diagrams
• Atoms share electrons in order to complete their “octet” of
electrons
• Some don't achieve an octet as they don’t have enough
electrons - Al in AlCl3, B in BCl3
• Others share only some - if they share all valence
electrons, their octet is exceeded - ie NH3, NCl3 and H2O,
OF2 (cf PCl5, SF6)
• Atoms of elements in the 3rd period onwards can exceed
their octet because they are not restricted to 8 electrons in
their outer shell (cf n=3 shell can use 3d for bonding).
39. Dative (coordinate) covalent
bonding
• A dative covalent (or coordinate) bond is a
covalent bond in which both the electrons shared
come from the same atom.
• Donor species will have lone pairs in their outer
shells
• Acceptor species will be short of their “octet” or
maximum
43. Dative (coordinate) covalent
bonding - Q1
solutions have the same effect on litmus.
What is element X?
A sodium
B magnesium
C aluminium
D phosphorus
15 Aluminium chloride sublimes at 178o
C.
Which structure best represents the species in the vapour at this temperature?
Cl
Cl
Cl
Cl
Al
Cl
Cl
Al
Cl
Cl
Cl
Cl
Al
Cl
Cl
Al
Al + 3Cl Al3+
(Cl –
)3
A B C D
16 Use of the Data Booklet is relevant to this question.
What mass of solid residue can be obtained from the thermal decomposition of 4.10 g of
anhydrous calcium nitrate?
44. Dative (coordinate) covalent
bonding - Q2
AlCl3 reacts with LiAlH4 and (CH3)3N to give (CH3)3NAlH3.
Which statement about (CH3)3NAlH is correct?
A. It contains hydrogen bonding.
B. It is dimeric.
C. The Al atom has an incomplete octet of electrons.
D. The bonds around the Al atom are tetrahedrally
arranged.
48. Dative (coordinate) covalent
bonding - Q4
Which diagram correctly shows the bonding in the
ammonium ion, NH4
+?
What is the same in an atom of 4
He and an atom of 3
H?
A the number of electrons
B the number of neutrons
C the number of protons
D the relative atomic mass
2 Which diagram correctly shows the bonding in the ammonium ion, NH4
+
?
N
H
H
H H
+
A
N
H
H
H H
+
B
N
H
H
H H
+
C
N
H
H
H H
+
D
key
N electron
H electron
50. Dative (coordinate) covalent
bonding - Q5
Why does aluminium chloride, Al2Cl6, sublime at the
relatively low temperature of 180°C?
(1) The intermolecular forces between the Al2Cl6
molecules are weak.
(2) The co-ordinate bonds between aluminium and
chlorine are weak.
(3) The covalent bonds between aluminium and
chlorine are weak.
52. Dative (coordinate) covalent
bonding - Q6
In the gas phase, aluminium chloride exists as the dimer,
Al2Cl6. By using this information, which of the following
are structural features of the Al2Cl6 molecule?
(1) Each aluminium atom is surrounded by four chlorine
atoms.
(2) There are twelve non-bonded electron pairs in the
molecule.
(3) Each aluminium atom contributes electrons to four
covalent bonds.
54. Dative (coordinate) covalent
bonding - Q7
Aluminium chloride catalyses certain reactions by forming
carbocations (carbonium ions) with chloroalkanes as shown.
Which property makes this reaction possible?
A. AlCl3 is a covalent molecule.
B. AlCl3 exists as the dimer Al2Cl6 in the vapour.
C. The aluminium atom in AlCl3 has an incomplete octet of
electrons.
D. The chlorine atom in RCl has a vacant p orbital.
56. Covalent bonding - simple
covalent ions
• Ammonium salt has two types of bonding
• Ionic bonding between ammonium ion and anion
• Covalent bonding (including dative covalent) within
ammonium ion
• Other examples include: nitrates, sulfates, phosphates,
carbonates, etc.
• How we know this? - easy - X-ray crystallographic data!
- bond length - compare single vs double bond length?
58. Molecular geometry -
VSEPR
• Valence shell electron pair repulsion theory
• What is an electron pair?
• What is a bonding electron pair?
• What is a non-bonding electron pair/lone pair?
• Electron clouds repel each other. Why?
• Extent of repulsion: LP/LP > BP/LP > BP/BP
61. Molecular geometry -
VSEPR
H
H
O
Bent line 2 2 104.5 OCl2, H2S, OF2 ,
SCl2
.. ..
O
H H
P
F
F
F
F
F
S
FF
FF
F
F
SCl2
Trigonal
Bipyramidal
5 0 120 and 90 PCl5
Octahedral 6 0 90 SF6
Occasionally more complex shapes are seen that are variations of octahedral and trigonal
bipyramidal where some of the bonds are replaced with lone pairs.
e.g XeF4 e.g. BrF5 e.g I3 e .g.ClF3 e.g. SF4
Remember lone pairs repel more than bonding pairs and so reduce bond angles
X
:
X:X
:
:X
::
62. Molecular geometry -
VSEPR
In this order,
a) state number of bonding pairs and lone pairs of electrons
b) state that electron pairs repel and try to get as far apart as
possible (or to a position of minimum repulsion)
c) IF there are no lone pairs, state that the electron pairs repel
equally
d) IF there are lone pairs of electrons, then state that lone
pairs repel more than bonding pairs.
e) state actual shape and bond angle.
63. Molecular geometry -
VSEPR
• Species with an odd number of valence electrons,
e.g. NO
• Electron deficient compound, e.g. BH3, BF3
• Species with expanded valence shells, e.g. SF6,
PCl5
68. Molecular geometry -
VSEPR - Q2
In which sequences are the molecules quoted in
order of increasing bond angle within the molecule?
1 H2O NH3 CH4
2 H2O SF6 BF3
3 CH4 CO2 SF6
70. Molecular geometry -
VSEPR - Q3
Chloroethene, CH2=CHCl, is the monomer of pvc.
What are the C-C-C bond angles along the polymeric
chain in pvc?
A. They are all 109 °.
B. Half are 109 ° and half are 120 °.
C. They are all 120 °.
D. They are all 180 °.
74. Molecular geometry -
VSEPR - Q5
Which of the following molecules and ions have a
regular trigonal planar shape?
(1) AlCl3
(2) CH3
+
(3) PH3
(4) BCl3
(5) NH3
76. Molecular geometry -
VSEPR - Q6
Methyl isocyanate, CH3NCO, is a toxic liquid which is used in
the manufacture of some pesticides. In the methyl isocyanate
molecule, the sequence of atoms is H3C-N=C=O.
What is the approximate angle between the bonds formed by
the N atom?
3
4 Methyl isocyanate, CH3NCO, is a toxic liquid which is used in the manufacture of some
pesticides.
In the methyl isocyanate molecule, the sequence of atoms is H3C—N C O.
What is the approximate angle between the bonds formed by the N atom?
A
104
B
109
C
120
D
180
N CH3
C O
N C
H3
C
ON C
H3
C
ON C
H3
C
O
5 At room temperature and pressure chlorine does not behave as an ideal gas.
At which temperature and pressure would the behaviour of chlorine become more ideal?
78. Molecular geometry -
VSEPR - Q7
Which statements about bond angles are correct?
(1) The bond angle in SO2 is smaller than the bond
angle in CO2.
(2) The bond angle in H2O is smaller than the bond
angle in CH4.
(3) The bond angle in NH3 is smaller than the bond
angle in BF3.
80. Molecular geometry -
VSEPR - Q8
The diagram shows an example of an organic nitrate
molecule.
What is the correct order of the bond angles shown in
ascending order?
8
B H3O+
C OD–
D OH–
s in photochemical smog can cause breathing difficulties.
hows an example of an organic nitrate molecule.
H C
H
H
O
C O O NO2
1
2
3
rect order of the bond angles shown in ascending order (smallest
B 2 → 1 → 3 C 3 → 1 → 2 D 3 → 2 → 1
82. Molecular geometry -
VSEPR - Q9
The CN- is widely used in the synthesis of organic
compounds. What is the pattern of electron pairs in
this ion?
No of bonding pairs of electrons = ?
No of lone pairs on carbon atom = ?
No of lone pairs on nitrogen atom = ?
84. Molecular geometry -
VSEPR - Q10
The antidote molecule shown can help to prevent liver
damage if someone takes too many paracetamol tablets.
What is the order of decreasing size of the bond angles
x, y and z?
A ionic radius
B ionisation energy
C neutron/proton ratio
D rate of reaction with water
he antidote molecule shown can help to prevent liver damage if someone takes too
aracetamol tablets.
H S C
H H H
H H H
C N
x y
z
represents a
lone pair
What is the order of decreasing size of the bond angles x, y and z?
largest smallest
A x y z
B x z y
85. Molecular geometry -
VSEPR limits
17 HF Hydrogen fluoride Hydrogen fluoride
17 HCl Hydrogen chloride Hydrogen chloride
17 HBr Hydrogen bromide Hydrogen bromide
17 HI Hydrogen iodide Hydrogen iodide
The electron-deficient hydrides are those that cannot complete an octet of electrons around the central atom. They are
chiefly the Group 13 elements, although the gas-phase only species BeH2 also fits this description. The electron precise
compounds are those that have an octet of electrons, while the electron-rich elements have additional electrons belonging to
the central atom that function as lone pairs. Note that although they are all electron rich, this group of compounds vary
greatly in the availability of these extra electrons.
Note that while the structures of these hydrides follow similar patterns, i.e. all the EH3 in group 15 are pyramidal, the
detailed structures differ significantly. Consider the following table of bond lengths for both groups 15 and 16:
Group 15 hydrogen compounds Bond angle Group 15 hydrogen compounds Bond angle
NH3 106.6° H2O 104.5°
PH3 93.8° H2S 92.1°
AsH3 91.8° H2Se 91°
SbH3 91.3° H2Te 89°
Source: A.F. Wells, Structural Inorganic Chemistry, Oxford University Press (1984)
In the last problem set,
you have explored the
origins of some of these
changes using molecular
orbital methods. The
bending of H–E–H
systems is due to a Jahn-
Teller effect that lowers
the energy of what would
be a degenerate set of E
orbitals in the linear
molecule. The greater
bending in the third and
subsequent periods is due
to a second-order Jahn-
89. Giant molecular compounds
• Some covalent compounds are not discrete molecular
compounds.
• Giant molecular compounds include,
(a) diamond
(b) SiO2, silicon (IV) oxide, sand - all same
(c) graphite
(d) BN (isoelectronic to carbon)
(e) silicon - same like diamond
90. Giant molecular compounds
- diamond
• mp of diamond = 3350 °C
• mp of silicon = 1410 °C
• For melting in giant molecular compounds, a
great amount of energy is required to overcome
the strong covalent bonds between the atoms.
• Bond strength: BDE(C-C) > BDE(Si-Si)
91. Giant molecular compounds
- graphite
• C-C (between layers) < C-C (same adjacent layer)
• Longer bond lengths => weaker bonding between these atoms
• VDW forces of attraction holds layers of graphite together, not
actual covalent bonds
92. Chemical bonding - Q1
Which solid exhibits more than one kind of chemical
bonding?
A. brass
B. copper
C. diamond
D. ice
94. Chemical bonding - Q2
Which of the following solids contain more than one
type of chemical bond?
1. brass (an alloy of copper and zinc)
2. graphite
3. ice
98. Chemical bonding - Q4
Which diagrams represent part of a giant molecular
structure?
The responses A to D should be selected on the basis of
A B C D
1, 2 and 3
are
correct
1 and 2
only are
correct
2 and 3
only are
correct
1 only
is
correct
No other combination of statements is used as a correct response.
31 Which diagrams represent part of a giant molecular structure?
1 2 3
= C = C = Na
= C
32 Which reactions are redox reactions?
1 CaBr2 + 2H2SO4 → CaSO4 + Br2 + SO2 + 2H2O
100. Covalent bonding - Bond
dissociation enthalpy (BDE)
• Energy required
• to break one mole of covalent bond
• to give separated atoms
• with everything being in the gas state
• Eqn example?
101. Covalent bond - bond
energy
• Bond dissociation energy (BDE) is the energy required
to break the bond between two covalently bonded atoms.
• Polyatomic molecules?
• X-ray diffraction - covalent bond length and covalent
bond radii
• By halving the interatomic distances obtained for diatomic
elements = covalent bond radii
• Suggest when would actual covalent radii be very
different to that predicted by tabulated covalent radii.
102. Covalent bond - bond
polarity
• Bonding electron pair is shared equally between
two same atoms that form a covalent bond.
• For any two unlike atoms, the bonding electron pair
sharing is always unequal.
• Unequal sharing of electrons can have two
extremes - both leading to ionic bond.
bond radii, but these theoretical values often differ from the experi-
mental values; the greatest deviations occur when elements ofwidely
different electronegativities are joined together.
ELECTRONEGATIVITY
If two like atoms form a covalent bond by sharing an electron pair,
for example
x F * FxX X X X
it is clear that the pair will be shared equally. For any two unlike
atoms, the sharing is always unequal and depending on the nature
of the two atoms (A and B say)we can have two extreme possibilities
or A :B i.e. A+
B"
A + B->A : B ^
equal
sharing X
or A ; g ie A- B +
and an ionic bond is formed. There are many compounds which lie
103. Bond polarity -
Electronegativity
• Many compounds are intermediate between truly covalent
(equal sharing) and truly ionic.
• Electronegativity is a measure of the tendency of an atom to
attract a bonding pair of electrons.
• Fajan’s rules state that covalent character in ionic compounds
increases if,
• a) small cations - highly polarising
• b) large anions - highly polarisable
• c) high charge on cations and anions
104. Covalent bond - Electronegativity
• Electronegativity is a measure of the tendency
of an atom to attract a bonding pair of electrons.
106. Covalent bond - Electronegativity
- Factors affecting it
• Think of Zeff
• The attraction that a bonding pair of electron feels
for a particular nucleus depends on:
• a) the number of protons in the nucleus
• b) the distance from the nucleus
• c) the amount of screening by inner electrons
107. Electronegativity - diagonal
relationship
• Boron is a Gp III non-metal with some properties like silicon.
• Beryllium is a Gp II metal with some properties resembling
aluminium.
• Electronegativity reasoning:
• Be = 1.5, B = 2.0
• B = 2.0, Al = 1.5
• Similar electronegativity, likely to form similar types of
bonds, hence similar chemistry
108. Ionic or covalent? - Q1
Which pairs of compounds contain one that is giant
ionic and one that is simple molecular?
(1) Al2O3 and Al2Cl6
(2) SiO2 and SiCl4
(3) P4O10 and PCl3
(4) B(NMe3)3 and BN
(5) C(CH3)4 and diamond
110. Ionic or covalent? - Q2
Which chlorine compound has bonding that can be
describe as ionic with some covalent character?
A. NaCl
B. MgCl2
C. AlCl3
D. SiCl4
114. Ionic or covalent? - Q4
When barium metal burns in oxygen, the ionic
compound barium peroxide, BaO2 is formed. Which
dot-and-cross diagram represents the electronic
structure of the peroxide anion in BaO2?
3
4 When barium metal burns in oxygen, the ionic compound barium peroxide, BaO2, is formed.
Which dot-and-cross diagram represents the electronic structure of the peroxide anion in BaO2?
A C DB
electron from
first oxygen atom
electron from
second oxygen atom
electron from
barium atom
key
5 In this question, the methyl group, CH3, is represented by Me.
116. Dipole moment
• Unequal distribution of charge produced when
elements of different electronegativities polarises a
covalent bond joining them.
• Unless this polarity is balanced by an equal and
opposite polarity, the molecule will be a dipole and have
a dipole moment (i.e. hydrogen halide - HF, HCl, HBr,
HI)
• Examples where dipole moments get cancelled out -
there is no net dipole moment => non-polar molecule.
117. How to determine if a
molecule is polar?
As long as a molecule has one of the standard shapes shown
below, with identical bonding atoms and no lone pairs on the
central atom, then the molecule has no net dipole moment.
=> Linear
=> Trigonal planar
=> Tetrahedral
=> Trigonal bipyramidal
=> Octahedral
119. How to determine if a
molecule is polar?
• If the molecule contains one lone pair on the central atom, then the
molecule has a net dipole moment.
• But, a molecule with more than one lone pair on the central atom
may or may not be polar.
• Eg. NF3, H2O, XeF4
120. How to determine if a
molecule is polar?
• Polar bonds ≠ Polar molecule
• A polar molecule has a net dipole moment -
hence permanent dipole intermolecular forces
between their molecules.
121. Polar molecule - Q1
Which of the following molecules has no permanent
dipole?
A. CCl2F2
B. CHCl3
C. C2Cl4
D. C2H5Cl
123. Polar molecule - Q2
Which molecule has the largest overall dipole?
4
5 Which molecule has the largest overall dipole?
C C
C C
C C
A B
O C
C
O C
D
O C O
C
C
H
H
6 The first stage in the industrial production of nitric acid from ammonia can be represented by the
following equation.
4NH3(g) + 5O2(g) 4NO(g) + 6H2O(g)
Using the following standard enthalpy change of formation data, what is the value of the standard
enthalpy change, ∆Ho
, for this reaction?
125. Polar molecule - Q3
Which molecule has the largest overall dipole?
CO2(g) + 3H2(g) CH3OH(g) + H2O(g) ∆H = –49kJmol
What would increase the equilibrium yield of methanol in this process?
A adding a catalyst
B adding an excess of steam
C increasing the pressure
D increasing the temperature
10 Which molecule has the largest overall dipole?
CH3
CH3
H
H
C C
A
CH3
CH3
O C
B
Cl
Cl
O C
C
H3C
Cl
Cl
CH3
C C
D
11 In which substance does nitrogen exhibit the highest oxidation state?
A NO B N2O C N2O4 D NaNO2
12 Red lead oxide, Pb3O4, is used in metal priming paints. It can be made by heating PbO in air.
127. Polar molecule - Q4
The greater the difference between the electronegativities of
the two atoms in a covalent bond, the more polar is the bond.
Which pair will form the most polar covalent bond between
the atoms?
A. chlorine and bromine
B. chlorine and iodine
C. fluorine and chlorine
D. fluorine and iodine
129. Intermolecular forces
• Intermolecular attractions are attractions between
one molecule and a neighbouring molecule.
• The forces of attraction which hold an individual
molecule together are the intramolecular attractions
(chemical bonds - like covalent bonds!)
130. Intermolecular forces - van
der Waals forces
• Between simple covalent molecules and also separate atoms in
noble gases
• VDW (monatomic noble gas atoms) < VDW (simple covalent
molecules) despite similar electronic count [eg. He vs H2]
• In any molecule, the electrons are moving constantly and
randomly. Electron density can fluctuate and parts of the
molecule become more or less negative - small temporary
dipoles
• These instantaneous dipole induces dipoles of opposite sign
on neighbouring molecules.
131. Intermolecular forces - van
der Waals forces
• Two factors affecting strength of VDW:
a) the number of electrons in the molecule
=> the more electrons, the bigger the electron cloud, the
more polarisable is the electron cloud, the stronger is the id-
id interactions
b) the surface area for contact of the molecule
=> the greater the surface area of contact possible between
the molecules, the greater is the extent of the id-id
interactions
132. Intermolecular forces - van
der Waals forces
Room temperature is 298 K (25 °C). bp of Br2 is 59 °C
while bp of Cl2 is -34 °C.
Account for why bromine is a liquid at r.t.p. while
chlorine is a gas.
134. Intermolecular forces - van
der Waals forces
Pentane has lower boiling point than its structural
isomer, dimethylpropane.
bp of pentane is 36 °C while bp of dimethylpropane
is 10 °C.
138. Intermolecular forces -
permanent dipoles
• Only between polar molecules
• Uneven distribution of electrons in polar bonds,
permanent separation of charges (dipoles) found
within polar molecules
• Two criteria required,
• a) there must be polar bonds within the molecules.
• b) there must be a net dipole moment for the
molecule.
139. Intermolecular forces -
hydrogen bond
• Present between molecules that have at least one
highly electronegative atom - F, O or N -
covalently bonded to an H atom.
• Sketch hydrogen bonding.
140. Intermolecular forces -
hydrogen bond
• Strength of a hydrogen bond depends on:
a) dipole moment of the H-X bond [X = F, O, N]
F-H - - - F-H > O-H - - - O-H > N-H - - - N-H
b) ease of donation of a lone pair on Y [Y= F, O, N]
N-H - - - N-H > O-H - - - O-H > F-H - - - F-H
• Overall, hydrogen bond strength is in this order,
F-H - - - F-H > O-H - - - O-H > N-H - - - N-H
141. Intermolecular forces -
hydrogen bond
• In terms of boiling point,
HF: 20 °C [But HF forms the strongest hydrogen bonds, so why?]
H2O: 100 °C
H3N: - 33 °C
• H2O has the highest bp because it can form more extensive
hydrogen bonding than NH3 and HF.
• H2O can form two hydrogen bonds per water molecule
whereas both NH3 and HF can only form one hydrogen bond
per molecule (why? sketch?)
148. Hydrogen bonding -
implications/importance
• Stabilisation of structure of proteins - alpha helix and beta
sheets
• Different base pairings (A’ Level biology - Adenine-Thymine
Guanine-Cytosine)
149. Intermolecular forces - Q1
A crystal of iodine produces a purple vapour when
gently heated. Which pair of statements correctly
describes this process?
type of bond broken formula of purple species
A covalent I
B covalent I2
C induced dipole-dipole I2
D permanent dipole-dipole I2
151. Intermolecular forces - Q2
Which types of intermolecular forces can exist between
adjacent urea molecules?
1. hydrogen bonding
2. permanent dipole-dipole forces
3. temporary induced dipole-dipole forces
e responses A to D should be selected on the basis of
A B C D
1, 2 and 3
are
correct
1 and 2
only are
correct
2 and 3
only are
correct
1 only
is
correct
other combination of statements is used as a correct response.
Which types of intermolecular forces can exist between adjacent urea molecules?
H2N
C
O
NH2
urea
1 hydrogen bonding
2 permanent dipole-dipole forces
3 temporary induced dipole-dipole forces
Ethanol is manufactured by reacting ethene gas and steam in the presence of phosphoric(V)
acid.
C H (g) + H O(g) C H OH(g) ∆H = –45kJmol–1
153. Intermolecular forces - Q3
What is involved when a hydrogen bond is formed
between two molecules?
1. a hydrogen atom bonded to an atom less
electronegative than itself
2. a lone pair of electrons
3. an electrostatic attraction between opposite
charges
154. Intermolecular forces - Q4
The three statements that follow are all true.
Which of these can be explained, at least in part, by
reference to hydrogen bonding?
1. At 0 °C, ice floats on water.
2. The boiling point of propan-2-ol is 82 °C. The
boiling point of propanone is 56 °C.
3. At 20 °C, propanone and propanal mix completely.
155. Intermolecular forces - Q5
Which compound is the only gas at room temperature
and pressure?
A CH3CH2CH2NH2 Mr = 59.0
B CH3CH2CH2OH Mr = 60.0
C CH2OHCH2OH Mr = 62.0
D CH3CH2Cl Mr = 64.5
157. Metallic bonding
• Metallic bond is the electrostatic attraction between the
positive ions and the delocalised valence electrons
• Strength of metallic bonding depends on:
• a) number of valence electrons available for bonding (across
a period, how?)
• b) size of the metal cation (down a group, how?)
159. Metallic bonding
• Malleable and ductile
=> Ability of cations to move over one another without breaking
metallic bonds
• High mp and bp
=> Large amount of energy required to overcome strong metallic bond
(electrostatic attraction) between the positively charged ions and the
“sea of delocalised valence electrons”.
• Good thermal and electrical conductivity
=> Due to presence of delocalised electrons in the metallic lattice
160. Metallic bonding - Q1
Which of the following are features of the structure of
metallic copper?
(1) ionic bonds
(2) delocalised electrons
(3) lattice of ions