Chemistry A Molecular Approach AP Edition Chapter 9 

Chemical Bonding I: The Lewis Model 

9.1 Bonding Models and AIDS Drugs

• Researches used X-ray crystallography to determine the structure of the HIV-protease. This protein is crucial to the cause of AIDS. Without this protein, HIV cannot spread and replicate. WIthout HIV-protease, AIDs cannot develop 
• Scientists tried to develop a molecule that would disable the HIV-protease. They used bonding theories to map out what the molecule would look like and how it would act.
• The molecules developed were called protease inhibitors. They did not cure AIDS but they helped lower infection rates
• Bonding theories explain how atoms bond, why some bonds are stable but others are not, and the shapes of the molecules 
• The Lewis Model, also called Lewis Structures, are made up of dots, lines, and elemental symbols show what a molecule might look like
• Valence electrons are represented with dots
• Lines represent bonds 
• Elemental symbols represent the element  

9.2 Types of Chemical Bonds 

• Chemical bonds are crucial to everyday life. They are what make different compounds rather than singular elements 
• The electrons of one atom are attracted to the protons of another atom and vice versa. At the same time, the electrons of one atom are repelling the electrons of the other atom. The same goes for the protons. Bonds can form when the attraction outways the repulsion. 
• Metals and nonmetals form ionic bonds. Ionic bonds deal with electrons being transferred
• Metals become cations and nonmetals are anions 
• Electrons get transferred to the nonmetal
• Nonmetals and nonmetals form covalent bonds. These atoms share electrons 
• They share electrons because nonmetals typically have high ionization energies 
• Metals and metals form metallic bonds. These bonds have pooled electrons 
• Metals have low ionization energies, so they lose their electrons very easily 
• The electron sea model models metallic bonding 

9.3 Representing Valence Electrons with Dots 

• Valence electrons are in the outermost energy level 
• Lewis Models use dots to represent valence electrons 
• Atoms with eight valence electrons are stable because they have a full outer level 
• Having 8 valence electrons is called an octet 
• Helium is an exception because it can only hold 2 valence electrons. This is called a duet 
• A chemical bond is the sharing or transferring of an electron in order to reach a stable electron configuration. Since the stable configuration is typically 8 electrons, this is called the octet rule 
• Lewis Models best represent covalent bonds, but they can also show ionic 

9.4 Ionic Bonding: Lewis Symbols and Lattice Energies 

Ionic Bonding and Electron Transfer 

• The Lewis symbol of an anion is represented with brackets around the model and the charge in the upper right corner 
• Losing an electrons causes a cation 
• Gaining an electron causes an anion 
• Lewis Models can predict the correct chemical formulas for ionic compounds 

Lattice Energy: The Rest of the Story 

• The transfer of an electron causes energy to be absorbed 
• Lattice energy causes the reaction to be exothermic 
• Lattice energy is the energy associated with the formation of a crystalline lattice of alternation cations and anions from the gaseous state 
• A lattice looks similar to a cube 
• Energy is emitted as heat when a lattice is formed 
• Lattice energy can be calculated with the Born-Haber Cycle 

The Born-Haber Cycle 

• This cycle is a series of steps that represent the formation of an ionic compound 
• The change in enthalpy is known so that lattice energy can be solved for 
• Here are the steps for the NaCl example: First, form gaseous sodium from solid sodium. Then form a chlorine atom from a chlorine molecule. Next ionize the gaseous sodium. The enthalpy change found in this step is the ionization energy of sodium. Then add an electron to the gaseous chlorine. The enthalpy change is the electron affinity of chlorine. Finally form the crystal from the gaseous ions. The enthalpy change is the lattice energy 
• The value of lattice energy will be a large number 
• Hess's law can set the overall enthalpy formation for NaCl 

Trends in Lattice Energies: Ion Size 

• Lattice energy decrease as you go down a column. This is because the radius increases and the potential energy decreases
• As the radius gets larger, atoms cannot attract and get close to each other, therefore not releasing as much energy  

Trends in Lattice Energies: Ion Charge 

• Lattice energy is larger when the magnitude of the ionic charge increases 

Ionic Bonding: Models and Reality 

• Lewis models show that ionic compounds have high melting points because the forces are nondirectional 
• The model does not include any free electrons that might conduct electricity 
• The model does not account for the nonconductitvty of ionic solids 

9.5 Covalent Bonding: Lewis Structures 

Single Covalent Bonds 

• A sharing pair of electrons is called a bonding pair 
• A lone pair is not involved in bonding
• Lone pairs are also called nonbonding electrons 
• Dashes/lines represent a bonding pair  

Double and Triple Covalent Bonds 

• When two electron pairs are shared between atoms it is called a double bond 
• Double bonds are shorter but stronger than single bonds 
• Three electron pairs are called triple bonds
• Triple bonds are shortest and strongest 

Covalent Bonding: Models and Reality 

• Lewis Models explain why atoms form molecules in a particular combination 
• Lewis Models show when an atom as an octet 
• It also explains why covalent bonds are highly directional. The atoms are attracted to each other because they are sharing electron pair(s) in the small space between them 
• Covalent bonds are typically weaker than ionic bonds 

9.6 Electronegativity and Bond Polarity 

• Lewis models are limited because they make it seem like electrons are equally shared between two atoms. In reality, one atom will have a more positive charge while the other has a negative charge 
• A polar covalent bonds is the intermediate nature between a pure covalent bond and an ionic bind 
• Pure covalent and pure ionic are extremes 
• Most covalent bonds are polar covalent 

Electronegativity 

• Electronegativity is an ions ability to attract electrons to itself 
• Electronegativity increases across a period 
• Electronegativity decreases going down a column 
• Fluorine is the most electronegative element 
• Francium is the least electronegative element 
• The larger the atom, the less ability it has to attract electrons 

Bond Polarity, Dipole Moment, and Percent Ionic Character 

• The degree of polarity depends on the electronegativity difference between the two atoms 
• The greater the difference, the more polar the bond 
• If atoms have identical electronegativities and form a covalent bond, they will share the electrons equally and the bond will be nonpolar 
• If there is a large difference between the atoms and the electrons is almost completely transferred to the nonmetal, it is an ionic bond 
• Intermediate differences result in covalent bonds 
• Small difference ranges from 0-0.4 
• Intermediate is 0.4-2.0
• Large is 2.0+ 
• The polarity of a bond is quantified by the size of the dipole moment 
• A dipole moment will occur any time there is a separation between a positive and negative charge 
• The debye (D) is used as the dipole moment unit 
• The percent ionic character is the ratio of a bonds actual dipole moment to the dipole moment it would have if the electron was completely transferred 

9.7 Lewis Structures of Molecular Compounds and Polyatomic Ions

Writing Lewis Structures for Molecular Compounds 

• The first step is to write the correct skeletal structure. Hydrogen atoms can never be in the center. The more electronegative element should be in the center 
• Next calculate the total number of valence electrons of the compound
• Then distribute the electrons and create duets and octets where necessary. Give two electrons to each outer atom. Any lone pairs should be given to the middle atom 
• If you cannot form an octet for the middle atom, form double or triple bonds where necessary  

Writing Lewis Structures for Polyatomic Ions 

• Follow the same procedure, but be careful when counting electrons. If there is a negative charge, add an electron. If there is a positive charge subtract an electron. You will add and subtract the number of electrons listed in the charge 
• After drawing the structure, surround it with brackets and put the charge in the upper-right-hand corner 

9.8 Resonance and Formal Charge

Resonance 

• For some molecules, there is more than one valid Lewis Structure 
• The structures can change because the location of bonds is different 
• It is important to represent both ways. This can be done by drawing resonance structures. This is done by drawing both structures with double headed arrows between them 
• The actual structure of the molecule is called a resonance hybrid 
• Resonance stabilization is the lowering of the potential energy of electrons through delocalization 
• Resonance stabilization stables many molecules 

Formal Charge 

• Formal charge is fictitious 
• Formal charge is assigned to each Lewis Structure
• Formal charge is the charge a Lewis structure would have if all bonding electrons were shared equally between the atoms 
• When calculating the charge, all effects of electronegativity are ignored 
• Formal charge can be found by finding the difference between the number of valence electrons in the atom and the number of electrons that it "owns" in a Lewis Structure 
• Formal charge allows us to distinguish the best (most stable) of the two competing resonance structures 
• The sum of all formal charges in a neutral molecule must be 0
• The sum of all formal charges in an ion must equal the charge of the ion
• Small formal charges are better than large ones on an individual atom 
• Negative formal charges reside on the most electronegative atom 
• The best structures will have the least electronegative atom in the middle 

9.9 Exceptions to the Octet Rule: Odd-Electron Species, Incomplete Octets, and Expanded Octets

• The octet rule has a few exceptions including odd-electrons species, incomplete octets, and expanded octets 

Odd-Electron Species 

• Free radicals are molecules and ions that have an odd number of electrons 
• Because there is an odd number of electrons, octets for both atoms cannot be formed 
• Lewis models are not sophisticated enough to handle every single type of molecule on earth 
• Free radicals can be expressed as Lewis structures because resonance shows the different places the electron could be located 
• Free radicals are typically unstable and reactive 

Incomplete Octets 

• Incomplete octets are molecules or ions with less than 8 electrons 
• The most important element with an incomplete octet is Boron. Boron has 6 electrons 
• There are no additional electrons to add to the bonding region 
• Trying to move lone pairs to fix the octet will give you a negative formal charge on one of the atoms, this is highly unfavorable 
• An octet could be filled during a chemical reaction 

Expanded Octets 

• Expanded octets are molecules or ions with more than 8 electrons 
• They usually have up to 12, but in some cases 14, electrons 
• Third-period elements and beyond are typically expanded octets because the d orbitals exist. Therefore there is room for the extra electrons to go 
• Expanded octets never occur in second-period elements 
• Expanding an octet can lead to a lower formal charge 

9.10 Bond Energies and Bond Lengths

• Bond energies are used to estimate the enthalpy changes of a reaction  

Bond Energy

• The bond energy is the energy required to break 1 mole of the bond in the gas phase
• Bond energies are always positive. This is because energy is required to break a bond  
• A bond that requires a lot of energy to break is "stronger" 
• A bond can have a different bond energy depending on what molecule it is in 
• The average bond energy is calculated by taking the bond energies of that bond is many compounds 
• Bond energies also depend on if it is a single, double, or triple bond 
• Triple bonds are the strongest and single bonds are the weakest 

Using Average Bond Energies to Estimate Enthalpy Changes for Reactions 

• When bonds break, the process is endothermic. This gives a positive value 
• When bonds form, the process is exothermic, giving a negative value 
• The enthalpy of the reaction can be found by adding the two values together 
• A reaction is exothermic when weak bonds break and strong bonds form 
• A reaction is endothermic when strong bonds break and weak bonds form 
• Braking a chemical bond always requires energy 
• The forming of a chemical bond always releases energy 

Bond Lengths 

• The average bond length is the average length of the bond between two particular atoms in multiple compounds 
• Bond length depends on the atoms and the type of bond 
• Triple bonds are the shortest while single bonds are the longest 
• As a bond gets longer, it gets weaker (but of course there are a few exceptions to this trend) 

9.11 Bonding in Metals: The Electron Sea Model 

• Metals typically lose electrons 
• Metals have low ionization energies 
• Each metal donates one or more electrons to an electron sea when they bond together 
• Metals conduct electricity because the electrons are free to move. This also accounts for why metals are good at producing heat 
• Since there are no real bonds in metals, this accounts for the malleability and ductility of different metals