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Chapter 10 Review

8 Pages

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CHEM 110
Ariel Fenster

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th October 6 , 2011 CHAPTER 10 Chemical Bonding 10.1 Lewis Theory: 1. Electrons (especially valence electrons) play a fundamental role in chemical bonding 2. Electrons are transferred from one atom to another. Positive and negative ions are formed and attract each other through electrostatic forces called ionic bonds. 3. One or more pairs of electrons are shared between atoms, this is called a covalent bond. 4. Electrons are transferred or shared in such a way that they may acquire a stable electron configuration or an octet. Binary ionic compounds: consist of monatomic cations and monatomic anions Ternary ionic compounds: consist of monatomic and polyatomic ions Bonding between atoms within the polyatomic ions are covalent Rules for Writing Lewis Structures: 1. Hydrogen atoms are always terminal atoms 2. Central atoms are generally those with the lowest EN 3. Carbon atoms are always central atoms 4. Molecules and polyatomic ions generally have compact, symmetrical structures 10.2 Types of Bonding: Ionic: transfer of electrons from one atom to another Cation and anion pair Positive charge is smaller than the negative charge Examples: NaCl, KBr, MgO (salts) Covalent: sharing of electrons between atoms Example: H ,2CH ,4H 2, graphite, diamond Bond Pair: a pair of electrons in a covalent bond Lone pair: electron pairs that are not involved in bonding Coordinate covalent bond: a single atom contributes both of the electrons to a shaied pair Multiple Covalent Bonds: double and triple covalent bonds Triple bonds are very strong and difficult to break Metallic Bonding: delocalization of electrons between atoms Good conductions Example: Zn, Cu, Fe, metal alloys Pure Covalent Bonds: both atoms share equally the bonding electrons Polar Covalent Bonds: electrons unequally shared Example: HF Hydrogen Fluoride: Fluorine has a greater electron affinity than hydrogen, so there is a greater electron density around fluorine o Sharing of electrons, however bond is polarized o Greater pull for electrons due to more nuclei and protons in F th October 6 , 2011 CHAPTER 10 Chemical Bonding Dipole Moment: polar molecules are characterized by a dipole moment, directed from the positive to the negative end of the molecules (permanent dipole and unequal change in distribution) Dipole moment = partial charge x distance Degree of Polarity is determined by the Difference in Electronegativity 10.3 Electrostatic Potential Map: A way to visualize the charge distribution within a molecule; the work done in moving a unit of positive charge at a constant speed from one region of a molecule to another Red represents the most negative electrostatic potential Blue represents the most positive electrostatic potential The potential increases from red, through yellow, to blue Electronegativity Scales: Mulliken: based on ionization energy and electron affinity Pauling: based on bond energy calculations Electronegativity: an atoms ability to compete for electrons with other atoms to which it is bonded Large EN differences = more metallic and non metallic elements (bonds are essentially ionic) Small EN differences = two nonmetal atoms (bonds are essentially covalent) Trends in Electronegativity: Difference in EN Type of Bonding |a -b | < 0.5 Pure covalent 0.5 < |a -b | < 1.7 Polar Covalent |a -b | > 1.7 Ionic 1.7 is arbitrary!October 6 , 2011 CHAPTER 10 Chemical Bonding 10.4 Formal Charges How to calculate a formal charge? FC= # of valence electrons - # of lone pairs number of bonded pairs (x ) Rules for Formal Charges: 1. Sum of the formal charges in a Lewis structure must equal zero for a neutral atom Exceptions to the Octet Rule and must equal theBFa3nitude of the charge for a polyatomic ion. 2. They should be as small as possible 1. Incomplete Octet 3. Negative formal charges usually appear on the most EN atoms; positive formal charges are on the least electronegative atoms 4. Structures having formal charges of the same sign on adjacent atoms are unlikely 10.5 Resonance: The term use to describe the existence of two or more forms of an element that differ in their bonding and molecular structure is allotropy (2.e. O3and O ) Resonance: two or more plausible Lewis structures contribute to the correct structure Same skeletal structure (atomic positions cannot change) but they can differ only in how many electrons are distributed within the structure 10.6 Exceptions to the Octet Rule BF 3 Odd-Electron SpeciBF:3 If the number of valence electrons in a Lewis structure is odd, there must be an unpaired
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