• Document: Chapter Nine. Chemical Bonding I
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Chapter Nine Chemical Bonding I 1 The Ionic Bond and Lattice Energies 2 Lewis Dot Symbols 3 Consists of atomic symbol surrounded by 1 dot for each valence electron in the atom Only used for main group elements # valence electrons = group number 4 Ionic Bonding Electrons are transferred from 1 atom to another Metal atoms: Lose electrons to form cations Nonmetal atoms: Gain electrons to form anions Electrostatic force bonds ions into an ionic compound Form an ionic salt with repeating structure: NaCl, LiF Ionic Bonds follow the octet rule Atoms lose or gain valence e- to make an octet (8e-) 8 valence e- = Noble gas configuration Li Li+ + e- 1s22s1→ 1s2 e- + F F - 1s22s22p5 →1s22s22p6 Li+ + F - Li+ F - 5 The Born Haber Cycle Formation of a salt can be tracked by individual steps Enthalpies associated with each step are additive Lattice energies are exothermic: salt is more stable than ions (∆H) associated with gaseous atoms, so convert to gas Ionization energies: Electron loss Electron affinity: Electron gain Lattice energy: Energy released when forming a crystal Formation energy (∆Hf) not based on gaseous atoms Energy associated with forming 1 mole of a compound from its elements in their normal states Sum of energy transfers called “Born-Haber cycle” Can use the cycle to the enthalpy of any step in cycle 6 Energy Changes in the Born-Haber Cycle Li (s)  Li(g) ∆Hsubl = +155 kJ 1/2 F2(g)  F(g) ∆Hdiss = +75 kJ Li (g)  Li+(g)+ e- ∆H IE = + 520kJ F(g) + e-  F-(g) ∆HEA = - 328 kJ Li+ (g) + F-(g)  LiF (s) ∆Hlattice E = ? kJ Li(s) + 1/2F2(g)  LiF(s) ∆H °f = - 594 kJ Hess’s Law: ∆H °f =∆Hsubl+ ∆Hdiss + ∆H IE + ∆HEA + ∆Hlattice E -594kJ =155kJ + 75kJ + 520kJ - 328kJ + ∆Hlattice E ∆Hlattice E = -1016kJ 7 Lattice Energies of Common Salts The Covalent Bond and Electronegativity 8 9 Lewis Structures Lewis structures represent covalent bond formation Shared pairs of electrons give both atoms an octet F + F F F 1s22s22p5 →1s22s22p6 7e- 7e- 8e- 8e- Bonding Pairs: Shared electrons count for both atoms Represented by a dash (-) between bonding atoms Lone Pairs: Non-shared electrons count for 1 atom Represented by a pair of dots (••) around atom lone pairs F F lone pairs single covalent bond 10 Multiple Bonds More than one pair of electrons is shared between atoms so each atom can form an octet. Single Bond: 1 shared pair: 1 dash (-) Double bonds: 2 shared pairs: 2 dashes (=) Triple bonds: 3 shared pairs: 3 dashes (≡) Allows atoms in the Lewis structure to share extra electrons if there are not enough for the central atom O C O N N double bonds triple bond 11 Electronegativity and Polar Covalent Bonds Electronegativity The ability of an atom to attract electrons F is the most electronegative atom H F Nonmetals high electronegativies Polar Covalent Bonds Differences in electronegativity result in unequal sharing of electrons between atoms More electronegative atom has a partial negative charge Percent Ionic Character Measure of polarity of bond 100% ionic is full transfer of electron, no sharing 100% covalent is equal sharing, H2, Cl2, etc. 12 The Electronegativities of Common Elements 9.5 Writing Lewis Structures 13 14 Writing Lewis Structures: General Information Electronegativity Central atom usually has the lowest electronegativity (atom lower or to the left in periodic table) Terminal atoms (except H) have higher electronegativities Terminal Atoms Bonded to only one other atom Hydrogen atoms are terminal atoms Bonding Hydrogen atoms are bonded to oxygen atoms in oxoacids Make the molecule as symmetrical as possible 15 Write the Lewis

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