Copyright©2000 by Houghton Mifflin Company. All rights reserved. 1 Bonds Forces that hold groups of atoms together and make them function as a unit.

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 2 Ionic Bonds 4 Formed from electrostatic attractions of closely packed, oppositely charged ions. 4 Formed when an atom that easily loses electrons reacts with one that has a high electron affinity.

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 3 Ionic Bonds Q 1 and Q 2 = numerical ion charges r = distance between ion centers (in nm)

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 4 Bond Energy 4 It is the energy required to break a bond. 4 It gives us information about the strength of a bonding interaction.

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 5 Bond Length The distance where the system energy is a minimum.

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 6 Figure 8.1 Interaction of Two H Atoms and the Energy Profile

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 7 Polarity A molecule, such as HF, that has a center of positive charge and a center of negative charge is said to be polar, or to have a dipole moment.

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 8 Figure 8.2 The Effect of an Electric Field on Hydrogen Fluoride Molecules

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 9 Electronegativity The ability of an atom in a molecule to attract shared electrons to itself.  = (H  X) actual  (H  X) expected

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 10 Figure 8.4 Dipole Moment for H 2 O

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 11 Figure 8.3 The Pauling Electronegativity Values

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 12 Figure 8.5 Dipole Moment for NH 3

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 13 Figure 8.6 (a) Carbon Dioxide (b) Opposed Bond Polarities

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 14 Achieving Noble Gas Electron Configurations (NGEC) Two nonmetals react: They share electrons to achieve NGEC. A nonmetal and a representative group metal react (ionic compound): The valence orbitals of the metal are emptied to achieve NGEC. The valence electron configuration of the nonmetal achieves NGEC.

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 15 Isoelectronic Ions Ions containing the the same number of electrons (O 2 , F , Na +, Mg 2+, Al 3+ ) O 2  > F  > Na + > Mg 2+ > Al 3+ largest smallest

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 16 Figure 8.7 Sizes of Ions Related to Positions of the Elements in the Periodic Table

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 17 Lattice Energy The change in energy when separated gaseous ions are packed together to form an ionic solid. M + (g) + X  (g)  MX(s) Lattice energy is negative (exothermic) from the point of view of the system.

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 18 Formation of an Ionic Solid 1.Sublimation of the solid metal M(s)  M(g) [endothermic] 2.Ionization of the metal atoms M(g)  M + (g) + e  [endothermic] 3.Dissociation of the nonmetal 1 / 2 X 2 (g)  X(g) [endothermic]

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 19 Formation of an Ionic Solid (continued) 4.Formation of X  ions in the gas phase: X(g) + e   X  (g) [endo/exo] 5.Formation of the solid MX M + (g) + X  (g)  MX(s) [quite exothermic]

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 20 Figure 8.8 The Energy Changes Involved in the Formation of Solid Lithium Fluoride from its Elements

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 21 Figure 8.9 The Structure of Lithium Fluoride

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 22 Figure 8.10 Comparison of the Energy Changes Involved in the Formation of Solid Sodium Fluoride and Solid Magnesium Oxide

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 23 Q 1, Q 2 = charges on the ions r = shortest distance between centers of the cations and anions k = a fudge factor (accounts for e - configuration and the crystal packing arrangement).

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 24 Figure 8.11 Three Possible Types of Bonds

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 25 Figure 8.12 The Relationship Between the Ionic Character of a Covalent Bond and the Electronegativity Difference of the Bonded Atoms

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 26 Models Models are attempts to explain how nature operates on the microscopic level based on experiences in the macroscopic world.

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 27 Fundamental Properties of Models 4 A model does not equal reality. 4 Models are oversimplifications, and are therefore often wrong. 4 Models become more complicated as they age. 4 We must understand the underlying assumptions in a model so that we don’t misuse it.

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 28 Bond Energies Bond breaking requires energy (endothermic). Bond formation releases energy (exothermic).  H =  D( bonds broken )   D( bonds formed ) energy requiredenergy released

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 29 Localized Electron Model A molecule is composed of atoms that are bound together by sharing pairs of electrons using the atomic orbitals of the bound atoms.

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 30 Figure 8.14 The Molecular Structure of Methane

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 31 Localized Electron Model 1.Description of valence electron arrangement (Lewis structure). 2.Prediction of geometry (VSEPR model). 3.Description of atomic orbital types used to share electrons or hold lone pairs.

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 32 Lewis Structure 4 Shows how valence electrons are arranged among atoms in a molecule. 4 Reflects central idea that stability of a compound relates to noble gas electron configuration.

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 33 Comments About the Octet Rule 4 2nd row elements C, N, O, F must observe the octet rule. 4 2nd row elements B and Be often have fewer than 8 electrons around themselves - they are very reactive. 4 3rd row and heavier elements CAN exceed the octet rule using empty valence d orbitals. 4 When writing Lewis structures, satisfy octets first, then place electrons around elements having available d orbitals.

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 34 Resonance Occurs when more than one valid Lewis structure can be written for a particular molecule. These are resonance structures. The actual structure is an average of the resonance structures.

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 35 Formal Charge It is the difference between the number of valence electrons (VE) on the free atom and the number assigned to the atom in the molecule. We need: 1.# VE on free neutral atom 2.# VE “belonging” to the atom in the molecule

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 36 Formal Charge #1 is easy to find.For #2, VE assigned = (# lone pair e - ) + ½ (# shared e - ) Consider the O atoms in sulfate, p All the same environment (single bonds), so, VE assigned to each O atom = 6 + ½(2) = 7 Then, Formal charge on each O: 6 – 7 = -1 Formal charge on S: 6– 4 = 2

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 37 Formal Charge Now see the second possible Lewis structure, p. 388 (with the two double bonds). For singly-bonded O, 6 – 7 = -1 For doubly-bonded O, 6 – 6 = 0, And for S, 6 – 6 = 0. So, which set of formal charges mo’ betta’?

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 38 Formal Charge Two fundamental assumptions about formal charges are used to evaluate valid possible Lewis structures. 1)Atoms in molecules are most stable when formal charges are as close to zero as possible. 2)Any negative formal charges are assigned to the most electronegative atom(s). See the rules, p. 389.

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 39 Formal Charge Not as good Mo’ Betta’

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 40 VSEPR Model The structure around a given atom is determined principally by minimizing electron pair repulsions.

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 41 Figure 8.14 The Molecular Structure of Methane

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 42 Figure 8.15 The Molecular Structure of NH 3

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 43 Figure 8.16 The Molecular Structure of H 2 O

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 44 Figure 8.17 The Bond Angles in the CH 4, NH 3, and H 2 O Molecules

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 45 Figure 8.18 Electrons in Bonding

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 46 Figure 8.19 Possible Electron Pair Arrange ments for XeF 4

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 47 Figure 8.20 Three Possible Arrangements of the Electron Pairs in the I 3 - Ion

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 48 Figure 8.21 The Molecular Structure of Methanol

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 49 Predicting a VSEPR Structure 1.Draw Lewis structure. 2.Put pairs as far apart as possible. 3.Determine positions of atoms from the way electron pairs are shared. 4.Determine the name of molecular structure from positions of the atoms.