2 Models of Chemical Bonding 9.1 Atomic Properties and Chemical Bonds9.2 The Ionic Bonding Model9.3 The Covalent Bonding Model9.4 Between the Extremes: Electronegativity and Bond Polarity9.5 An Introduction to Metallic Bonding
3 A general comparison of metals and non-metals Figure 9.1A general comparison of metals and non-metals
4 Types of Chemical Bonding 1. Metal with non-metalelectron transfer and ionic bonding2. Non-metal with non-metalelectron sharing and covalent bonding (localized)3. Metal with metalelectron pooling and metallic bonding (delocalized)
6 Lewis Electron-Dot Symbols For Main Group elements:The group number gives the number of valence electrons.Place one dot per valence electron on each of the four sides of the element symbol.Pair the dots (electrons) until all of the valence electrons are used.Example:Nitrogen (N) is in Group 5A and therefore has 5 valence electrons.N:..N::N.:N.
7 Lewis electron-dot symbols for elements in Periods 2 and 3 Figure 9.3
8 General RulesFor a metal, the total number of dots equals the maximum numberof electrons it loses to form a cation.For a non-metal, the number of unpaired dots equals the numberof electrons that become paired either through electron gain orelectron sharing. The number of unpaired dots equals either thenegative charge of the anion an atom forms or the number ofcovalent bonds it forms.
9 The Ionic Bonding Model Involves the transfer of electrons from metal to non-metal to formions that come together in a solid ionic compoundThe Octet RuleWhen atoms bond, they lose, gain or share electrons to attain a filledouter shell of eight (or two) electronsIn ionic bonding, the total number of electrons lost by the metal atomsequals the total number of electrons gained by the non-metal atoms.
10 SAMPLE PROBLEM 9.1Depicting Ion FormationPROBLEM:Use partial orbital diagrams and Lewis symbols to depict the formation of Na+ and O2- ions from the atoms, and determine the formula of the compound.PLAN:Draw orbital diagrams for the atoms and then move electrons to make filled outer levels. It can be seen that two sodiums are needed for each oxygen.SOLUTION:2s2pO2-3s3pNa2s2pO2 Na+3s3pNa:Na+ O.2Na+ + O 2-:
11 Figure 9.4Three ways to represent the formation of Li+ and F- through electron transfer1. Electron configurationsLi 1s22s1+F 1s22s22p5Li+ 1s2+F- 1s22s22p62. Orbital diagramsLi+1s2s2pLi1s2s2pF1s2s2p+F-1s2s2p+3. Lewis electron-dot symbols.+F:LiLi++F -:
12 Ionic Bonding and Lattice Energy The electron transfer process is an endothermic process, but ionic compoundformation is an exothermic process.Li(g) Li+ + e IE1 = 520 kJF(g) + e F-(g) EA = -328 kJLi(g) F(g) Li+(g) + F-(g) IE1 + EA = 192 kJBut ∆Hfo for solid LiF = kJ/mol!Li+(g) + F-(g) LiF(g) ∆Ho = kJ(an exothermic process due to the attraction ofoppositely charged ions)
13 Even more energy is released when the gaseous ions coalesce into a crystalline solid. Thus….Li+(g) + F-(g) LiF(s) ∆Holattice of LiF = lattice energy = kJThe lattice energy is the enthalpy change that occurs when gaseousions coalesce into an ionic solid.How do we measure lattice energy experimentally? UseHess’s law in a Born-Haber cycle
14 The Born-Haber cycle for lithium fluoride Figure 9.6
15 Working the Numbers STEP 1: Enthalpy of Li atomization = 161 kJ STEP 2: 1/2 the bond energy of F2(g)= 0.5(159 kJ) = 79.5 kJSTEP 3: IE1 for Li(g) = 520 kJSTEP 4: EA of F(g) = -328 kJThe enthalpy change for the overall process, ∆Hfo, = -617 kJOnly the lattice energy is unknown, and it is equal to the enthalpychange of the overall process minus the sum of the above foursteps = kJ
16 Central PointIonic solids exist only because the lattice energy drives theenergetically unfavorable electron transfer.
17 Periodic Trends in Lattice Energy Coulomb’s Lawcharge A x charge Belectrostatic force adistance2But energy = force x distance. Therefore,charge A x charge Belectrostatic energy adistancecation charge x anion chargea DHolatticeelectrostatic energy acation radius + anion radius
19 Effect of Ionic Charge on Lattice Energy Compare LiF and MgO: Li+ and Mg2+ have similar radii, andF- and O2- have similar radii.∆Holattice (LiF) = kJ/mol ∆Holattice (MgO) = kJ/molThe nearly four-fold larger value for MgO reflects the difference inthe product of the charges (12 vs 22) in the numerator of theelectrostatic energy equation (monovalent vs divalent ions).
20 Does the ionic model explain the properties of ionic compounds?
21 Electrostatic forces and the reason ionic compounds crack Figure 9.8
22 Electrical Conductance and Ion Mobility Molten ionic compoundSolid ionic compoundIonic compound dissolved in waterElectrical Conductance and Ion MobilityFigure 9.9
23 Table 9.1 Melting and Boiling Points of Some Ionic Compounds mp (oC)bp (oC)CsBr6361300NaI6611304MgCl27141412KBr7341435CaCl2782>1600NaCl8011413LiF8451676KF8581505MgO28523600
25 The Covalent Bonding Model Each atom in a covalent bond “counts” the shared electronsas belonging entirely to itself.An electron pair that is part of an atom’s valence shell but notinvolved in bonding is called a lone pair, or unshared pair.Bond order: the number of electron pairs being shared betweenany two bonded atomssingle bond (H2) - bond order of 1double bond (H2C=CH2) - bond order of 2triple bond (N2) - bond order of 3
27 The attractive and repulsive forces in covalent bonding Figure 9.12
28 Properties of Covalent Bonds Bond energy (bond enthalpy or bond strength): the energy requiredto overcome the mutual attraction between the bonded nuclei andthe shared electrons.Bond breakage is an endothermic process; bond energy is alwayspositive.Bond formation is an exothermic process.
29 Bond LengthFor a given pair of atoms, a higher bond order results in ashorter bond length and a higher bond energy.A shorter bond is a stronger bond.
34 SAMPLE PROBLEM 9.2Comparing Bond Length and Bond StrengthPROBLEM:Using the periodic table, rank the bonds in each set in order of decreasing bond length and bond strength:(a) S - F, S - Br, S - Cl(b) C = O, C - O, C OPLAN:(a) Bond order =1 for all and sulfur is bonded to halogens; bond length should increase and bond strength should decrease with increasing atomic radius. (b) Similar atoms (C) are bonded but bond order changes; bond length decreases as bond order increases, and bond strength increases as bond order increases.SOLUTION:(a) Atomic size increases moving down a group.(b) Using bond orders we get:Bond length: S - Br > S - Cl > S - FBond length: C - O > C = O > C OBond strength: S - F > S - Cl > S - BrBond strength: C O > C = O > C - O
35 Properties of Covalent Compounds Weak forces between molecules, not the strong covalentbonds within each molecule, are responsible for thephysical properties of covalent compounds.Covalent compounds have relatively low melting and boilingpoints.Most covalent compounds are poor electrical conductors.
36 Strong forces within molecules, weak forces between them Strong covalent bonding forces within moleculesFigure 9.14Weak intermolecular forces between molecules
37 Network Covalent Solids No separate molecules; held together by covalent bonds thatextend throughout the samplequartz: melts at 1550 oC.diamond: melts at 3550 oC.These examples illustrate the strength of covalent bonds.
38 Covalent bonds of network covalent solids Figure 9.15
39 The Concept of Electronegativity (EN) Defined as the relative ability of a bonded atom to attractshared electrons (not the same as EA)Bond energy of H2 = 432 kJ/molBond energy of F2 = 159 kJ/molBond energy of HF = 565 kJ/mol, not 296 kJ/molThe stronger-than-expected HF bond is due to unequalsharing of electrons, with F bearing a partial negativecharge and H bearing a partial positive charge. Theattraction between the partial charges strengthens thebond.
40 The Pauling electronegativity (EN) scale Figure 9.16
41 Trends in Electronegativity In general, electronegativity is inversely related toatomic size.For main-group elements, EN generally increases upa group and across a period.Non-metals are more electronegative than metals.The least electronegative (most electropositive) non-radioactive element is Cs (lower left-hand corner ofthe Periodic Table).
43 Electronegativity and Oxidation Number (a) The more electronegative atom in a bond is assignedall of the shared electrons; the less electronegative atom isassigned none of the shared electrons.(b) Each atom in a bond is assigned all of its unsharedelectrons.(c) The oxidation number is given by:O.N. = # valence e- - (# shared e- + # unshared e-)e.g.: HCl: Cl more electronegative than H; has 7 valenceelectrons; has an O.N. of = -1H has 1 valence electron; has an O.N. of = +1
44 Polar Covalent Bonds and Bond Polarity Covalent bonds involving atoms with different electronegativities: generate partial (+) and (-) charges; defined as polar covalent bonds (e.g., HCl)Polar covalent bonds: depicted by a polar arrow ( ) thatpoints toward the negative poleH2 and F2: examples of nonpolar covalent bonds
45 SAMPLE PROBLEM 9.3Determining Bond Polarity from EN ValuesPROBLEM:(a) Use a polar arrow to indicate the polarity of each bond: N-H, F-N, I-Cl.(b) Rank the following bonds in order of increasing polarity: H-N, H-O, H-C.PLAN:(a) Use Figure 9.16 to find EN values; the arrow should point toward the negative end.(b) EN increases across a period.SOLUTION:(a) The EN of N = 3.0, H = 2.1, F = 4.0, I = 2.5, Cl = 3.0N - HF - NI - Cl(b) The order of increasing EN is C < N < O; all have an EN larger than that of H.H-C < H-N < H-O
46 Partial Ionic Character of Polar Covalent Bonds Related directly to the electronegativity difference (∆EN) betweenthe bonded atomsThe greater the ∆EN, the larger the partial charges and the higherthe partial ionic character (PIC).Thus LiF has more PIC than HF; HF has more PIC than F2.
47 Boundary ranges for classifying the ionic character of chemical bonds 3.0DEN2.0Boundary ranges for classifying the ionic character of chemical bonds0.0Figure 9.18
48 Percent ionic character as a function of electronegativity difference (DEN) Figure 9.19
49 Li F Charge density of the LiF molecule (an ionic compound) No bond has 100%ionic character; electronsharing occurs to some extentFigure 9.20
50 Ionic-To-Covalent Bonding Continuum Across a Period Consider bonding between a metal and non-metal in Period 3NaCl, MgCl2, AlCl3, SiCl4, PCl3, S2Cl2, and Cl2Increasing covalent character (decreasing ionic character)from NaCl to Cl2Underlying factor: As ∆EN becomes smaller, the bond becomesmore covalent.
51 Properties of the Period 3 chlorides Figure 9.21Properties of the Period 3 chlorides
52 Metallic Bonding The electron-sea model: all metal atoms in the sample contribute their valence electrons to form an “electronsea” that is delocalized throughout the substanceThe metal atoms are not held in place as rigidly as arethe ions of an ionic solid.
53 Table 9.5 Melting and Boiling Points of Some Metals elementmp (oC)bp (oC)lithium (Li)1801347tin (Sn)2322623aluminum (Al)6602467barium (Ba)7271850silver (Ag)9612155copper (Cu)10832570uranium (U)11303930
54 Melting points of the Group 1A and Group 2A elements Figure 9.23
55 The reason metals deform metal is deformedFigure 9.24
56 Infrared Spectroscopy Tools of the LaboratoryInfrared SpectroscopyFigure B9.1Some vibrational modes in general diatomic and triatomic molecules
57 Tools of the Laboratory Infrared SpectroscopySome vibrational modes in general diatomic and triatomic moleculesFigure B9.1
58 Some vibrational modes in general diatomic and triatomic molecules. Tools of the LaboratoryInfrared SpectroscopySome vibrational modes in general diatomic and triatomic molecules.Figure B9.1
59 Tools of the Laboratory Figure B9.2The infrared (IR) spectrum of acrylonitrile