Chapter 8 Chemical Bonding I: Basic Concepts

Chapter 8Chemical Bonding I:Basic ConceptsCopyright McGraw-Hill 20091

8.1 Lewis Dot Symbols Valence electrons determine anelement’s chemistry. Lewis dot symbols represent thevalence electrons of an atom as dotsarranged around the atomic symbol. Most useful for main-group elementsCopyright McGraw-Hill 20092

Lewis Dot Symbols of the Main Group ElementsCopyright McGraw-Hill 20093

Write Lewis dot symbols for the following:(a) N(b) S2 (c) K Copyright McGraw-Hill 20094

Write Lewis dot symbols for the following:(a) N(c) K S 2 (b)S2 N K Copyright McGraw-Hill 20095

8.2 Ionic Bonding Ionic bond: electrostatic force that holdsoppositely charge particles together Formed between cations and anions ExampleNa Cl Na Cl IE1 EA1 496 kJ/mol 349 kJ/mol 147 kJ/molm.p. 801oC f H – 410.9 kJ/molCopyright McGraw-Hill 20096

Microscopic View of NaCl FormationCopyright McGraw-Hill 20097

Lattice energy the energy required tocompletely separate one mole of a solid ioniccompound into gaseous ions -- -NaCl(s) Na (g) Cl (g) - - Hlattice 788 kJ/molBecause they are defined as an amount of energy,lattice energies are always positive.Copyright McGraw-Hill 20098

Coulombic attraction:Q1 Q2F d2Q1 Q amount of charged distance of separationQ2 d Lattice energy (like a coulombic force) depends on Magnitude of charges Distance between the chargesCopyright McGraw-Hill 20099

Lattice energies of alkali metal iodidesCopyright McGraw-Hill 200910

The ionic radii sums for LiF and MgO are 2.01 and 2.06 Å,respectively, yet their lattice energies are 1030 and 3795 kJ/mol.Why is the lattice energy of MgO nearly four times that of LiF?Copyright McGraw-Hill 200911

Born-Haber cycle: A method to determinelattice energiesCopyright McGraw-Hill 200912

Born-Haber cycle for CaOCa(s)#1 Ca(g)#2 Ca2 (g) (1/2) O2(g)#3 O(g)#4 O2 (g)#6 CaO(s)#5#1 Heat of sublimation Hf [Ca(g)] 178 kJ/mol#2 1st & 2nd ionization energies I1(Ca) I2(Ca) 1734.5 kJ/mol#3 (1/2) Bond enthalpy (1/2) D(O O) Hf [O(g)] 247.5 kJ/mol#4 1st & 2nd electron affinities EA1(O) EA2(O) 603 kJ/mol#5 (Lattice Energy) Hlattice[CaO(s)] (the unknown)#6 Standard enthalpy of formation Hf [CaO(s)] 635 kJ/mol Hlattice 3398 kJ/mol 178 1734.5 247.5 603 Hlatt 635Copyright McGraw-Hill 200913

8.3 Covalent Bonding Atoms share electrons to form covalentbonds. HHH H orH–H In forming the bond the atoms achieve amore stable electron configuration.Copyright McGraw-Hill 200914

Octet: Eight is a “magic” number of electrons. OctetRule: Atoms will gain, lose, orshare electrons to acquire eightvalence electronsExamples:Na Copyright McGraw-Hill 2009 HOH H H O Cl Na Cl 15

Lewis Structures HH–H Cl– Cl Cl Cl Cl Cl HH H Shared electrons BondsNon-bonding valence electrons Lone pairsCopyright McGraw-Hill 200916

Multiple Bonds- The number of shared electron pairs is the numberof bonds. Single Bond O C O O C O Double BondN N N N Cl– Cl Cl Cl Triple BondCopyright McGraw-Hill 200917

Bond strength and bond lengthbond strengthsingle double triplebond lengthsingle double tripleBond StrengthBond LengthN–NN NN N163 kJ/mol418 kJ/mol941 kJ/mol1.47 Å1.24 Å1.10 ÅCopyright McGraw-Hill 200918

8.4 Electronegativity andPolarity Nonpolar covalent bond electronsare shared equally by two bondedatoms Polar covalent bond electrons areshared unequally by two bonded atomsCopyright McGraw-Hill 200919

Electron density distributions red high electron densitygreen intermediate electron densityblue low electron density -H–FH–FalternaterepresentationsCopyright McGraw-Hill 200920

Electronegativity: ability of an atomto draw shared electrons to itself.- More electronegative elements attract electrons morestrongly. relative scale related to IE and EA unitless smallest electronegativity:Cs 0.7 largest electronegativity:F 4.0Copyright McGraw-Hill 200921

Electronegativity: The Pauling ScaleCopyright McGraw-Hill 200922

Variation in Electronegativity with Atomic NumberCopyright McGraw-Hill 200923

Polar and nonpolar bonds2.1 - 2.1 0.04.0 - 2.1 1.94.0 - 0.9 3.1nonpolarcovalentpolarcovalentionic 2.0 is ionicCopyright McGraw-Hill 200924

Dipole moments and partial charges- Polar bonds often result in polar molecules.- A polar molecule possesses a dipole.- dipole moment ( ) the quantitative measure of adipole - QrH–F Q –QrSI unit: coulomb meter (C m)common unit: debye (D)1 D 3.34 10 30 C mCopyright McGraw-Hill 2009HFHClHBrHI1.82 D1.08 D0.82 D0.44 D25

8.5 Drawing Lewis Structures1) Draw skeletal structure with the central atom beingthe least electronegative element.2) Sum the valence electrons. Add 1 electron for eachnegative charge and subtract 1 electron for eachpositive charge.3) Subtract 2 electrons for each bond in the skeletalstructure.4) Complete electron octets for atoms bonded to thecentral atom except for hydrogen.5) Place extra electrons on the central atom.6) Add multiple bonds if atoms lack an octet.Copyright McGraw-Hill 200926

What is the Lewis structure of NO3 ?–O1) Draw skeletal structure with centralatom being the leastO–N–Oelectronegative.2) Sum valence electrons. Add 1 for each negativecharge and subtract 1 for each positive charge.NO (1 5) (3 6) 1 24 valence e 324 e 6 e 4) Complete electron octets for atomsbonded to the central atom exceptfor hydrogen.18 e : ::O – N –O:: ––:O::: :Copyright McGraw-Hill 2009:O:: :5) Place extra electrons on thecentral atom.6) Add multiple bonds if atomslack an octet.: –3) Subtract 2 for each bond in the skeletal structure.:O – N O:24 e 27

Copyright McGraw-Hill 200928

8.6 Lewis Structures andFormal Charge The electron surplus or deficit, relative to the free atom,that is assigned to an atom in a Lewis structure.Formal ChargeTotal nonTotal bonding 11 bonding2 electronselectrons: :Example:TotalvalenceelectronsH2O H : O : HH: orig. valence e 1 non-bonding e 0 1/2 bonding e 1formal charge 0O: orig. valence e non-bonding e 1/2 bonding e formal charge 6 4 20Formal charges are not “real” charges.Copyright McGraw-Hill 200929

Example: Formal charges on the atoms in ozoneO O OO 6 4 12 4 OO 0 6 2 12 6 1 6 6 12 2 1Copyright McGraw-Hill 200930

Formal charge guidelines A Lewis structure with no formal charges isgenerally better than one with formal charges. Small formal charges are generally better thanlarge formal charges. Negative formal charges should be on themore electronegative atom(s).HExample:HCOHorHO?HAnswer: CC OHHCopyright McGraw-Hill 2009CH O 31

Identify the best structure for the isocyanate ion below:(a):C N O: 2(b)0:C N – O: 1(c) 1 1 1– 1:C – N O: 3–– 1Copyright McGraw-Hill 200932

Identify the best structure for the isocyanate ion below:(a):C N O: 2(b)0:C N – O: 1(c) 1 1 1– 1:C – N O: 3–– 1Copyright McGraw-Hill 200933

8.7 Resonance Resonance structures are used when two ormore equally valid Lewis structures can be written.::: :Example: NO2:O – N O:–These two bonds are known to be identical.: :::O – N O:–:::: :Solution::O N – O:–Two resonance structures, their average or theresonance hybrid, best describes the nitrite ion.The double-headed arrow indicates resonance.Copyright McGraw-Hill 200934

Additional ExamplesCarbonate: CO32 Benzene: C6H6orCopyright McGraw-Hill 200935

8.8 Exceptions to the OctetRule Exceptionsto the octet rule fall into threecategories: Molecules with an incomplete octet Molecules with an odd number ofelectrons Molecules with an expanded octetCopyright McGraw-Hill 200936

Incomplete OctetsExample: BF3 (boron trifluoride)BF3 (1 3) (3 7) 24 val. e :::F::F::–::–::F – B F: 1::::F – B – F:-1no octet Common with Be, B and Al compounds, but theyoften dimerize or polymerize.ClExample:BeClBeClCopyright McGraw-Hill 2009ClBeClBeCl37

Odd Numbers of ElectronsExample: NO (nitrogen monoxide or nitric oxide)NO (1 5) (1 6) 11 valence e 0. 0 1 1.Are these both:N O::N O:equally good?betterExample: NO2 (nitrogen dioxide)NO2 (1 5) (2 6) 17 val. e 000.0.:O N – O:00:O – N O:best0 1. 1:O N – O: 1 1.0:O – N O:Are these all equally good?Copyright McGraw-Hill 200938

Expanded Octet Elements of the 3rd period and beyond haved-orbitals that allow more than 8 valence electrons.:F:F:F:48 valence e (S has 12 valenceelectrons )F:::::F – Xe – F::XeF2 FS–SF6 –F22 valence e (Xe has 10 valenceelectrons)Copyright McGraw-Hill 200939

8.9 Bond Enthalpy Bond enthalpy is the energy associated with breaking aparticular bond in one mole of gaseous molecules. Bond enthalpy is one measure of molecular stability. Symbol: Ho For diatomic molecules these are accuratelymeasured quantities.Cl2(g) Cl(g) Cl(g) Ho 243.4 kJHCl(g) H(g) Cl(g) Ho 431.9 kJsingle bondsO2(g) O(g) O(g) Ho 495.0 kJdouble bondN2(g) N(g) N(g)triple bond Ho 945.4 kJCopyright McGraw-Hill 200940

Bond enthalpies for polyatomic molecules dependupon the bond’s environment.–––H–HH–C–H H–C HHH–––––H H––H HH – C– C – H H – C– C H–H H H 435 kJH H H 410 kJ6% less Average bond enthalpies are used for polyatomicmolecules. Provide only estimatesCopyright McGraw-Hill 200941

Prediction of bond enthalpyenthalpyatoms BE(p)BE(r)productsreactants Ho BE(reactants) BE(products)Copyright McGraw-Hill 200942

Example: Calculate the enthalpy of reaction forCH4(g) Br2(g) CH3Br(g) HBr(g)Solution:Consider ONLY bonds broken or formed.HH Br – Br ––H–C–HH – C – Br––HH H – Br Hrxn [BE(C–H) BE(Br–Br)] – [BE(C–Br) BE(H–Br)] [ (413) (193) ] – [ (276) (366) ] – 36 kJ/molCopyright McGraw-Hill 200943

Copyright McGraw-Hill 200944

Key Points Lewis dot symbolsIonic bondingLattice energyBorn-Haber cycleCovalent bondingOctet ruleLewis structuresBond orderBond polarityCopyright McGraw-Hill 200945

Key Points ElectronegativityDipole momentDrawing lewis structuresFormal chargeResonance structuresIncomplete octetsOdd numbers of electronsExpanded octetsBond enthalpyCopyright McGraw-Hill 200946

non-bonding e 0 1/2 bonding e 1 formal charge 0 O: orig. valence e 6 non-bonding e 4 1/2 bonding e 2 formal charge 0 Example: H 2 O H:O:: Total valence electrons Formal Charge Total non-bonding