CH_10_6_per_page - 1 Dr Berk& Dr Koel CHM 25 Spring 2007 1 of 76 2 of 76 Contents 10-1 10-2 10-3 10-4 10-5 10-6 10-7 10-8 10-9 Lewis Theory An

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Unformatted text preview: 1 Dr. Berk & Dr. Koel CHM 25 Spring 2007 1 of 76 2 of 76 Contents 10-1 10-2 10-3 10-4 10-5 10-6 10-7 10-8 10-9 Lewis Theory: An Overview Covalent Bonding: An Introduction Polar Covalent Bonds Writing Lewis Structures Resonance Exceptions to the Octet Rule The Shapes of Molecules Bond Order and Bond Lengths Bond Energies Chapter 10: Chemical Bonding I: Basic Concepts Chemical Bonding Problems and questions — How is a molecule or polyatomic ion held together? Why are atoms distributed at strange angles? Why are molecules not flat? Can we predict the structure? How is structure related to chemical and physical properties? 3 of 76 Forms of Chemical Bonds 4 of 76 • There are 2 extreme forms of connecting or bonding atoms: •Ionic— complete transfer of 1 or more electrons from one atom to another •Covalent— some valence electrons shared between atoms •Most bonds are somewhere in between. Ionic Compounds Metal of low IE 5 of 76 Covalent Bonding The bond arises from the mutual attraction of 2 nuclei for the same electrons. Electron sharing results. 6 of 76 Nonmetal of high EA HA + H B HA HB The bond is a balance of attractive and repulsive forces. 2 Na(s) + Cl2(g) ---> 2 Na+ + 2 Cl- CHM 25 Sp07 CHM 25 Sp07 2 Bond Formation A bond can result from a “ head-to-head” overlap of atomic orbitals on neighboring atoms. • • • • • • 7 of 76 Chemical Bonding: Objectives Objectives are to understand: 1. valence e- distribution in molecules and ions. 2. molecular structures 3. bond properties and their effect on molecular properties. 8 of 76 H + Cl • • H Cl • • • • Overlap of H (1s) and Cl (2p) Note that each atom has a single, unpaired electron. 9 of 76 10 of 76 Lewis Theory: An Overview • Valence e- play a fundamental role in chemical bonding. • e- transfer leads to ionic bonds. • Sharing of e- leads to covalent bonds. • e- are transferred of shared to give each atom a noble gas configuration –the octet. Electron Distribution in Molecules Electron Distribution in Molecules • Electron distribution is depicted with Lewis electron dot structures • • • Valence electrons are distributed as shared or BOND PAIRS and unshared or LONE PAIRS. H Cl • • • • shared or bond pair lone pair (LP) This is called a LEWIS ELECTRON DOT structure. Valence Electrons Valence Electrons Electrons are divided between core and valence 11 of 76 12 of 76 Lewis Symbols •A chemical symbol represents the nucleus and the core e-. •Dots around the symbol represent valence e-. • •• Si • • • •N • • • • •P • • • • •Se • • • • • As • • • • I• • • • • electrons B 1s2 2s2 2p1 Core = [He] , valence = 2s2 2p1 • • •• Sb • • • • • •• Bi • Br [Ar] 3d10 4s2 4p5 Core = [Ar] 3d10 , valence = 4s2 4p5 •Al • • Ar • • CHM 25 Sp07 CHM 25 Sp07 • • • • 3 13 of 76 14 of 76 Ionic bonding • Exchange of electrons results in a pair of electrostatically attracted ions. This type of bond is thus called an ionic bond. Covalent bonding •Sharing of valence electrons results in a covalent bond. 15 of 76 Writing Lewis Structures •All the valence e- of atoms must appear. •Usually, the e- are paired. •Usually, each atom requires an octet. –H only requires 2 e-. Building a Dot Structure Ammonia, NH3 1. Decide on the central atom; never H. Central atom is atom of lowest affinity for electrons. Therefore, N is central 2. Count valence electrons H = 1 and N = 5 Total = (3 x 1) + 5 = 8 electrons / 4 pairs 16 of 76 •Multiple bonds may be needed. –Readily formed by C, N, O, S, and P. Building a Dot Structure 3. Form a single bond between the central atom and each surrounding atom 4. Remaining electrons form LONE PAIRS to complete octet as needed. 3 BOND PAIRS and 1 LONE PAIR. 17 of 76 HNH H • • Double and even triple bonds are commonly observed for C, N, P, O, and S 18 of 76 H2CO HNH H C2F4 • • SO3 Note that N has a share in 4 pairs (8 electrons), while H shares 1 pair. O • • C O • • • • CHM 25 Sp07 CHM 25 Sp07 4 19 of 76 20 of 76 Sulfur Dioxide, SO2 Sulfur Dioxide, SO2 1. Central atom = S 2. Valence electrons = 18 or 9 pairs • • • • • • • • Sulfur Dioxide, SO2 Sulfur Dioxide, SO2 bring in left pair • • • • • • OR bring in right pair • • O • • S O • • • • O • • • • S O • • • • • • • • • • This leads to the following structures. • • • • • • 3. Form double bond so that S has an octet — but note that there are two ways of doing this. bring in left pair OR bring in right pair • • • • • • • • • • O • • S O • • • • O S • • O • • O • • S O • • These equivalent structures are called RESONANCE STRUCTURES. The true electronic structure is a HYBRID of the two. 21 of 76 Formal Atom Charges • Atoms in molecules often bear a charge (+ or -). • The predominant resonance structure of a molecule is the one with charges as close to 0 as possible. Carbon Dioxide, CO2 2 22 of 76 +6 - ( 1 / 2 ) ( 4 ) - 4 •• • • •• = 0 • Formal charge = Group number – 1/2 (no. of bonding electrons) - (no. of LP electrons) O C O = • • +4 - ( 1 / 2 ) ( 8 ) - 0 0 23 of 76 24 of 76 Violations of the Octet Rule Usually occurs with B and elements of higher periods. Sulfur Tetrafluoride, SF4 • Central atom = • Valence electrons = ___ or ___ pairs. • Form sigma bonds and distribute electron pairs. • • • • • • • • F • • • • • • S • • F • • • • • • BF3 3 SF4 4 F F • • • • 5 pairs around the S 5 pairs around the S atom. A common atom. A common occurrence outside the occurrence outside the 2nd period. 2nd period. CHM 25 Sp07 CHM 25 Sp07 5 25 of 76 26 of 76 MOLECULAR GEOMETRY Electron Pair Geometries VSEPR •Valence Shell Electron Pair Repulsion theory. • Most important factor in determining geometry is relative repulsion between electron pairs. Molecule adopts Molecule adopts the shape that the shape that minimizes the minimizes the electron pair electron pair repulsions. repulsions. 27 of 76 28 of 76 Electron Pair Geometries Structure Determination by VSEPR • • Ammonia, NH3 H 1. Draw electron dot structure 2. Count BP’ and LP’ = 4 s s 3. The 4 electron pairs are at the corners of a tetrahedron. lone pair of electrons in tetrahedral position H H N H H N H Structure Determination by VSEPR Ammonia, NH3 There are 4 electron pairs at the corners of a tetrahedron. •• • • 29 of 76 Structure Determination by VSEPR Ammonia, NH3 The electron pair geometry is tetrahedral. N H H H lone pair of electrons in tetrahedral position 30 of 76 HNH H H N H H lone pair of electrons in tetrahedral position The ELECTRON PAIR GEOMETRY is tetrahedral. The MOLECULAR GEOMETRY — the The — the positions of the atoms — is PYRAMIDAL.. positions of the atoms — is CHM 25 Sp07 CHM 25 Sp07 6 31 of 76 Geometries for Four Electron Pairs Structure Determination by VSEPR Formaldehyde, CH2O 1. Draw electron dot structure • • 32 of 76 O C • • 2. Count BP’ and LP’ at C s s 3. There are 3 electron “ lumps” around C at the corners of a planar triangle. • • H H O C • • The electron pair geometry The electron pair geometry is PLANAR TRIGONAL with is PLANAR TRIGONAL with 120oo bond angles. 120 bond angles. H H 33 of 76 34 of 76 Structure Determination by VSEPR Formaldehyde, CH2O • • • • Structures with Central Atoms with More Than or Less Than 4 Electron Pairs O • • • • C H H The electron pair The electron pair geometry is PLANAR geometry is PLANAR TRIGONAL TRIGONAL The molecular geometry The molecular geometry is also planar trigonal.. is also planar trigonal Often occurs with Group 3A elements and with those of 3rd period and higher. 35 of 76 36 of 76 Boron Compounds • • Compounds with 5 or 6 Pairs Around the Central Atom • • Consider boron trifluoride, BF3 • • • • F 90 Þ F P F T rig o na l b ip y r a m id F F 12 0Þ F The B atom is surrounded by only • • • • B • • 3 electron pairs. Bond angles are 120o Bond angles are 120o Geometry described as F • • • • F 5 electron pairs planar trigonal CHM 25 Sp07 CHM 25 Sp07 7 37 of 76 38 of 76 Molecular Geometries for Five Electron Pairs All based on trigonal bipyramid Sulfur Tetrafluoride, SF4 Sulfur Tetrafluoride, SF4 • Number of valence electrons = 34 • Central atom = S • Dot structure • • • F • • • • • • • F • • • • S •• F •• • • • • • F• • • Electron pair geometry --> trigonal bipyramid (because there are 5 pairs around the S) 90Þ • • F S F F F 1 2 0Þ 39 of 76 40 of 76 Sulfur Tetrafluoride, SF4 Sulfur Tetrafluoride, SF4 Lone pair is in the equator because it requires more room. 90Þ 90Þ • • • • F F S S F F F F F F 1 2 0Þ 1 2 0Þ • • • F • • • • •• • F • • • • S •• F •• • • • • • F• • • Molecular Geometries for Six Electron Pairs All are based on the 8sided octahedron 41 of 76 Compounds with 5 or 6 Pairs Around the Central Atom 90 Þ F F F 6 electron pairs Bond Properties • What is the effect of bonding and structure on molecular properties? 42 of 76 F S Oc tahe dr on F F 90Þ Free rotation around C–C single bond No rotation around C=C double bond CHM 25 Sp07 CHM 25 Sp07 8 Bond Order 43 of 76 # of bonds between a pair of atoms Bond Order Bond Order 44 of 76 Double bond Fractional bond orders occur in molecules with resonance structures. Consider NO2• • • • Single bond N N • • •• • • • • • • O O• • O O • • • • • • • • The N— O bond order = 1.5 Acrylonitrile Triple bond Bond order = Total # of e - pairs used for a type of bond Total # of bonds of that type Bond order = 3 e - pairs in N — O bonds 2 N — O bonds Bond Order Bond Order Bond order is proportional to two important bond properties: 45 of 76 46 of 76 Bond Length • Bond length is the distance between the nuclei of two bonded atoms. (a) (b) bond strength bond length 414 kJ 123 pm 110 pm 745 kJ 47 of 76 48 of 76 Bond Length Bond length depends on size of bonded atoms. Bond Length Bond length depends on bond order. H— F H— Cl Bond distances measured in Bond distances measured in Angstrom units where 1 A = Angstrom units where 1 A = 10--2 pm. 10 2 pm. Bond distances measured in Bond distances measured in Angstrom units where 1 A = Angstrom units where 1 A = 10-22pm. 10- pm. H— I CHM 25 Sp07 CHM 25 Sp07 9 49 of 76 50 of 76 Bond Strength • — measured by the energy req’ to break a bond. d H— H C— C C=C CC NN Bond Strength • — measured by the energy req’ to break a bond. d BOND STRENGTH (kJ/mol) 436 346 602 835 945 The GREATER the number of bonds (bond order) the HIGHER the bond strength and the SHORTER the bond. 51 of 76 52 of 76 Using Bond Energies Estimate the energy of the reaction H— H(g) + Cl— Cl(g) ----> 2 H— Cl(g) Net energy = ∆Hrxn = = energy required to break bonds - energy evolved when bonds are made H— H = 436 kJ/mol H— H = 436 kJ/mol Cl— Cl = 242 kJ/mol Cl— Cl = 242 kJ/mol H— Cl = 432 kJ/mol H— Cl = 432 kJ/mol 53 of 76 Using Bond Energies Using Bond Energies Estimate the energy of the reaction H— H + Cl— Cl ----> 2 H— Cl H— H = 436 kJ/mol H— H = 436 kJ/mol Cl— Cl = 242 kJ/mol Cl— Cl = 242 kJ/mol H— Cl = 432 kJ/mol H— Cl = 432 kJ/mol Molecular Polarity Water Boiling point = 100 ˚C 54 of 76 Methane Boiling point = -161 ˚C Sum of H-H + Cl-Cl bond energies = 436 kJ + 242 kJ = +678 kJ 2 mol H-Cl bond energies = 864 kJ Net = ∆H = +678 kJ - 864 kJ = -186 kJ Why do ionic compounds dissolve in water? Why do water and methane differ so much in their boiling points? CHM 25 Sp07 CHM 25 Sp07 10 55 of 76 56 of 76 Polar Covalent Bonds and Electrostatic Potential Maps Polar Molecules 57 of 76 58 of 76 Bond Polarity Bond Polarity HCl is POLAR because it has a positive end and a negative end. + - • • • • Bond Polarity Bond Polarity • Three molecules with polar, covalent bonds. • Each bond has one atom with a slight negative charge (-) and and another with a slight positive charge (+ ) H Cl • • Cl has a greater share in bonding electrons than does H. Cl has slight negative charge (-) and H has slight positive charge (+ ) + - • • H C l• • • • Bond Polarity Bond Polarity Due to the bond polarity, the H— Cl bond energy is GREATER than expected for a “ pure” covalent bond. ENERGY ENERGY 339 kJ/mol calc’ d 339 kJ/mol calc’ d 432 kJ/mol measured 432 kJ/mol measured 59 of 76 60 of 76 Electronegativity, (EN) is a measure of the ability of an atom in a molecule to attract electrons to itself. BOND BOND “pure” bond pure ” bond “ real bond real bond Concept proposed by Linus Pauling 1901-1994 ELECTRONEGATIVITY, . Difference = 92 kJ. This difference is Difference = 92 kJ. This difference is proportional to the difference in proportional to the difference in CHM 25 Sp07 CHM 25 Sp07 11 Linus Pauling, 1901-1994 61 of 76 62 of 76 Electronegativity The only person to receive two unshared Nobel prizes (for Peace and Chemistry). Chemistry areas: bonding, electronegativity, protein structure 63 of 76 Percent Ionic Character Bond Polarity Which bond is more polar (or DIPOLAR)? O— H O— F 3.5 - 2.1 3.5 - 4.0 1.4 0.5 OH is more polar than OF 64 of 76 - O + H + O - F and polarity is “ reversed.” Molecular Polarity Molecules— such as HI and H2O— can be POLAR (or dipolar). 65 of 76 66 of 76 Molecular Polarity The magnitude of the dipole is given in Debye units. Named for Peter Debye (1884 1966). Rec’ 1936 d Nobel prize for work on x-ray diffraction and dipole moments. They have a DIPOLE MOMENT. The polar HCl molecule will turn to align with an electric field. CHM 25 Sp07 CHM 25 Sp07 12 Dipole Moments 67 of 76 Molecular Polarity Molecules will be polar if a) bonds are polar AND b) the molecule is NOT “ symmetric” 68 of 76 Why are some molecules polar but others are not? All above are NOT polar 69 of 76 Polar or Nonpolar? Compare CO2 and H2O. Which one is polar? • • Carbon Dioxide 70 of 76 O • • C O • • • • • CO2 is NOT polar even though the CO bonds are polar. • CO2 is symmetrical. Positive C atom Positive C atom is reason CO22 is reason CO and H22O react to and H O react to give H22CO33 give H CO -0.75 +1.5 -0.75 71 of 76 72 of 76 Polar or Nonpolar? • Consider AB3 molecules: BF3, Cl2CO, and NH3. Consequences of H2O Polarity CHM 25 Sp07 CHM 25 Sp07 13 Is CH3F Polar? 3 73 of 76 74 of 76 CH4 … CCl4 Polar or Not? C— F bond is very polar. C— F bond is very polar. Molecule is not symmetrical and Molecule is not symmetrical and so is polar. so is polar. • Only CH4 and CCl4 are NOT polar. These are the only two molecules that are “ symmetrical.” Substituted Ethylene 75 of 76 Substituted Ethylene 76 of 76 • C— F bonds are MUCH more polar than C— H bonds. • Because both C— F bonds are on same side of molecule, molecule is POLAR. • C— F bonds are MUCH more polar than C— H bonds. • Because both C— F bonds are on opposing ends of molecule, molecule is NOT POLAR. CHM 25 Sp07 CHM 25 Sp07 ...
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This note was uploaded on 10/24/2009 for the course CHEM 025 taught by Professor X during the Spring '06 term at Lehigh University .

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