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1/22/20141 Formation of Covalent Bonds Two different theories which attempt to explain covalent molecular/ionic structure/shape Localized Electron (LE)

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Presentation on theme: "1/22/20141 Formation of Covalent Bonds Two different theories which attempt to explain covalent molecular/ionic structure/shape Localized Electron (LE)"— Presentation transcript:

1 1/22/20141 Formation of Covalent Bonds Two different theories which attempt to explain covalent molecular/ionic structure/shape Localized Electron (LE) Model-electron pairs are still localized around specific atoms, but orbitals around central atom are modified Molecular Orbital (MO) Model-all electrons in molecule are combined into set of molecular orbitals which describe bonding in entire molecule

2 Why hybrid orbitals? VSEPR model does good job at predicting molecular shape, despite fact that it has no obvious relationship to filling and shapes of atomic orbitals Based on shapes and orientations of 2s/2p orbitals on carbon atom, not obvious why CH 4 molecule should have tetrahedral geometry Used to reconcile covalent bonds formed from overlap of atomic orbitals with molecular geometries from VSEPR model 2 1/22/2014

3 Hybridization/Hybrid orbitals Hybridization (orbitals in covalently bonded atoms) Mixing of 2/more atomic orbitals of similar energies on same atom to produce new orbitals of equal energies Hybrid orbitals Valence bond theory creates hybrid orbitals that are linear combinations of s/p orbitals in valence shell (d if necessary) # atomic orbits = # hybrid orbitals Each hybrid orbital equivalent to others but large lobes point in different directions Atomic orbitals not used to make hybrids unaffected Mixtures of atomic orbitals with intermediate energy 3 1/22/2014

4 4 One 2s electron promoted to empty (2p) orbital 2 occupied orbitals blend to form 2 sp hybrid orbitals/2 remaining p orbitals unchanged No unpaired electrons

5 sp has 50% s/50% p character 2 sp hybrids point in opposite directions at 180 o to each other Require energy to promote 2s 2p orbital Large lobe of hybrid orbital can be directed at other atoms better than unhybridized atomic orbital Overlap more strongly/stronger bonds result Energy released by bond formation offsets energy expended to promote electrons Each sp hybrid involved in s bond/remaining p orbitals forms 2p bonds (All contain single unpaired electron) 5 1/22/2014

6 6 +

7 7 2s electron promoted to 2p orbital Have 3 unpaired electrons that form 3 sp 2 orbitals Creates 3 identical orbitals of intermediate energy/length Leaves one unhybridized p orbital

8 8 1/22/ Large lobes of orbitals lie in plane at angles of 120 o and point toward corners of triangle

9 9 1/22/2014 Promotion of 2s electron to 2p orbital results in valence shell with 4 unpaired electrons in four sp 3 hybrid orbitals. 4 orbitals form one 2s/three 2p orbitals (s 1 p 3 ) tetrahedral Four sp 3 orbitals identical in shape

10 10 1/22/2014

11 D-orbital hybridization Central atoms located in Period 3 and above can use empty d orbitals to receive promoted s electron 1/22/201411

12 12 1/22/2014 dsp 3 Hybridization 5 effective pairs around central atom Trigonal bipyramidal shape Lobes have bond angles of 90 o & 120 O PCl 5 example

13 13 1/22/2014 Each dsp 3 orbital also has small lobe not shown in diagram Phosphorus uses set of 5 dsp 3 orbitals to share electron pairs with sp 3 orbitals on 5 chlorine atoms Other sp 3 orbitals on each chlorine atom hold lone pairs

14 14 1/22/2014 d 2 sp 3 Hybridization Six effective pairs around central atom Octahedral structure Lobes have angles of 90 o SF 6 example

15 15 1/22/2014 Atomic Hybrid orbital setorbital set_ s/p 2 sp s/p/p 3 sp 2 s/p/p/p 4 sp 3 s/p/p/p/d 5 sp 3 d s/p/p/p/d/d 6 sp 3 d 2 Examples BeF 2, HgCl 2 BF 3, SO 3 CH 4, NH 3, H O, NH 4 + PF 5, SF 4, BrF 3, SbCl 5 2- SF 6, ClF 5, XeF 4, PF 4 -

16 16 1/22/2014 Basic DerivedHybridBondingNonbonding Structure structuree- pairse- pairs Linearsp20 Trigonal planar sp 2 30 Trigonal planar Bent sp 2 21 Tetrahedronsp 3 40 Tetrahedron Triangular pyramidsp 3 31 Tetrahedron Bentsp 3 22 Trigonal bipyramidsp 3 d50 Trigonal bipyramid Distorted tetrahedronsp 3 d41 Trigonal bipyramid T-shapesp 3 d32 Trigonal bipyramid Linearsp 3 d23 Octahedronsp 3 d 2 60 Octahedron Square pyramidsp 3 d 2 51 Octahedron Square planarsp 3 d 2 42

17 Strength of sigma bonds p-p > p-s > s-s p-orbitals allow overlap to greater extent as compared to p-s which is larger as compared to s-s overlap 17 1/22/2014 (s/single /δ) (Head-to-head overlap) Lobes of bonding orbital point toward each other. Overlap of two S orbitals to form sigma bond (green) Overlap of two P orbitals to form sigma bond (green)

18 Maximum electron density lies along bond (electron pair shared in area centered on line connecting nuclei) Line joining 2 nuclei passes through middle of overlap region (between nuclei) Maximum overlap forms strongest- possible sigma bond Atoms arrange themselves to give greatest-possible orbital overlap 18 1/22/2014

19 19 1/22/2014 (p/ ) Axes parallel to each other but perpendicular to internuclear axis Occupies space above/below internuclear axis (imaginary line connecting nuclei of two atoms) Electron density zero along bond Atomic orbitals interact above and below nuclei (perpendicular) Formed only in addition to sigma bond Always present in molecules with double/triple bonds Occur only w/sp or sp 2 hybridization present on central atom, but not sp 3

20 20 1/22/2014 Pi bonds are superimposed on sigma bonds so they simply modify dimensions of molecule. Significantly less overlap between component p-orbitals due to parallel orientation Weaker than sigma bonds-electrons farther from nucleus, so more reactive

21 Multiple Bonds Double Bonds Consist of 1 s bond (overlap of 2 sp orbitals) and 1 p bond (overlap of 2 p orbitals) s bond where electron pair located directly between atoms p bond where shared pair occupies space above and below s bond Triple Bonds Consist of 1 s / 2 p bonds Side-to-side overlap makes p bond electrons more reactive Electron density no longer located on internuclear axis (one electron cloud above and one below) Bond is weaker-p orbitals do not overlap as much in p bond as s bond 21 1/22/2014

22 Double bonds Triple bonds 22 1/22/2014 Consists of sigma (green)/2 pi bonds (red)

23 23 1/22/2014 (a) sp hybridized nitrogen atom (b) s bond in N 2 molecule (c) 2 p bonds in N 2 are formed when electron pairs are shared between two sets of parallel p orbitals (d) Total bonding picture for N 2

24 24 1/22/2014 Carbon is unique sp 3 hybrid orbitals = single bonds sp 2 hybrid orbitals = double bonds sp hybrid orbitals = triple bonds

25 Carbon atom of methane (CH 4 ) Made up of 4 C-H sigma (σ) bonds Each hybrid sp 3 orbitals of carbon undergoes end-on overlap with s-orbitals of H atoms Tetrahedral geometry 25 1/22/2014

26 Carbon atoms of ethyne (acetylene - C 2 H 2 ) 2 C-H σ bonds, 1 C-C σ bond, 2 C-C π bonds Δ bonds (gray) are linear in arrangement Unhybridized p-orbitals (green/purple) interact with each other laterally, resulting in bond formation 26 1/22/2014

27 Carbon atoms of ethene (ethylene - C 2 H 4 ) 4 C-H σ bonds, 1 C-C σ bond, 1 C-C π bond Hybrid orbital overlap end-on with s-orbitals of H atoms (δ bond in gray) Unhybridized p-orbitals (purple) at right angles to plane of hybrids, overlap laterally (π bond-double) 27 1/22/2014

28 Formaldehyde has following Lewis structure: Describe it bonding in terms of appropriate hybridized/unhybridized orbitals VSEPR predicts trigonal planar geometry which suggests sp 2 hybrid orbitals on C 28 1/22/2014

29 Acetonitrile molecule Predict bond angles around each C Approximately 109° around left C and 180° on right C Give hybridization of each C sp 3, sp Determine total number of δ/π bonds 5 δ bonds and 2 π bonds 29 1/22/2014

30 30 1/22/2014 orbitals of sp hybridized carbon atom orbital arrangement for sp 2 hybridized oxygen atom hybrid orbitals in the CO 2 molecule (a) orbitals in carbon dioxide-carbon-oxygen double bonds each consist of one s bond and one p bond. (b) Lewis structure for carbon dioxide

31 31 1/22/2014 Determine the total number of sigma and pi bonds in each of the following: Using the simple Lewis structure, also determine the hybridization for each: CH 3 Cl 4 δ, 0 Π, sp 3 PH 3 2 δ, 2 Π, sp H 2 S 3 δ, 0 Π, sp 3 CO δ, 0 Π, sp 3

32 32 1/22/2014 SO δ, 0 Π, sp 3 CS 2 3 δ, 1 Π, sp 2 SiF 4 3 δ, 1 Π, sp 2 NO δ, 0 Π, sp 3 PO δ, 0 Π, sp 2 ClO δ, 0 Π, sp 3

33 VSEPR did not explain why bonds exist between atoms Pair of electrons attracted to both atomic nuclei Bond is formed As extent of overlap increases, strength of bond increases Electronic energy drops as atoms approach each other Begin to increase again when they become too close Optimum distance (observed bond distance) at which total energy is at minimum 33 1/22/2014

34 34 1/22/2014 Basic s, p x, p y, and p z orbitals unsatisfactory for two reasons Orbitals not directed in particular direction-tend to spread out in all directions Geometry of orbitals rarely consistent with molecular geometry Could not adequately explain fact that some molecules contain two equivalent bonds with bond order between that of single and double bonds

35 35 1/22/2014 Atomic orbitals explain bonding/ account for molecular geometries Mathematical descriptions of where electrons most likely found Obtained by solving Schrödinger equation As angular momentum (m l ) and energy of electron increases, it tends to reside in differently shaped orbitals Orbitals corresponding to three lowest energy states (s, p, and d, respectively)

36 36 1/22/2014 Localized Electron Model (Valence Bond Theory) Describes structure of covalent bonds (how bonding occurs) Atoms in molecule bond together w/shared electrons Lewis structure shows valence electron arrangement Use VSEPR model to predict molecular geometry Atoms use atomic orbitals to share electrons/hold lone pairs (new set of hybridized orbitals can form) Lone pairs Lone pairs: electron pairs localized on atom Bonding pairs Bonding pairs: electron pairs in space between atoms

37 Combine Lewiss notion of electron-pair bonds with atomic orbitals Lewis theory Lewis theory: covalent bonding occurs when atoms share electrons (concentrates electron density between nuclei) Valence-bond theory Valence-bond theory: buildup of electron density between two nuclei occurs when valence atomic orbital of one atom shares space (overlaps) with that of another atom Shortcomings of LE Model Electrons not actually localized Does not deal effectively w/molecules containing unpaired electrons Gives no direct information about bond energies 37 1/22/2014

38 Direct overlap (electron sharing) of two atomic orbitals Local view of bonding (how adjacent atoms share electrons) How 2 electrons of opposite spin share space between nuclei/form bond Paired electrons localized in specific internuclear spaces between bonded atoms or remain unshared (lone pairs) As bond is formed, paired electrons spread out over molecule to form final electron cloud surrounding nuclei Description confirmed by many chemistry/physics experiments, including actual Scanning Tunneling Microscope picture of p-orbital Electronic structure/geometry is best compromise between maximum overlap (electron-nucleus attraction) and repulsion (electron-electron/nucleus-nucleus) 38 1/22/2014

39 Bonds (δ-head on/π- sideways) made by overlap of atomic (s,p) or hybridized (sp 2 ) orbitals antibonding orbitals (δ*/π*) Formation of bonding orbital accompanied by formation of antibonding orbitals (δ*/π*) which remain unoccupied and does not contribute to structure of molecule δ bonds have cylindrical symmetry Formed between pair of atoms within molecule Increased electron density on internuclear axis Rotation around bond does not change overlap of contributing atomic orbitals Lower energy than π bonds 39 1/22/2014

40 π bonds are not cylindrically symmetric May cover more than 2 nuclei (resonance) Increased electron density above/below internuclear axis (not on axis itself) Rotation breaks bond Electrons not always shared equally (EN) Skeletal structures (Lewis structures) correspond to valence-bond model 40 1/22/2014

41 Resonance structures represent molecules not adequately described by single structure, because electrons shared by more than 2 nuclei Rules for drawing reasonable resonance structures All resonance structures must be valid Lewis structures In all possible resonance structures atomic nuclei must not change their positions All atoms must not change their hybridization Only electron distribution may be changed All resonance structures must have same # unpaired electrons All atoms involved in resonance (electron sharing)/atoms directly bonded to them must lie in (or nearly in) same plane 41 1/22/2014

42 42 1/22/2014 Homework: Read 9.1, pp Q pp , #12, 14, 16, 21, 22, 27, 28

43 1/22/ Molecular Orbital (MO) Theory (Model)

44 44 1/22/2014 Molecular Orbital Theory Explains distributions (organization of valence electrons) and energy of electrons in molecules Molecule is similar to atom (have distinct energy levels that electrons can populate) Takes global view of bonding (all electrons in molecule are needed to describe how bonding occurs) Useful for describing properties of compounds Bond energies, electron cloud distribution, and magnetic properties Solution to VB problems by creating new set of orbitals w/intermediate between those of basic orbitals used to construct them

45 Basic principles of MO Theory 2 atomic orbitals w/similar energies overlap to form 2 molecular orbitals Electrons from each element participating in bond occupy new molecular orbitals MOs delocalized over many atoms (dont directly correspond to specific bond as in VB theory) No hybridization (all available atomic orbitals mixed into multiple combination Molecular orbitals have different energies depending on type of overlap Bonding orbitals (lower energy than corresponding AO) Nonbonding orbitals (same energy as corresponding AO) Antibonding orbitals (higher energy than corresponding AO) 1/22/

46 Molecular orbitals (MOs) made of fractions of atomic orbitals All atoms in molecule provide atomic orbitals to make MOs, but not all atomic orbitals must participate in all MOs MO filled by all available electrons (2 per orbital), starting with lowest energy MO orbital π bonds perpendicular to δ bonds, so cant mix Reflects orbital geometry of molecule 46 1/22/2014

47 Principles for formation of MOs 1. # MOs formed = # atomic orbitals combined 2. Atomic orbitals combine best with other atomic orbitals of similar energy 3. How effectively 2 atomic orbitals combine proportional to their overlap As overlap increases, energy of bonding MO lowered, energy of antibonding MO raised Lower energy molecular orbitals fill first 4. Each MO accommodates 2 electrons w/opposite spins (Pauli exclusion principle) 5. When MOs of same energy populated, Hunds rule is followed (equal energy orbitals ½ filled before pairing up) 6. Electron in antibonding orbital cancels its corresponding electron in bonding orbital 47 1/22/2014

48 In atoms, electrons occupy atomic orbitals, In molecules they occupy molecular orbitals which surround molecule Two atomic orbitals combine to form two molecular orbitals, one bonding ( ) and one antibonding ( *) Each line in diagram represents orbital 1/22/

49 49 1/22/2014 Bonding Orbital MOs w/lower energy than its corresponding original atomic orbitals Promotes formation of stable bond High electron density along internuclear axis

50 50 1/22/2014 Antibonding Orbital MOs w/higher energy than its corresponding atomic orbitals Destabilizes (negative impact) formation of bond Much lower electron density along internuclear axis

51 51 1/22/2014 When two atomic orbitals overlap in side-by-side fashion, molecular orbitals called π molecular orbitals Antibonding molecular orbital is designated as π* molecular orbital

52 52 1/22/2014 Bond Order (BO) Describes nature of bond formed by molecular orbitals Refers to average number of bonds that atom makes in all of its bonds to other atoms Bond order above 0 considered stable because it has excess of bonding electrons If bond order = 0 (BE = ABE), species does not exist Larger bond order = Greater bond strength = Greater bond energy = Shorter bond length

53 53 1/22/2014 Pg , Sample 9.6 For the species O 2, O 2 +, and O 2 -, give the electron configuration and the bond order for each. Which has the strongest bond?

54 54 1/22/2014 Magnetism can be induced in some nonmagnetic materials when in presence of magnetic field 4 Paramagnetism: (oxygen-2 unpaired es) 4 Unpaired electrons 4 Attracted to induced magnetic field 4 Much stronger than diamagnetism 4 Diamagnetism 4 Paired electrons 4 Repelled from induced magnetic field 4 Much weaker than paramagnetism

55 55 1/22/2014 NO + ion has total of = 14 electrons to place in molecular orbitals as follows Calculate bond order of NO + ion ½(10 - 4) = 3 Is the NO + ion diamagnetic or paramagnetic? 0 unpaired electrons so ion is diamagnetic

56 56 1/22/2014 Hydrogen Atom (H) 1 bonding electron/0 antibonding electrons Stable (lower energy, greater stability) Bond order of ½ One unpaired electron- paramagnetic

57 57 1/22/ bonding electrons in δ 1s molecular orbital /0 antibonding electrons Stable Bond order of 1 Electrons in line w/2 nuclei, so are s molecular orbitals No unpaired electrons- diamagnetic Hydrogen Molecule (H 2 )

58 Does He 2 exist? He 2 has four electrons, two in s 1s orbital and two in s* 1s orbital Bond order of 0, so He 2 does not exist 58 1/22/2014

59 59 1/22/2014 Bonding in Homonuclear Diatomic Molecules (composed of two identical atoms) In order to participate in molecular orbitals, atomic orbitals must overlap in space Larger bond order is favored When molecular orbitals are formed from p orbitals, s orbitals are favored over p orbitals (stronger) Electrons are closer to nucleus = lower energy

60 60 1/22/2014 Molecular orbital energy-level diagrams, bond orders, bond energies, and bond lengths for diatomic molecules Note that for O 2 and F 2, 2p orbital lower in energy than the π 2p orbitals

61 61 1/22/2014 Bonding in Heteronuclear Diatomic Molecules (different atoms) For atoms adjacent to each other in periodic table Use molecular orbital diagrams for homonuclear molecules Significantly different atoms Each molecule must be examined individually Use only electrons that are going to be involved in bonding No universally accepted molecular orbital energy order

62 N 2 NO bond order = 3 bond order = 2.5 1/22/

63 63 1/22/2014 Outcomes of MO Model Strengths Correctly predicts relative bond strength and magnetism of simple diatomic molecules Accounts for bond polarity Correctly portrays electrons as being delocalized in polyatomic molecules Disadvantages Difficult to apply quantitatively to polyatomic molecules

64 64 1/22/2014 Combining Localized Electron and Molecular Orbital Model Resonance Attempt to draw localized electrons in structure in which electrons not localized s (δ) bonds can be described using localized electron model p (π) bonds (delocalized) must be described using molecular orbital model

65 65 1/22/2014 Benzene s bonds (C - H and C - C) are sp 2 hybridized Localized model p bonds result of remaining p orbitals above/ below plane of benzene ring Delocalizing gives stability

66 66 1/22/2014 Homework: Read , pp Q pp , #32, 33, 38, 40, 46 Do 1 additional exercise and 1 challenge problem Submit quizzes by to me: dahl/ace/launch_ace.html?folder_path=/chemistry/book_content/ _zumdahl/ace&layer=act&src=ch09_ace1.xml dahl/ace/launch_ace.html?folder_path=/chemistry/book_content/ _zumdahl/ace&layer=act&src=ch09_ace2.xml dahl/ace/launch_ace.html?folder_path=/chemistry/book_content/ _zumdahl/ace&layer=act&src=ch09_ace3.xml

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