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Formation of Covalent Bonds

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1 Formation of Covalent Bonds
Localized Electron (LE) Model-electron pairs are still localized around specific atoms, but orbitals around central atom are modified Formation of Covalent Bonds Two different theories which attempt to explain covalent molecular/ionic structure/shape 3/25/2017 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 CH4 molecule should have tetrahedral geometry Used to reconcile covalent bonds formed from overlap of atomic orbitals with molecular geometries from VSEPR model 3/25/2017

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 3/25/2017

4 One 2s electron promoted to empty (2p) orbital
No unpaired electrons One 2s electron promoted to empty (2p) orbital 2 occupied orbitals blend to form 2 sp hybrid orbitals/2 remaining p orbitals unchanged 3/25/2017

5 2 sp hybrids point in opposite directions at 180o to each other
sp has 50% s/50% p character 2 sp hybrids point in opposite directions at 180o 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) 3/25/2017

6 + 3/25/2017

7 2s electron promoted to 2p orbital
2s electron promoted to 2p orbital Have 3 unpaired electrons that form 3 sp2 orbitals Creates 3 identical orbitals of intermediate energy/length Leaves one unhybridized p orbital 3/25/2017

8 Large lobes of orbitals lie in plane at angles of 120o and point toward corners of triangle
+ 3/25/2017

9 Four sp3 orbitals identical in shape
Promotion of 2s electron to 2p orbital results in valence shell with 4 unpaired electrons in four sp3 hybrid orbitals. Four sp3 orbitals identical in shape tetrahedral 4 orbitals form one 2s/three 2p orbitals (s1p3) 3/25/2017

10 3/25/2017

11 D-orbital hybridization
Central atoms located in Period 3 and above can use empty d orbitals to receive promoted s electron 3/25/2017

12 dsp3 Hybridization 5 effective pairs around central atom
Trigonal bipyramidal shape Lobes have bond angles of 90o & 120O PCl5 example 3/25/2017

13 Each dsp3 orbital also has small lobe not shown in diagram
Phosphorus uses set of 5 dsp3 orbitals to share electron pairs with sp3 orbitals on 5 chlorine atoms Other sp3 orbitals on each chlorine atom hold lone pairs 3/25/2017

14 d2sp3 Hybridization Six effective pairs around central atom
Octahedral structure Lobes have angles of 90o SF6 example 3/25/2017

15 Atomic Hybrid orbital set orbital set_ s/p 2 sp s/p/p 3 sp2 s/p/p/p 4 sp3 s/p/p/p/d 5 sp3d s/p/p/p/d/d 6 sp3d2 Examples BeF2, HgCl2 BF3, SO3 CH4, NH3, H O, NH4+ PF5, SF4, BrF3, SbCl52- SF6, ClF5, XeF4, PF4- 3/25/2017

16 Trigonal planar Bent sp2 2 1 Tetrahedron sp3 4 0
Basic Derived Hybrid Bonding Nonbonding Structure structure e- pairs e- pairs Linear sp 2 0 Trigonal planar sp2 3 0 Trigonal planar Bent sp2 2 1 Tetrahedron sp3 4 0 Tetrahedron Triangular pyramid sp3 3 1 Tetrahedron Bent sp3 2 2 Trigonal bipyramid sp3d 5 0 Trigonal bipyramid Distorted tetrahedron sp3d 4 1 Trigonal bipyramid T-shape sp3d 3 2 Trigonal bipyramid Linear sp3d 2 3 Octahedron sp3d2 6 0 Octahedron Square pyramid sp3d2 5 1 Octahedron Square planar sp3d2 4 2 3/25/2017

17 Strength of sigma bonds p-p > p-s > s-s
(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) 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 3/25/2017

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 3/25/2017

19 (p/) (perpendicular) 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 Formed only in addition to sigma bond Always present in molecules with double/triple bonds Occur only w/sp or sp2 hybridization present on central atom, but not sp3 3/25/2017

20 Significantly less overlap between component p-orbitals due to parallel orientation
Weaker than sigma bonds-electrons farther from nucleus, so more reactive Pi bonds are superimposed on sigma bonds so they simply modify dimensions of molecule. 3/25/2017

21 Multiple Bonds Double Bonds Triple 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 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 3/25/2017

22 Consists of sigma (green)/2 pi bonds (red)
Double bonds Triple bonds Consists of sigma (green)/2 pi bonds (red) 3/25/2017

23 sp hybridized nitrogen atom s bond in N2 molecule
2 p bonds in N2 are formed when electron pairs are shared between two sets of parallel p orbitals Total bonding picture for N2 3/25/2017

24 Carbon is unique sp3 hybrid orbitals = single bonds
sp2 hybrid orbitals = double bonds sp hybrid orbitals = triple bonds 3/25/2017

25 Carbon atom of methane (CH4)
Made up of 4 C-H sigma (σ) bonds Each hybrid sp3 orbitals of carbon undergoes end-on overlap with s-orbitals of H atoms Tetrahedral geometry 3/25/2017

26 Carbon atoms of ethyne (acetylene - C2H2)
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 3/25/2017

27 Carbon atoms of ethene (ethylene - C2H4)
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) 3/25/2017

28 Formaldehyde has following Lewis structure:
Describe it bonding in terms of appropriate hybridized/unhybridized orbitals VSEPR predicts trigonal planar geometry which suggests sp2 hybrid orbitals on C 3/25/2017

29 Acetonitrile molecule
Predict bond angles around each C Approximately 109° around left C and 180° on right C Give hybridization of each C sp3, sp Determine total number of δ/π bonds 5 δ bonds and 2 π bonds 3/25/2017

30 hybrid orbitals in the CO2 molecule
orbitals of sp hybridized carbon atom orbital arrangement for sp2 hybridized oxygen atom (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 3/25/2017

31 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: CH3Cl 4 δ, 0 Π, sp3 PH3 2 δ, 2 Π, sp H2S 3 δ, 0 Π, sp3 CO32- 3/25/2017

32 SO32- 2 δ, 0 Π, sp3 CS2 3 δ, 1 Π, sp2 SiF4 NO3- 4 δ, 0 Π, sp3 PO43-
ClO4- 4 δ, 0 Π, sp3 3/25/2017

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 3/25/2017

34 Basic s, px, py, and pz 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 3/25/2017

35 Atomic orbitals explain bonding/ account for molecular geometries
Mathematical descriptions of where electrons most likely found Obtained by solving Schrödinger equation As angular momentum (ml) 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) 3/25/2017

36 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: electron pairs localized on atom Bonding pairs: electron pairs in space between atoms 3/25/2017

37 Combine Lewis’s notion of electron-pair bonds with atomic orbitals
Lewis theory: covalent bonding occurs when atoms share electrons (concentrates electron density between nuclei) 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 3/25/2017

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) 3/25/2017

39 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 Bonds (δ-head on/π-sideways) made by overlap of atomic (s,p) or hybridized (sp2) orbitals Formation of bonding orbital accompanied by formation of antibonding orbitals (δ*/π*) which remain unoccupied and does not contribute to structure of molecule 3/25/2017

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 3/25/2017

41 Rules for drawing reasonable resonance structures
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 3/25/2017

42 Homework: Read 9.1, pp Q pp , #12, 14, 16, 21, 22, 27, 28 3/25/2017

43 Molecular Orbital (MO) Theory (Model)

44 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 3/25/2017

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 (don’t 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) 3/25/2017

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 can’t mix Reflects orbital geometry of molecule 3/25/2017

47 Principles for formation of MOs
# MOs formed = # atomic orbitals combined Atomic orbitals combine best with other atomic orbitals of similar energy 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 Each MO accommodates 2 electrons w/opposite spins (Pauli exclusion principle) When MOs of same energy populated, Hund’s rule is followed (equal energy orbitals ½ filled before pairing up) Electron in antibonding orbital “cancels” its corresponding electron in bonding orbital 3/25/2017

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 (s) and one antibonding (s*) Each line in diagram represents orbital 3/25/2017

49 Bonding Orbital MOs w/lower energy than its corresponding original atomic orbitals Promotes formation of stable bond High electron density along internuclear axis 3/25/2017

50 Antibonding Orbital MOs w/higher energy than its corresponding atomic orbitals Destabilizes (negative impact) formation of bond Much lower electron density along internuclear axis 3/25/2017

51 Antibonding molecular orbital is designated as π* molecular orbital
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 3/25/2017

52 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 3/25/2017

53 Pg , Sample 9.6 For the species O2, O2+, and O2-, give the electron configuration and the bond order for each. Which has the strongest bond? 3/25/2017

54 Paramagnetism: (oxygen-2 unpaired e’s)
Magnetism can be induced in some nonmagnetic materials when in presence of magnetic field Paramagnetism: (oxygen-2 unpaired e’s) Unpaired electrons Attracted to induced magnetic field Much stronger than diamagnetism Diamagnetism Paired electrons Repelled from induced magnetic field Much weaker than paramagnetism 3/25/2017

55 Calculate bond order of NO+ ion ½(10 - 4) = 3
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  3/25/2017

56 Hydrogen Atom (H) 1 bonding electron/0 antibonding electrons
Stable (lower energy, greater stability) Bond order of ½ One unpaired electron-paramagnetic 3/25/2017

57 Hydrogen Molecule (H2) 2 bonding electrons in δ1s 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 3/25/2017

58 Does He2 exist? He2 has four electrons, two in s1s orbital and two in s*1s orbital Bond order of 0, so He2 does not exist 3/25/2017

59 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 3/25/2017

60 Molecular orbital energy-level diagrams, bond orders, bond energies, and bond lengths for diatomic molecules Note that for O2 and F2, 2p orbital lower in energy than the π2p orbitals 3/25/2017

61 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 3/25/2017

62 N2 NO bond order = 3 bond order = 2.5

63 Outcomes of MO Model Strengths Disadvantages
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 3/25/2017

64 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 3/25/2017

65 Benzene s bonds (C - H and C - C) are sp2 hybridized
Localized model p bonds result of remaining p orbitals above/ below plane of benzene ring Delocalizing gives stability 3/25/2017

66 Homework: Read 9.2-9.5, pp. 426-441 Q pp. 443-444, #32, 33, 38, 40, 46
Do 1 additional exercise and 1 challenge problem Submit quizzes by to me: 3/25/2017

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