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Reaction of O, S and N with H AtomsThe complete electron configurations for S could be written asSulfur 1s22s22p63px23py13pz1

Again, we can pair up all electrons if S and two H atoms combine to form H2S. The valence bond picture suggests that all bond angles in H2O, NH3 and H2S should be 90o. This is close to the value seen in H2S (92o) but significantly underestimates bond angles in H2O (105o) and NH3 (107o).

Bonding in H2S represented by atomic orbital overlap

FIGURE 11-3Copyright 2011 Pearson Canada Inc.General Chemistry: Chapter 11Slide 2 of 57Chemistry 140 Fall 20022For sulfur only 3p orbitals are shown. The phases of the lobes of the sulfur 3p orbitals are shown in red and blue for positive and negative. However, we do not know which lobe is positive or negative; all we know is that they are opposite. Although the angular part of the p orbital is circular, when multiplied by the radial part of the wave function, each lobe is more pear shaped, similar to the probability distribution shown in Figure 8-27(c). Bond formation occurs between orbitals that are in phase (same color), although the hydrogen 1s orbital is colored yellow for clarity.The Methane ProblemThe ground state configurations for C can be written as 1s22s22px12py1 Using the valence bond picture and the concept of paired electrons in molecular orbitals we might expect C to react with H atoms to form CH2. The CH2 molecule does form but is unstable (a transient species). However, carbon happily reacts with H to form the methane, CH4.

Methane and HybridizationBy experiment, as previously discussed, methane has a regular tetrahedral geometry four equal bond distances and all bond angles of 109.5o. The regular geometry of methane and its ability to form four bonds can be explained using the concept of hybridization. How have we explained carbons tendency to form four bonds previously? Methane and Hybridization contd:In the hybridization picture we imagine methane being formed from C and H atoms in three steps. In the first step we take a ground state C atom and excite one electron (from the 2s orbital) to form the lowest lying ( or first) excited state.Carbon Ground State: 1s22s22px12py1Carbon Excited State: 1s22s12px12py12pz1

Methane and Hybridization contd:In the second step we imagine combining the single occupied 2s orbital and the three occupied 3p orbitals in the excited to form four equivalent sp3 hybrid orbitals (each containing a single unpaired electron). In step three the hybridized C atom reacts with four H atoms to form a CH4 molecule. The process is represented on the next few slides. Hybridization of Atomic Orbitals

Copyright 2011 Pearson Canada Inc.General Chemistry: Chapter 11Slide 7 of 57If we try to extend the unmodified valence-bond method of Section 11-2 to a greater number of molecules, we are quickly disappointed. In most cases, our descriptions of molecular geometry based on the simple overlap of unmodified atomic orbitals do not conform to observed measurements. For example, based on the ground-state electron configuration of the valence shell of carbon We have been describing bonded atoms as though they have the same kinds of orbitals (that is, and so on) as isolated, nonbonded atoms. This assumption worked rather well for and but we have no reason to expect these unmodified pure atomic orbitals to work equallywell in all cases.The mathematical process of replacing pure atomic orbitals with reformulated atomic orbitals for bonded atoms is called hybridization, and the new orbitals are called hybrid orbitals.

Chemistry 140 Fall 20027The sp3 hybridization scheme

FIGURE 11-6Copyright 2011 Pearson Canada Inc.General Chemistry: Chapter 11Slide 8 of 57Bonding and structure of CH4

FIGURE 11-7Copyright 2011 Pearson Canada Inc.General Chemistry: Chapter 11Slide 9 of 57The four carbon orbitals are sp3 hybrid orbitals (blue). Those of the hydrogen atoms (yellow) are 1s. The structure is tetrahedral, with H-C-H bond angles of 109.5 (more precisely, 109.471). Remember that the hydrogen orbitals and the carbon hybrid orbitals have the same phase, but we have colored the hydrogen orbitals yellow for clarity.

Chemistry 140 Fall 20029 Organic Compounds and Structures: An OverviewCopyright 2011 Pearson Canada Inc.General Chemistry: Chapter 26Slide 10 of 75FIGURE 26-1Representations of the methane molecule

The simplest organic compounds are those of carbon and hydrogen hydrocarbonsand the simplest hydrocarbon is methane, the chief constituent of natural gas.Tetrahedral structure showing bond angle.Dashed-wedged line structure convention used to suggest a three-dimensional structure through a structural formula. The solid lines represent bonds in the plane of the page. The dashed wedge projects away from the viewer (behind the plane of the page), and the heavy wedge projects toward the viewer (out of the page).Ball-and-stick model. Space-filling model.Chemistry 140 Fall 200210Organic ChemistryThe next two slides illustrate what starts to happen when two or more sp3 hybridized carbons are linked. The chemistry of carbon is infinitely varied and organic compounds are part of all of living things, important energy sources, key pharmaceuticals and so on.The ethane molecule C2H6

FIGURE 26-2Copyright 2011 Pearson Canada Inc.General Chemistry: Chapter 26Slide 12 of 75(a) Structural formula. (b) Dashed-wedged line structure. (c) Space-filling model.Chemistry 140 Fall 200212The propane molecule, C3H8

FIGURE 26-3Copyright 2011 Pearson Canada Inc.General Chemistry: Chapter 26Slide 13 of 75(a) Structural formula. (b) Condensed structural formula. (c) Ball-and-stick model. (d) Space-filling model. (e) Dashed-wedged line notation.Chemistry 140 Fall 200213Alkanes from petroleumCopyright 2011 Pearson Canada Inc.Slide 14 of 75General Chemistry: Chapter 26

Hybridization in NH3 and H2OThe sp3 hybridization picture can also be used to discuss the bonding in NH3 and H2O. The neutral N and O atoms have more valence electrons than does C. We thus end up putting either one lone pair of electrons (for N) or two lone pairs of electrons (for O) into sp3 hybrid orbitals. The following slides represent the process for N (NH3).Bonding in H2O and NH3

Copyright 2011 Pearson Canada Inc.General Chemistry: Chapter 11Slide 16 of 57Notice that hybrid orbitals can accommodate lone-pair electrons as well as bonding electrons.sp3 hybrid orbitals and bonding in NH3

FIGURE 11-8Copyright 2011 Pearson Canada Inc.General Chemistry: Chapter 11Slide 17 of 57An sp3 hybridization scheme conforms to a molecular geometry in close agreement with experimental observations. Excluding the orbital occupied by a lone pair of electrons, the centers of the atoms form a trigonal pyramid. The hydrogen orbitals are colored yellow for clarity, but they have the same phase as the nitrogen hybrid orbitals.

Chemistry 140 Fall 200217Class Examples1. Jumping the gun just a bit lets draw structural formulas for ethane (H3C-CH3), methyl amine (H3C-NH2), methanol (CH3-OH) and propane (H3C-CH2-CH3). What structural features do these molecules have in common? Does the octet rule still hold?

Finding sp3 Hybridized Atoms2. Using the molecular structure of the morphine molecule shown on the next slide, find (a) an sp3 hybridized C atom, (b) an sp3 hybridized N atom and (c) an sp3 hybridized O atom. There may be more than one example of each. The example is a little unfair since many non-terminal H atoms are not shown in this structure.Copyright 2011 Pearson Canada Inc.General Chemistry: Chapter 26Slide 20 of 75

Morphine , a very powerful and addictive painkiller, can be isolated from the opium poppy (Papaver somniferum).Hybridization in B and Be CompoundsA hybridization scheme can be invoked for B that involves exciting a B atom from its ground electronic state 1s22s22px1 (say) to its first excited state 1s22s12px12py1. The hybrid orbitals formed here (from the combination of a single s orbital and two p orbitals are called sp2 hybrid orbitals. The sp2 hybridization scheme is invoked for the BF3 molecule.sp2 Hybrid OrbitalsCopyright 2011 Pearson Canada Inc.Slide 22 of 57General Chemistry: Chapter 11

Carbons group 13 neighbor, boron, has four orbitals but only three electrons in its valence shell. For most boron compounds, the appropriate hybridization scheme combines the 2s and two 2p orbitals into three sp2 hybrid orbitals andleaves one p orbital unhybridized.

In hybridization, molecular orbital, and valence-bond theory, not only is energy conserved, but also the number of orbitals is conserved. sp2 hybrid orbitals still have one p orbital left over and this is particularly important for carbon. Carbon readily uses the leftover p orbitals to form bonds (see page 461). In contrast, silicon, the element one below carbon, does not use the p orbitals as readily because silicon is larger and the p orbitals do not project out far enough to form double and triple bonds. Chemistry 140 Fall 200222The sp2 hybridization scheme

FIGURE 11-9Copyright 2011 Pearson Canada Inc.General Chemistry: Chapter 11Slide 23 of 57sp Hybridizationsp hybridization is important for molecules such as H-CN and H-CC-H. Acetylene (ethyne) is the first member of an important series of organic compounds, the alkynes.sp HybridizationCopyright 2011 Pearson Canada Inc.Slide 25 of 57General Chemistry: Chapter 11

Borons group 2 neighbor, beryllium, has four orbitals and only two electrons in its valence shell. In the hybridization scheme that best describes certain gaseous beryllium compounds, the 2s and one 2p orbital of Be are hybridized into two sp hybrid orbitals, and the remaining two orbitals are left unhybridized.Chemistry 140 Fall 200225The sp hybridization scheme

FIGURE 11-10Copyright 2011 Pearson Canada Inc.General Chemistry: Chapter 11Slide 26 of 57Class Example3. The sp3 hybridization scheme can be invoked to explain the bonding in the silane (SiH4) molecule. What atomic orbitals on silicon would be used to construct hybrid orbitals?Covalent Bonds Orbital Overlap Sigma and Pi BondsThe formation of both sigma bonds ( bonds) and pi bonds ( bonds) is likely familiar.Sigma bonds are formed (for a pair of atoms) by the overlap of atomic orbitals pointing towards, in each case, the other bonded atom. We can use s and p orbitals (etc) to form sigma bonds.

Covalent Bonds Orbital Overlap Sigma and Pi Bonds contd:In sigma bonds the two atomic orbitals used to construct the molecular orbital overlap in the most spatially direct manner possible.The overlap of a H 1s orbital and S 3p orbitals is shown on the next slide. As an aside, one can see that it difficult to point an s orbital in any direction. Why?Bonding in H2S represented by atomic orbital overlap

FIGURE 11-3Copyright 2011 Pearson Canada Inc.General Chemistry: Chapter 11Slide 30 of 57Chemistry 140 Fall 200230For sulfur only 3p orbitals are shown. The phases of the lobes of the sulfur 3p orbitals are shown in red and blue for positive and negative. However, we do not know which lobe is positive or negative; all we know is that they are opposite. Although the angular part of the p orbital is circular, when multiplied by the radial part of the wave function, each lobe is more pear shaped, similar to the probability distribution shown in Figure 8-27(c). Bond formation occurs between orbitals that are in phase (same color), although the hydrogen 1s orbital is colored yellow for clarity.Pi BondsPi () bonds are formed when the bonding electron pair is placed in a molecular orbital formed (frequently) by p orbitals on adjacent atoms overlapping. The p orbitals on the bonded atoms are oriented perpendicular to the internuclear axis (which makes orbital overlap slightly less favourable). Carbon-Carbon Double BondsThe carbon-carbon double bonds in common organic molecules are comprised of one bond and one bond. The simplest molecule of this type is ethylene, H2C=CH2. Ethylene is a planar symmetric molecule (four identical C-H bonds and all bond angles near 120o). Aside: Hydrocarbons containing C=C double bonds are called unsaturated. They can react with H2 to from saturated hydrocarbons. Carbon-Carbon Double Bonds contd:In saturated hydrocarbons (such as propane H3C-CH2-CH3 the bonding behaviour of all C atoms is well explained using the sp3 hybridization scheme. In many unsaturated hydrocarbons the bonding of C atoms joined by double bonds is rationalized using an sp2 hybridization scheme.Carbon-Carbon Double Bonds contd:For the sp3 hybridization scheme (for carbon!) we imagined distributing four valence electrons among a 2s and the three 2p atomic orbitals and scrambling these orbitals together to form four hybrid orbitals. For the sp2 hybridization scheme we will scramble the single 2s orbital and two 2p orbitals to form three hybrid sp2 orbitals. A single 2p orbital remaining is used to from a pi bond. Multiple Covalent BondsEthylene has a double bond in its Lewis structure.VSEPR says trigonal planar at carbon.Copyright 2011 Pearson Canada Inc.General Chemistry: Chapter 11Slide 35 of 57

Bonding in C2H4HHHHCCBonding in Ethylene an Information Packed Slide!The next slide contains lots of information.Upper left corner a representation (for C) showing the structure and disposition in space of the three hybrid sp2 orbitals and the left over carbon p atomic orbital.Upper right corner picture showing how the five sigma bonds in ethylene are formed using four H 1s orbitals and six carbon sp2 orbitals.Bonding in Ethylene an Information Packed Slide contd!Lower left corner. The sigma bonds are in place and the p orbitals on the two C atoms have not yet overlapped to from a pi bond. In fact the p orbitals have been drawn slightly smaller than life for clarity. Bottom center and right. Two representations of the pi bond in ethylene. Note the electron density well away from the C-C internuclear axis.Sigma (s) and pi () bonding in C2H4

FIGURE 11-14Copyright 2011 Pearson Canada Inc.General Chemistry: Chapter 11Slide 38 of 57Class examples4. The ethanoic acid (vinegar) moleculae and the methyl ethanoate molecule (an ester) shown on the next slide contain C=O double bonds. What is the hybridization of the C and O atoms in these double bonds? (Mention acetone, acetaldehyde, formaldehye?)

EstersCH3CO2CH2(CH2)6CH3The distinctive aroma and flavor of oranges are due in part to the ester octyl acetate,Copyright 2011 Pearson Canada Inc.General Chemistry: Chapter 26Slide 40 of 75