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UT CH 301 - Worksheet 7 Answer Key

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CH301 Fall 2009 Worksheet 7 Answer Key1. How can you decide whether a bond is polar or not? If a molecule has polar bonds, doesthat make the molecule polar?A polar bond is between two atoms with different electronegativities. One atom with hog theelectrons, giving it a slightly negative charge, denoted δ-. The other atom will have aslightly positive charge, δ+. A molecule with polar bonds may not be polar overall.Symmetry of the molecule can cancel out the dipoles created by the polar bonds. Forexample, BF3 has three polar bonds but all these polar bonds are in the same plane andcancel each other out. If you replace one of the F with another atom then there would bepolarity in the molecule because the electronegativities would not cancel out.2. Fill in the chart and then rank the molecules in order of increasing polarity.NH3 O2 NF3 C2H5OHDraw structureCalculate ΔEN ofall bondsΔEN = 3-2.2 =.8ΔEN=2.5-2.5= 0 ΔEN = 4 - 3 = 1ΔEN C-H =2.5-2.2 = .3ΔEN C-O =3.5-2.5 = 1ΔEN O-H =3.5-2.2 =1.3Polar or nonpolarmolecule?Polar Nonpolar PolarPolar andnonpolar endsO2 < C2H5OH < NH3 < NF3Homonuclear diatomic molecules are never polar.3. Rank the following in order of increasing polarity:D = B = A < CMolecule C is the only one where the dipoles do not cancel out.4. Explain how carbon, with only 2 unpaired electrons, can form 4 bonds to fulfill its octet.Account for any energy changes (if you're putting in energy, where is it then released).What kind of hybrid orbitals will carbon then form?One of the electrons in the 2s orbital is promoted to the 2p orbital, giving carbon theelectronic configuration of [He]2s12p3. Now there are 4 unpaired electrons that can bond.Promoting an electron takes energy, but the amount of energy released when bonds formcompensates this. Also, only a small amount of energy is needed to initially promote theelectron to the 2p orbital. If the electrons stay in the 2s and 2p orbitals then we would see 2different types of bonds, but we know that we only see one type. In order to form 4 bondsthese electrons will then hybridize their orbitals to form four sp3 hybrid orbitals.5. Group 14 elements can all form four bonds. However, as you go down the periodic tableit gets harder for them to form multiple bonds with one another like carbon can. Explainwhy carbon is so special.As you go down the periodic table the atomi radii increases (see these trends are stillimportant!), which means that they are too big to have significant pi overlap.6. Explain how each of the hybrid orbitals are constructed. Account for the electronarrangement in the shape of molecules.spsp2sp3sp3dsp3d2First recognize which atomic orbitals are blending to form hybrid orbitals. Remember thatelectrons like to be as far away as possible from each other, so the hybrid orbitals are goingto distribute with the maximum angle of separation as possible.sp: a linear arrangement of electron pairs requires two hybrid orbitals, and so an s-orbitaland a p-orbital are mixed to form two sp hybrid orbitals.sp2: seen in trigonal planar molecules, one s-orbital and two p-orbitals mix to form threesp2 hybrid orbitals. These hybrids all lie in the same plane and point toward the corners ofan equilateral triangle.sp3: tetrahedral, an s-orbital and three p-orbitals mix together to give four sp3 hybridorbitals that point to a corner of a tetrahedron.sp3d: trigonal bipyramidal, a d orbital is included in this hybrid orbital because the centralatom has to accommodate for 5 electron pairs. There are only four sp3 hybrid orbitals, sowe must include the d now. An s-orbital, three p-orbitals and a d-orbital blend together.sp3d2: octahedral, central atoms that have 6 electron pairs to account for must use anotherd orbital. So, an s-orbital, three p-orbitals and two d-orbitals blend together to make sixhybrid sp3d2 hybrid orbitals.7. Is it possible to have more hybrid orbitals than atomic orbitals of an atom?No, orbitals are not added or reduced when hybridizing atomic orbitals.8. For formic acid, HCOOH:Draw the Lewis structureUse VSEPR to determine the geometryIdentify bond anglesIdentify hybrid orbitalsHow many sigma an pi bonds? (Hint: There is an -OH group.)Answer:The carbon atom has three sigma bonds and no lone pairs, the oxygen atoms andhydrogen surround it in a trigonal planar planar arrangement, with bond angles of 120. Theoxygen of the -OH group is in the plane with carbon, but the hydrogen of the -OH is not.Oxygen has two lone pairs and two sigma bonds, so it has four regions of electron density.Thus, the oxygen of the -OH group has a tetrahedral electron arrangement, with bondangles of 109.5.Since there is a double bond between carbon and one of the oxygens a p-orbital from theoxygen will form a pi bond with the non-hybridized p orbital of carbon. The three sigmabonds around carbon are sp2 hybridized. The O of the -OH group has four regions ofelectron density, so it is sp3 hybridized.9. The bond angle of an sp3 hybridized atom is 109.5 and that of an sp2 hybridized atom is120. Should the bond angle increase or decrease between two hybrid orbitals when the s-character of a hybrid orbital increases?Increase. As more orbitals are hybridized the hybrid orbital looks less like an s orbital. So ansp hybrid orbital looks most like an s-orbital and an sp3d2 hybrid orbital looks least like an sorbital. Since sp looks the most like an s-orbital and there are only two sp hybrid orbitalsthey have the highest bond angle.10.Why does valence bond theory not account for bonding in polyatomic molecules likeCH4?VB theory says that bonds are formed when unpaired electrons in valence-shell atomicorbitals pair by atomic orbital overlap end-to-end to form sigma bonds or side-to-side toform pi bonds. However, carbon only has 2 unpaired valance electrons but as we see withmethane, CH4, it forms four bonds. This is why we see hybridization of orbitals, so that wecan have four unpaired electrons. In the case of methane these hybrid orbitals are sp3.11. Explain the difference between sigma and pi bonds. What types of orbitals are involvedin each?Sigma bonds are areas of electron density along the axis between the contributing atoms.Sigma bonds can form between s orbitals, p orbitals, and all hybrid orbitals. Pi bonds areares of electron density off the axis between the contributing atoms. Pi bonds can be formedby p orbitals.12. Explain why hybrid orbitals make sigma bonds.S orbitals can only form sigma bonds because they form areas of electron


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UT CH 301 - Worksheet 7 Answer Key

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