4_Hybridization.pdf

Biol 112L ◍ Hydrocarbons: those compounds that only contain carbon (C) and hydrogen(H). 2  Hybridization ◍ In many molecules, the electron orbitals surrounding one atom "overlap" with one or several of the orbitals surrounding another...

Biol 112L ◍ Hydrocarbons: those compounds that only contain carbon (C) and hydrogen(H). 2  Hybridization ◍ In many molecules, the electron orbitals surrounding one atom "overlap" with one or several of the orbitals surrounding another atom, forming a bond. In the process of overlapping, the orbitals are distorted. 3  Hybridization ◍ sp3 Hybrid orbitals – bond angles of approximately 109.5° (tetrahedral geometry), formed by the combination of one s atomic orbital and three p atomic orbitals ◍ Single-bonded hydrocarbons are formed when an sp3 hybrid orbital is formed between two carbon atoms. 4 ◍ Note that the superscript 3 in the name sp3 tells how many of each type of atomic orbital combine to form the hybrid, not how many electrons occupy it. 5 sp3 hybrid orbitals 6  sp3 hybrid orbitals ◍ The asymmetry of sp3 orbitals arises because the two lobes of a p orbital have different algebraic signs, + and - , in the wave function. ○ Thus, when a p orbital hybridizes with an s orbital, the positive p lobe adds to the s orbital but the negative p lobe subtracts from the s orbital. ○ The resultant hybrid orbital is therefore unsymmetrical about the nucleus and is strongly oriented in one direction. 7 Methane 8  ethane The carbon–carbon bond is formed by s overlap of sp3 hybrid orbitals. For clarity, the smaller lobes of the sp3 hybrid orbitals are not shown. 9  Hybridization ◍ Double-bonded hydrocarbons are formed when an sp2 orbital is formed. ◍ sp2 Hybrid orbitals – bond angles of approximately 120° (trigonal geometry), formed by the combination of one s atomic orbital and two p atomic orbitals 10 sp2 hybrid orbitals 11  sp2 hybrid ◍ When two carbons with sp2 hybridization approach each other, they form a strong s bond by sp2–sp2 head-on overlap. At the same time, the unhybridized p orbitals interact by sideways overlap to form what is called a pi bond. ◍ The combination of an sp2–sp2 s bond and a 2p–2p p bond results in the sharing of four electrons and the formation of a carbon–carbon double bond. 12 sp2 orbitals 13  Hybridization ◍ Triple-bonded hydrocarbons are formed when an sp orbital is formed. 14  Hybridization ◍ sp Hybrid orbitals – bond angles of approximately 180° (linear geometry), formed by the combination of an s atomic orbital and a p atomic orbitals 15 sp hybrid 16 The two sp hybrid orbitals are oriented 180° away from each other, perpendicular to the two remaining p orbitals (red/blue). 17  sp hybrid ◍ When two sp carbon atoms approach each other: ○ sp hybrid orbitals on each carbon overlap head-on to form a strong sp–sp s bond ○ the pz orbitals from each carbon form a pz–pz p bond by sideways overlap ○ the py orbitals overlap similarly to form a py–py p bond ○ The net effect is the sharing of six electrons and formation of a carbon–carbon triple bond. 18 acetylene 19 20 21  Nitrogen ◍ H-N-H bond angle in methylamine is 107.1°, and the C-N-H bond angle is 110.3°, (close to the 109.5° tetrahedral angle found in methane) ◍ four sp3-hybridized orbitals like carbon ○ unshared lone pair of electrons 22  Oxygen ◍ sp3-hybridized ◍ The C-O-H bond angle in methanol is 108.5°, very close to the 109.5° tetrahedral angle. ◍ Two of the four sp3 hybrid orbitals on oxygen are occupied by nonbonding electron lone pairs, and two are used to form bonds. 23  Phosphorus and sulfur ◍ Phosphorus often forms five covalent bonds, and sulfur often forms four. 24  Phosphorus and sulfur ◍ The O-P-O bond angle in such compounds is typically in the range 110 to 112°, implying sp3 hybridization for the phosphorus orbitals. 25  Phosphorus and sulfur ◍ sp3 hybridization 26  Effect of substituents polar covalent bonds the bonding electrons are attracted more strongly by one atom than the other so that the electron distribution between atoms is not symmetrical 27 ○ Bond polarity is due to differences in electronegativity the intrinsic ability of an atom to attract the shared electrons in a covalent bond electronegativity of elements 28  Electronegativity generally increases from left to right across the periodic table and decreases from top to bottom. 29 ◍ bonds between atoms whose electronegativities differ by less than 0.5 are nonpolar covalent ◍ bonds between atoms whose electronegativities differ by 0.5–2 are polar covalent ◍ bonds between atoms whose electronegativities differ by more than 2 are largely ionic 30 ◍ indicate the direction of bond polarity ◍ By convention, electrons are displaced in the direction of the arrow. ◍ The tail of the arrow (which looks like a plus sign) is electron-poor, and the head of the arrow is electron-rich. Question ◍ Which element in each of the following pairs is more electronegative? ◍ (a) Li or H (b) B or Br (c) Cl or I (d) C or H Question ◍ Look at the following electrostatic potential map of chloromethane, and tell the direction of polarization of the C-Cl bond:  Molecular polarity ◍ results from the vector summation of all individual bond polarities and lone-pair contributions in the molecule ○ strongly polar substances are often soluble in polar solvents like water, whereas less polar substances are insoluble in water ○ Net molecular polarity is measured by a quantity called the dipole moment  Dipole moment ◍ the magnitude of the charge Q (C, coulomb) at either end of the molecular dipole times the distance r (meter) between the charges ◍ unit: debyes Dipole moment (2nd formula)  strong dipole moments, lone pairs  zero dipole moments; symmetrical structures  Question ◍ Make three-dimensional drawings of the following molecules, and predict whether each has a dipole moment. If you expect a dipole moment, show its direction. *the atom with the higher electronegativity is the negative end of the dipole ◍ (a) H2C=CH2 (b) CHCl3 (c) CH2Cl2 (d) H2C=CCl2 end 41