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Brad Collins, DTCC Chemical Structures (Part 2) Chapter 11 Some images Copyright © The McGraw-Hill Companies, Inc. Organic Structures • Organic compounds do not follow the AXn general formula, but: • DO generally follow the octet rule. • Hydrogen and halogens are always outer atoms • In neutral atoms, N and P will retain their lone pair; O and S will retain their 2 lone pairs • Rarely see expanded octets 11.1 Organic Compounds • Use structural formulas to depict organic compounds • Isomer - compounds with the same molecular formula, but different structural formulas. • Ethanol: • Dimethyl ether: • Molecular: C2H6O • Molecular: C2H6O • Structural: • Structural: • Condensed structural: CH3CH2OH • Condensed structural: CH3OCH3 11.1 Organic Compounds • Condensed structural formulas • Hydrogens follow the atom they are bonded to. • • Single bonds are omitted • • CH3CH3OH, not CH3–CH3–O–H Multiple bonds ARE shown • • CH3, not H3C CH2=CH2 or CH CH Side chains are often shown: • but may be written in parenthesis: CH3CH(OH)CH3 11.1 Organic Compounds • Condensed structural formulas • Cyclic structures a drawn with the geometric figure they resemble: • Cyclohexane, C6H12 11.1 Practice Draw the Lewis structure for diethyl ether, CH3CH2OCH2CH3 11.1 Bond Enthalpy • The enthalpy required to break a particular bond in 1 mole of gaseous molecules Bond Enthalpy ΔH0 = 436.4 kJ H2 (g) H (g) + H (g) Cl2 (g) Cl (g) + Cl (g) ΔH0 = 242.7 kJ HCl (g) H (g) + Cl (g) ΔH0 = 431.9 kJ O2 (g) O (g) + O (g) ΔH0 = 498.7 kJ O O N2 (g) N (g) + N (g) ΔH0 = 941.4 kJ N N Bond Energies Single bond < Double bond < Triple bond 11.2 Bond Enthalpy • In polyatomic molecules (e.g., H2O) the average bond enthalpy is listed H2O (g) H (g) + OH (g) ΔH0 = 502 kJ OH (g) H (g) + O (g) ΔH0 = 427 kJ 502 + 427 = 464 kJ Average OH bond energy = 2 11.2 Bond Enthalpy Note: Forming the listed bond, releases the tabled amount of energy • • Breaking bonds is endothermic Forming bonds is exothermic 11.2 Bond Enthalpy (BE) and Enthalpy changes in reactions Imagine a reaction proceeding by breaking all bonds in the reactants and then using the gaseous atoms to form all the bonds in the products. ΔH0 = total energy input – total energy released = ΣBE(reactants) – ΣBE(products) 11.2 Bond Enthalpy and ∆Hrxn H2(g) + F2 (g) —> 2HF(g) ! ∆Hrxn = –543.2 kJ 11.2 Practice Use bond energies to calculate the enthalpy change for: H2 (g) + F2 (g) 2HF (g) Type of bonds broken Number of bonds broken Bond energy (kJ/mol) Energy change (kJ) H H 1 436.4 436.4 F F 1 156.9 156.9 Type of bonds formed H F Number of bonds formed Bond energy (kJ/mol) 2 568.2 Energy change (kJ) 1136.4 ΔH0 = 436.4 + 156.9 – 2 x 568.2 = -543.1 kJ 11.2 Shapes of Molecules • VSEPR Theory • Valence Shell Electron Pair Repulsion • Molecules will form shapes that minimize the repulsive force between the electron pairs (bonding pairs and lone pairs) • Arrangement describes how the electron pairs arrange themselves around an atom • Shape describes how the molecule actually looks in 3-dimensional space 11.3 Atomic Arrangements Two Items Around a Central Atom • Classify molecules according to how many atoms (B) are bonded to the central atom and how many lone pairs (E) are on the central atom. • Two outer atoms, no lone pairs (AB2) • Electron pairs align opposite the central atom, A Class # of atoms bonded to central atom # lone pairs on central atom AB2 2 0 Arrangement of electron pairs Molecular Shape linear B linear B B B 11.3 Beryllium Chloride Cl Be Cl lone pairs on to central atom 2 0atoms bonded central atom 10.1 Atomic Arrangements Three Items Around a Central Atom • Three outer atoms, no lone pairs (AB3) • Class AB3 Electron pairs align at 0, 120, and 360º around the central atom, A # of atoms bonded to central atom 3 # lone pairs on central atom 0 Arrangement of electron pairs Molecular Shape trigonal planar trigonal planar 120º 11.3 Boron Trifluoride Trigonal Planar 10.1 Atomic Arrangements Three Items Around a Central Atom • 2 outer atoms, one lone pair (AB2E) • Class AB2E Electron pairs align at 0, 120, and 360º around the central atom, A # of atoms bonded to central atom 2 # lone pairs on central atom Arrangement of electron pairs trigonal planar 1 B Molecular Shape bent B 11.3 Atomic Arrangements Four Items Around a Central Atom • Four outer atoms, no lone pairs (AB4) • Electron pairs align at ~109.5º angles around the central atom, A Class # of atoms bonded to central atom # lone pairs on central atom AB4 4 0 Arrangement of electron pairs Molecular Shape tetrahedral B tetrahedral B B B 11.3 Methane 11.3 Atomic Arrangements Four Items Around a Central Atom • Three outer atoms, one lone pair (AB3E) • Electron pairs align at ~109.5º angles around the central atom, A Class # of atoms bonded to central atom # lone pairs on central atom AB3E 3 1 Arrangement of electron pairs tetrahedral Molecular Shape trigonal pyramidal 11.3 Atomic Arrangements Four Items Around a Central Atom • Two outer atoms, two lone pairs (AB2E2) • Electron pairs align at ~109.5º angles around the central atom, A Class # of atoms bonded to central atom # lone pairs on central atom AB2E2 2 2 Arrangement of electron pairs tetrahedral Molecular Shape bent 11.3 Effect of Lone Pairs bonding-pair vs. < bonding pair repulsion lone-pair vs. bonding pair repulsion < lone-pair vs. lone pair repulsion 11.3 Atomic Arrangements Five Items Around a Central Atom • Five outer atoms, no lone pairs (AB5) • Electron pairs align at ~109.5º angles around the central atom, A Class # of atoms bonded to central atom # lone pairs on central atom AB5 5 0 Arrangement of electron pairs Molecular Shape Trigonal bipyramidal Trigonal bipyramidal B B B B B 11.3 Phosphorus Pentachloride 11.3 Atomic Arrangements Six Items Around a Central Atom • Six outer atoms, no lone pairs (AB6) • Electron pairs align at ~109.5º angles around the central atom, A Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Shape AB6 6 0 Octahedral Octahedral B B B B B B 11.3 11.3 Sulfur Hexafluoride 11.3 Atomic Arrangement Summary 11.3 Polarity • Definition: Unequal sharing of electrons in covalent bonds • Electrons spend more time with one atom in the bond • Bond has negative and positive end electron poor region electron rich region H F δ+ δ− 11.4 Polar Molecules • To be polar, molecules must meet 2 criteria: • Must have a polar bond (dipole) • Must have a molecular shape that allows a dipole moment to exist. • • Dipoles are vectors (have magnitude and direction) • Dipole moment is the vector sum of all the dipoles present. Examples: HF, NH3, and NF3 HF electron poor region electron rich region H F δ+ δ− 11.4 NH3 NF3 11.4 Which of the following molecules have a dipole moment? H2O, CO2, SO2, and CH4 O H H dipole moment polar molecule S O O dipole moment polar molecule H O C O no dipole moment nonpolar molecule H C H H no dipole moment nonpolar molecule 11.4 Does CH2Cl2 have a dipole moment? 11.4 Dipole Moments of Polar Molecules 11.4 Trends in Polarity • Molecules with a linear, trigonal planar, or tetrahedral arrangement are predicted to be nonpolar if: • There are no lone pairs on the central atom • All the outer atoms are the same • If one or more of these aren’t true, the molecule is usually polar. • Does NOT work for trigonal bipyramidal or octahedral arrangements. 11.4 Practice: Polar Molecules Predict the polarity of the following molecules: Chloroform, CHCl3 Phosphorus trichloride, PCl3 Bromine pentafluoride, BrF5 11.4