Representative Carbon Compounds 3 PDF

Chapter 2 Families of Carbon Compounds Functional Groups, Intermolecular Forces Chapter 2 In this chapter we will consider:  The major functional groups  The correlation between properties of functional groups and molecules and intermolecular forces...

Chapter 2 Families of Carbon Compounds Functional Groups, Intermolecular Forces Chapter 2 In this chapter we will consider:  The major functional groups  The correlation between properties of functional groups and molecules and intermolecular forces Chapter 2 1. Hydrocarbons  Hydrocarbons are compounds that contain only carbon and hydrogen atoms Alkanes  hydrocarbons that do not have multiple bonds between carbon atoms e.g. pentane cyclohexane Chapter 2  Alkenes  contain at least one carbon–carbon double bond e.g. propene cyclohexene Chapter 2  Alkynes  contain at least one carbon–carbon triple bond Chapter 2  Aromatic compounds  contain a special type of ring, the most common example of which is a benzene ring CH3 COOH e.g. benzene toluene benzoic acid Chapter 2 1A. Alkanes  The primary sources of alkanes are natural gas and petroleum  The smaller alkanes (methane through butane) are gases under ambient conditions  Methane is the principal component of natural gas  Higher molecular weight alkanes are obtained largely by refining petroleum H H H H Chapter 2 1B. Alkenes  Ethene and propene, the two simplest alkenes, are among the most important industrial chemicals produced in the United States  Ethene is used as a starting material for the synthesis of many industrial compounds, including ethanol, ethylene oxide, ethanal, and the polymer polyethylene H H C C H H Chapter 2  Propene is the important starting material for acetone, cumene, and polypropylene  Examples of naturally occurring alkenes -Pinene An aphid alarm (a component of pheromone turpentine) Chapter 2 1C. Alkynes  The simplest alkyne is ethyne (also called acetylene) H C C H  Examples of naturally occurring alkynes O Br Cl C C C C C CH3 Capillin O (an antifungal agent) Br Dactylyne (an inhibitor of pentobarbital metabolism) Chapter 2 1D. Benzene  All C C bond lengths are the same (1.39 Å) (compare with C–C single bond 1.54 Å, C=C double bond 1.34 Å)  Extra stabilization due to resonance  aromatic Chapter 2  3-Dimensional structure of benzene p-electrons above and below ring Planar structure All carbons sp2 hybridized Chapter 2  The lobes of each p orbital above and below the ring overlap with the lobes of p orbitals on the atoms to either side of it  the six electrons associated with these p orbitals (one electron from each orbital) are delocalized about all six carbon atoms of the ring Chapter 2 2. Polar Covalent Bonds Li F  Lithium fluoride has an ionic bond H H H C C H H H  Ethane has a covalent bond. The electrons are shared equally between the carbon atoms Chapter 2  electronegativity + - d d C C C O 2.5 3.5 equal sharing unequal sharing ⊖ ⊖ of e of e (non-polar bond) (polar bond) Chapter 2  Electronegativity (EN) The intrinsic ability of an atom to attract the shared electrons in a covalent bond Electronegativities are based on an arbitrary scale, with F the most electronegative (EN = 4.0) and Cs the least (EN = 0.7) Chapter 2 element (EN) H (2.1) Li Be B C N O F (1.0) (1.6) (2.0) (2.5) (3.0) (3.5) (4.0) Increasing EN Na Mg Si P S Cl (0.9) (1.2) (1.8) (2.1) (2.5) (3.0) K Br (0.8) (2.8) Rb I (0.8) (2.5) Cs (0.7) Increasing EN Chapter 2  + - + - d d d d C N C Cl 2.5 3.0 2.5 3.0 + - + - d d d d H C Si C 2.1 2.5 1.8 2.5 Chapter 2 3. Polar and Nonpolar Molecules Dipole distance between the =  moment the charges charge m=rQ  Dipole moments are expressed in debyes (D), where 1 D = 3.336  10–30 coulomb meter (C m) in SI units Chapter 2  - d Cl > C net dipole H (1.87 D) H H d+ Chapter 2  Molecules containing polar bonds are not necessarily polar as a whole, for example (1) BF3 (m = 0 D) (2) CCl4 (m = 0 D) Cl F o C 120 Cl B Cl F F Cl (trigonal planar) (tetrahedral) Chapter 2  Dipole moment of some compounds Dipole Dipole Compound Compound Moment Moment NaCl 9.0 H2O 1.85 CH3NO2 3.45 CH3OH 1.70 CH3Cl 1.87 CH3COOH 1.52 CH3Br 1.79 NH3 1.47 CH3I 1.64 CH4 0 CHCl3 1.02 CCl4 0 Chapter 2 3A. Dipole Moments in Alkenes cis- trans- 1,2-Dichloroethene 1,2-Dichloroethene H H H Cl C C C C Cl Cl Cl H resultant dipole moment (m = 1.9 D) (m = 0 D) Chapter 2  Physical properties of some cis-trans isomers m.p. b.p. Compound (m) (oC) (oC) cis-1,2-Dichloroethene -80 60 1.90 trans-1,2-Dichloroethene -50 48 0 cis-1,2-Dibromoethene -53 112.5 1.35 trans-1,2-Dibromoethene -6 108 0 Chapter 2 4. Functional Groups  Functional groups are common and specific arrangements of atoms that impart predictable reactivity and properties to a molecule Chapter 2 4A. Alkyl Groups and the Symbol R Chapter 2  These and others can be designated by R  General formula for an alkane is R–H Chapter 2 4B. Phenyl and Benzyl Groups  Phenyl group  Benzyl group CH2 or or C6H5CH2 or Bn Chapter 2 5. Alkyl Halides or Haloalkanes  R–X (X = F, Cl, Br, I) Examples Attached to Attached to Attached to 1 carbon atom 2 carbon atoms 3 carbon atoms C C C C Cl C Br C I a 1o chloride a 2o bromide a 3o iodide Chapter 2  an alkenyl bromide (Br bonded to an alkene carbon) an aryl chloride (Cl bonded to an aromatic ring) Chapter 2 6. Alcohols and Phenols  R–OH Examples Chapter 2  Alcohols may be viewed structurally in two ways: As hydroxyl derivatives of alkanes As alkyl derivatives of water ethyl group CH3CH2 H CH3CH3 109.5o O 104.5o O H hydroxyl H group Ethane Ethyl alcohol Water (ethanol) Chapter 2 7. Ethers  R–O–R Examples ~100o O O Acyclic Cyclic Chapter 2 8. Amines  R–NH2 H H CH3 N N N H3C H H3C CH3 H3C CH3 (1o) (2o) (3o) N N H (cyclic) (aromatic) Chapter 2 9. Aldehydes and Ketones O O R H R R (aldehydes) (ketones) O O O O H , , H ketone aldehyde Chapter 2  Aldehydes and ketones have a trigonal planar arrangement of groups around the carbonyl carbon atom o O 121 121o H H 108o Chapter 2 10. Carboxylic Acids, Esters, and Amides O O O R OH R OR R Cl (carboxylic (ester) (acid acid) chloride) O O O R NR2 R O R (amide) (acid anhydride) Chapter 2 11. Nitriles  R–C≡N Chapter 2 12. Summary of Important Families of Organic Compounds Chapter 2 Chapter 2 13. Physical Properties & Molecular Structure 13A. Ionic Compounds: Ion-Ion Forces  The melting point of a substance is the temperature at which an equilibrium exists between the well- ordered crystalline state and the more random liquid state Chapter 2  If the substance is an ionic compound, the ion–ion forces that hold the ions together in the crystalline state are the strong electrostatic lattice forces that act between the positive and negative ions in the orderly crystalline structure  A large amount of thermal energy is required to break up the orderly structure of the crystal into the disorderly open structure of a liquid Chapter 2  The boiling points of ionic compounds are higher still, so high that most ionic organic compounds decompose before they boil Chapter 2  Physical properties of selected compounds Compound Structure mp (oC) bp (oC) (1 atm) Ethane CH3CH3 -172 -88.2 Chloroethane CH3CH2Cl -138.7 13.1 Ethyl alcohol CH3CH2OH -114 78.5 Acetaldehyde CH3CHO -121 20 Acetic acid CH3CO2H 16.6 118 Sodium acetate CH3CO2Na 324 dec Chapter 2 13B. Intermolecular Forces (van der Waals Forces)  The forces that act between molecules are not as strong as those between ions  These intermolecular forces, van der Waals forces, are all electrical in nature Dipole-dipole forces Hydrogen bonds Dispersion forces Chapter 2  Dipole-dipole forces  Dipole-dipole attractions between polar molecules + - + d O d d O - d dipole-dipole attraction H H H d H C + - + d C Cl d Cl - d H H Chapter 2  Hydrogen bonds  Dipole-dipole attractions between hydrogen atoms bonded to small, strongly electronegative atoms (O, N, or F) and nonbonding electron pairs on other such electronegative atoms  Hydrogen bonds (bond dissociation -1 energies of about 4 – 38 kJ mol ) are weaker than ordinary covalent bonds but much stronger than the dipole–dipole interactions Chapter 2  Hydrogen bonds d+H d+H d-O d-O d+ H d+ H hydrogen bond + + d d H H d+ d+ H N- H N- d d +H +H d d Chapter 2  Hydrogen bonds  Hydrogen bonding explains why water, ammonia, and hydrogen fluoride all have far higher boiling points than methane (bp -161.6°C), even though all four compounds have similar molecular weights  One of the most important consequences of hydrogen bonding is that it causes water to be a liquid rather than a gas at 25°C Chapter 2  Hydrogen bonds  Calculations indicate that in the absence of hydrogen bonding, water would have a bp near -80°C and would not exist as a liquid at room temperature Chapter 2  Dispersion forces (London forces)  The average distribution of charge in a nonpolar molecule over a period of time is uniform  At any given instant, however, because electrons move, the electrons and therefore the charge may not be uniformly distributed  Electrons may, in one instant, be slightly accumulated on one part of the molecule, and, as a consequence, a small temporary dipole will occur Chapter 2  Dispersion forces (London forces)  This temporary dipole in one molecule can induce opposite (attractive) dipoles in surrounding molecules  These temporary dipoles change constantly, but the net result of their existence is to produce attractive forces between nonpolar molecules Chapter 2  Two important factors determine the magnitude of dispersion forces  The relative polarizability of electrons of the atoms involved  The electrons of large atoms such as iodine are loosely held and are easily polarized, while the electrons of small atoms such as fluorine are more tightly held and are much less polarizable Chapter 2  - stronger + - + d d d d - + dispersion forces - + d d d d + + d- I d d- I d + d- d- d d+ d- C I d+ d- C I d+ d - I d+ d - I d+ d - I d+ d - I d+ d- d+ d- d+ weaker dispersion forces d- F d+ d- F d+ d- d+ d- + C F C F d F F d- F d+ d- F d+ Chapter 2  The relative surface area of the molecules involved  The larger the surface area, the larger is the overall attraction between molecules caused by dispersion forces Chapter 2 e.g. Pentane vs. Neopentane (both C5H12) - d - - d d - - d- d- H H H d - - d d - - d H H H H d - H H d H H d larger surface C C d- H - area H d H H  stronger H C C C H + + dispersion + + d H H d d d d+ H + Hd+ forces H H H d d+ + d+ d+ + d d d+ d+ -- - d d d- d d - d- H H H H d - - d- d- H H H d - - d H H d smaller surface area H H H H - C C -  weaker d H H d H C C C H d+ d+ dispersion H H d H+ Hd+ + forces d d+ + d+ d+ d + + d d Neopentane + H H H Pentane (bp 36oC) (bp 9.5oC) d d+ Chapter 2 d+ 13C. Boiling Points  The boiling point of a liquid is the temperature at which the vapor pressure of the liquid equals the pressure of the atmosphere above it  the boiling points of liquids are pressure dependent, and boiling points are always reported as occurring at a particular pressure Chapter 2  Examples CH3 t Bu CH3 NO2 OH o bp: 245 C / 760 mmHg bp: 260oC / 760 mmHg (74oC / 1 mmHg) (140oC / 20 mmHg) 1 atm = 760 torr = 760 mmHg Chapter 2 13D. Solubilities  A general rule for solubility is that “like dissolves like” in terms of comparable polarities Polar and ionic solids are usually soluble in polar solvents Polar liquids are usually miscible Nonpolar solids are usually soluble in nonpolar solvents Nonpolar liquids are usually miscible Polar and nonpolar liquids, like oil and water, are usually not soluble to large extents Chapter 2 e.g. MeOH and H2O are miscible in all proportions H3C - d + d O H hydrogen bond +H H + d O d - d Chapter 2  Hydrophobic means incompatible with water Hydrophilic means compatible with water Hydrophilic Hydrophobic portion group OH Decyl alcohol O S O Na O O A typical detergent molecule Chapter 2 13E. Guidelines for Water Solubility  Organic chemists usually define a compound as water soluble if at least 3 g of the organic compound dissolves in 100 mL of water  Usually compounds with one to three carbon atoms are water soluble, compounds with four or five carbon atoms are borderline, and compounds with six carbon atoms or more are insoluble Chapter 2  END OF CHAPTER 2  Chapter 2