Chapter 1. Introduction to Organic Chemistry

Transcription

1 Chapter 1 Introduction to rganic Chemistry

2 rganic Chemistry - General Description A. The Study of Carbon Compounds B. rganic reminds us of plant or animal origins 1. Natural medicines: morphine, penicillin 2. Natural fibers: cotton, silk, wool 3. Foodstuffs: Fats, carbohydrates, proteins, vitamins 4. Natural rubber C. Man-made substances can also be organic 1. Drugs: xylocaine, aspirin, acetaminophen 2. Fibers: nylon, dacron, rayon 3. Polymers: Saran wrap, polyesters, teflon, nylon 4. Synthetic rubber, synfuels

3 Differences Between rganic and Inorganic Compounds Inorganic rganic Elements Present Metals and Nometals Mostly Carbon Bonding Covalent and Ionic Mostly Covalent Particles INS & molecules ions & MLECULES Melting Points Relatively igh Relatively Low Boiling Points Relatively igh Relatively Low Electrolytes STRNG to weak weak to NN

4 Comparison of Physical Properties of rganic and Inorganic Compounds Name Salt Ethyl alcohol Benzene Formula NaCl C26 C66 rganic/inorganic Inorganic rganic rganic Melting Point 804 ºC -117 ºC 5 ºC Boiling Point 1413 ºC 78 ºC 80 ºC Burns in 2? No Yes Yes Water Soluble? Yes Yes No

5 Review of General Chemistry Chemical Bonding

6 Types of Bonds Types of Atoms metals to nonmetals nonmetals to nonmetals metals to metals Type of Bond Ionic Covalent Metallic Bond Characteristic electrons transferred electrons shared electrons pooled

7 Types of Bonding

8 Covalent Bonds Nonmetal atoms have relatively high ionization energies. It is difficult to remove electrons from them. When nonmetals bond together, it is better in terms of potential energy for the atoms to share valence electrons.

9 Lewis Bonding Theory Emphasizes valence electrons to explain bonding Lewis structures - Electron Dot Structures Lewis structures allow us to predict many properties of molecules - molecular stability, shape, size, polarity

10 Dot Structures Emphasize Valence Electrons The column number on the Periodic Table will tell you how many valence electrons a main group atom has. We represent the valence electrons of main-group elements as dots surrounding the symbol for the element. We use the symbol of element to represent nucleus and inner electrons. We use dots around the symbol to represent valence electrons. IA VIIIA IIA III A IVA VA VIA VIIA e Li Be B C N F Ne Na Mg Al Si P S Cl Ar

11 Lewis Structures

12 Lewis Bonding Theory Atoms bond because it results in a more stable electron configuration. (lower potential energy) Atoms bond together by either transferring or sharing electrons. Usually this results in all atoms obtaining an outer shell with eight electrons (octet rule) There are some exceptions to this rule.

13 Common bonding patterns: Lewis Structures Be = 2 bonds & 0 lone pairs B = 3 bonds & 0 lone pairs C = 4 bonds & 0 lone pairs N = 3 bonds & 1 lone pair = 2 bonds & 2 lone pairs and halogens = 1 bond Be B C N F Structures that result in bonding patterns different from the common may have formal charges.

14 Single Covalent Bonds When two atoms share one pair of electrons it is called a single covalent bond (2 electrons) ne atom may use more than one single bond to fulfill its octet F F F F F F

15 Multiple Covalent Bonds N N N N When two atoms share two pairs of electrons the result is called a double covalent bond (four electrons) When two atoms share three pairs of electrons the result is called a triple covalent bond (six electrons)

16 Drawing Lewis structures of molecules 1. Write skeletal structure 2. Count valence electrons 3. Attach atoms together with pairs of electrons, and subtract from the total 4. Complete octets, outside-to-inside 5. Give extra electrons to the central atom 6. If central atom does not have octet, bring in electrons from outside atoms to share

17 N3 N N,,,, N N

18 P24 P P 4(1) + 2(5) = 14 valence electrons P P = 4 electrons remaining P P 4-4 = 0 electrons remaining

19 C24 C C C,C,,,,, C C C C

20 N3 (N2) N N N 1 + 3(6) + 5 = 24 valence electrons 24-8 = 16 electrons remaining = 0 electrons remaining N

21 C2 C = 16 valence electrons C 16-4 = 12 electrons remaining C = 0 electrons remaining C

22 N2 - N 2(6) = 18 electrons N 18-4 = 14 electrons remaining N = 0 electrons remaining N

23 SeF2 F Se F 2(7) = 26 valence electrons F Se F 26-6 = 20 electrons remaining F Se F = 0 electrons remaining

24 3P4 P 3(1) + 4(6) + 5 = 32 electrons P = 18 electrons remaining P = 0 electrons remaining

25 S3 2- S 6 + 3(6) + 2 = 26 electrons S 26-6 = 20 electrons remaining S = 0 electrons remaining

26 Structural Isomerism

27 Isomerism n-butane butane iso-butane 2-methylpropane

28 Isomerism A. A molecular formula may not convey a unique structure. B. Isomers - Compounds with identical molecular formulas, but different structural formulas. C. Example #1: C410 Compound C C C C C C C C Name Butane Isobutane Melting point -138 o C -160 o C Boiling point 0 o C -12 o C Density g/ml g/ml

29 Isomerism Example #2: C242 Compound C C C C Name Acetic Acid Methyl formate Melting point 16.6 o C -99 o C Boiling point 118 o C 31 o C Density 1.05 g/ml 0.97 g/ml Toxicity

30 Compounds with the Formula C410 3 C C 2 C 2 C 2 C 3 C 2 C 2 C 3 C 3 C 2 C C 3 C 3 C 2 C C 3 C 2 C 2 C 3 C 3 C 3 C 3 C C 3 3 C C 3 C C 3

31 Compounds with the Formula C48 C 3 3 C C 2 C C 3 3 C C C 2 C C 2 2 C C 2 C C 3 C C 3 C 3 2 C C 2 C 2 2 C C C 2 C2 C C 2 C C 2

32 Formal Charges

33 Formal Charge During bonding, atoms may end with more or fewer electrons than the valence electrons they brought in order to fulfill octets. This results in atoms having a formal charge. FC = (#valence e ) (#nonbonding e ) (½ #bonding e ) Sum of all the formal charges in a molecule = 0 Sum of all the formal charges in an ion = ionic charge

34 Writing Lewis Formulas of Molecules Assign formal charges to the atoms (fc = valence e lone pair e ½ bonding e ) N N For, fc = 1-0-½(2) = 0 For, fc = 6-4-½(4) = 0 For N, fc = 5-0-½(8) = +1 For, fc = 6-4-½(4) = 0 For, fc = 6-6-½(2) = -1

35 Common Bonding Patterns B C N C + N + + F F + B C N F

36 Practice - Assign formal charges P P C [ N ] P F Se F [ S ] (fc = valence e lone pair e ½ bonding e )

37 Minimizing Formal Charges Through Resonance - SeF2 1. First draw Lewis structure that maximizes octets 2. Assign formal charges 3. Move electron pairs from atoms with ( ) formal charge toward atoms with (+) formal charge 4. If (+) fc atom is 2 nd row, only move in electrons if you can move out electron pairs from multiple bond. -1 F Se F If (+) fc atom 3 rd row or below, keep bringing in electron pairs to reduce the formal charge, even if get expanded octet. F Se F

38 Minimizing Formal Charges Through Resonance - 3P4 1. First draw Lewis structure that maximizes octets 2. Assign formal charges 3. Move electron pairs from atoms with ( ) formal charge toward atoms with (+) formal charge 4. If (+) fc atom is 2 nd row, only move in electrons if you can move out electron pairs from multiple bond. P If (+) fc atom 3 rd row or below, keep bringing in electron pairs to reduce the formal charge, even if get expanded octet. P 0 0

39 Minimizing Formal Charges Through Resonance - S First draw Lewis structure that maximizes octets 2. Assign formal charges 3. Move electron pairs from atoms with ( ) formal charge toward atoms with (+) formal charge 4. If (+) fc atom is 2 nd row, only move in electrons if you can move out electron pairs from multiple bond. 5. If (+) fc atom 3 rd row or below, keep bringing in electron pairs to reduce the formal charge, even if get expanded octet. -1 [ S ] [ S ]

40 Polar Covalent Bonds

41 Polar Covalent Bonding Unequal sharing of the electrons 1) ne atom pulls the electrons in the bond closer to its side 2) ne end of the bond has higher electron density than the other The result is a polar covalent bond Bond polarity is indicated by (δ-) and (δ+)

42 Bond Polarity Most bonds have some degree of sharing and some degree of ion formation to them. Bonds are classified as covalent if the amount of electron transfer is insufficient for the material to display the classic properties of ionic compounds. If the sharing is unequal enough to produce a dipole in the bond, the bond is classified as polar covalent.

43 Electronegativity The ability of an atom to attract bonding electrons to itself Increases across period ( ) Decreases down group( ) 1) Fluorine is the most electronegative element. 2) Francium is the least electronegative element. 3) Noble gas atoms are not assigned values. 4) This is opposite to the trend in atomic size.

44 Electronegativity Scale

45 Electronegativity Difference ( EN) and Bond Classification The larger the difference in electronegativity, the more polar the bond. EN Bond Classification 0 pure covalent nonpolar covalent polar covalent >2.0 ionic

46 Bond Polarity EN Cl = = 0 Pure Covalent EN Cl = 3.0 EN = = 0.9 Polar Covalent EN Cl = 3.0 EN Na = = 2.1 Ionic

47 Electrostatic Potential Diagrams