Ch.2 Polar Bonds and Their Consequences. Ch.3 Organic Compounds: Alkanes and Cycloalkanes. Ch.4 Stereochemistry of Alkanes and Cycloalkanes
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2 h.1 Structure and Bonding h.2 Polar Bonds and Their onsequences h.3 Organic ompounds: Alkanes and ycloalkanes h.4 Stereochemistry of Alkanes and ycloalkanes h.5 An Overview of Organic Reactions h.6 Alkenes: Structure and Reactivity h.7 Alkenes: Reactions and Synthesis h.8 Alkynes: Introduction to Organic Synthesis h.9 Stereochemistry h.10 Alkyl alides h.11 Reaction of Alkyl alides h The E2 Elimination h.12 Mass and Infrared Spectroscopy ( )
3 Organic hemistry? mid-1700s: alchemist, no differences between substances obtained from living sources and those obtained from minerals 1770, Torbern Bergman: first express the difference of organic/inorganic ; organic chemistry = chemistry of compounds found in living organisms organic ~ vital force 1816, Michael hevreul: for the first time, one organic substance was converted into others Animal fat NaO 2 O Soap + Glycerin Soap 3 O + "Fatty acids"
4 1828, Friedrich WŐhler: convert "inorganic" salt to "organic" substance N ON Ammonium cyanate heat O 2 N N 2 Urea mid-1800s: no more vitalistic theory ~ today: no clear distinction between organic and inorganic chemistry ; unified chemistry ; but, all organic chemicals contain the element carbon
5 Organic hemistry? h.1 Structure and Bonding - the chemistry of compounds found in living organisms - the study of carbon compounds N O F Si P S l Br I - DNA, RNA, amino acid, proteins, carbohydrates - medicine, dyes, polymers, plastics, food additives, pesticides, fragrant..
6 1.1 Atomic Structures h.1 Structure and Bonding nucleus (protons + neutrons): dense, positively charged; occupies most of atomic mass d= ~10-15 m orbiting electrons: negligible mass, negatively charged ~10-10 m (1 A) atomic number (Z): # of protons mass number (A): # of protons and neutrons (neutrons are neutral) for neutral atoms 1 : 1 (proton) + 0 (neutron) = 1 2 e: 2 (proton) + 2 (neutron) = 4 atomic number = # of protons (positively charged) = # of electrons (negatively charged)
7 All the atoms in a given element have the same atomic number but they can have different mass numbers depending on how many neutrons they have. Isotope: same atomic number but different mass number (different number of neutrons) Atomic weight: the weighted average atomic mass units of an element s isotopes for hydrogen: 1 (99.985%, ), 2 (0.015%, ) for carbon: 12 (98.89%, ), 13 (1.11%, )
8 1.2 Atomic Structure: Orbitals h.1 Structure and Bonding - electron distribution: quantum mechanics - wave equation solution: wave function (ψ) or orbital: wave function of one electron atom ψ 2 ; the volume of space around a nucleus where an electron can most likely be found (probability) shapes of orbitals: s, p, d, f s orbital p orbital d orbitals
9 shells: atomic orbitals with the same principal quantum number - distribution of electrons in an atom: each orbital can occupy 2 e - 3rd shell (18 e - ) 3d 3p 3s energy 2nd shell (8e - ) 2p 2s 1st shell (2 e - ) 1s
10 shapes of 2 p orbitals y y y x x x z z z 2p x orbital 2p y orbital 2p z orbital node: a region of zero electron density nodal plane:
11 1.3 Atomic Structure: Electron onfiguration ground-state electron configuration: lowest-energy arrangement rule 1: The lowest-energy orbitals fill up first (Aufbau principle) 1s 2s 2p 3s 3p 4s 3d Energy 2p 2s 1s
12 rule 2: Only two electrons can occupy an orbital and they must be of opposite spin (Pauli exclusion principle): electrons spins ( ) or ( ) 2p 2s 1s
13 rule 3: If two or more orbitals of equal energy are available, one electron occupies each until all orbitals are half-full. Only then does a second electron occupy one of the orbitals (unt s rule): The electrons in the half-filled orbitals all have the same spin 2p 2s 1s
14 ground-state electron configurations of some elements ydrogen 1s atomic number 1 atomic number arbon 6 2p 2s 1s
15 ground-state electron configuration of oxygen atomic number Oxygen 8 2p 2s 1s
16 Practice hloride: atomic number 17 3p 3s 2p 2s 1s
17 1.4 Development of hemical Bonding Theory In 1858, Kekule-ouper: tetravalent carbon In 1865, Kekule-ouper: multiple bonding, ring systems (2-dimensional) In 1874, van't off: tetrahedral carbon (3-dimensional) in plane dashed line (behind the plane) heavy wedged line (out of the plane)
18 1.5 ovalent Bonds h.1 Structure and Bonding Why do atoms bond together? Because the compound is more stable, has less energy than the separate atoms ow can bonds be described electronically? difficult to answer Octet rule: 8 electrons in valence shell, special stability ; noble-gas configuration Ne (2 + 8); Ar ( ); Kr ( )
19 group 1A (alkali metal): loose 1 electron Na Na + + e - 1s 2 2s 2 2p 6 3s 1 1s 2 2s 2 2p 6 group 7A (halogen): gain 1 electron l + e - l - 1s 2 2s 2 2p 5 1s 2 2s 2 2p 6
20 ionic bond: held together by electrostatic attraction Na + l - ow about the middle elements? X e - 1s 2 2s 2 2p 2 1s 2 - satisfy octet by sharing electrons: covalent bond
21 molecule: the neutral collection of atoms held together by covalent bonds Lewis structure or electron-dot structure: - valence electrons are represented by dots 1s 1 2s 2 2p 2 2s 2 2p 3 2s 2 2p 4 2s 2 2p 5 N O F
22 O N F 4 2 O N 3 F O 3 O
23 Br l F O one bond two bond N B three bond four bond
24 nonbonding electrons (lone-pair electrons); valence electrons that are not used for bonding N 3 N N Kekule structure (line-bond structure)
25 Lewis structure O N F 4 2 O N 3 F Kekule' structure O N O 4 2 O N 3 3 O
26 Practice P? Gel? P P 3 l l Ge l l Gel 4 P : 5A Ge : 4A All?? l? l l Al l All 3 l l 2 l 2 Al : 3A : 4A
27 1.6 Valence Bond Theory and Molecular Orbital Theory ow does the electron sharing occur? two models for covalent bond formation valence bond theory "electrons are localized around the bond" molecular orbital theory "electrons are delocalized over the molecule" explains well σ-bonds explains well π-bonds
28 Valence Bond Theory key ideas ovalent bonds are formed by overlap of two atomic orbitals, each of which contains one electron. The spins of the two electrons are opposite. Each of the bonded atoms retains its own atomic orbitals, but the electron pair in the overlapping orbitals is shared by both atoms. The greater the amount of orbital overlap, the stronger the bond.
29 + 1s 1s 2 molecule circular cross-section - - bond is cylindrically symmetrical sigma ( ) bond: bonds formed by the head-on overlap of two orbitals along a line drawn between the nuclei
30 2 2 energy two hydrogen atoms 2 molecule 436 kj/mol (104 kcal/mol) released when bond formed absorbed when bond breaks ow close are the two nuclei in the 2 molecule? too close: nuclei repel too far: unable to share electrons optimum distance for maximum stability: bond length
31 (too close) + energy 0 (too far) _ 74 pm bond length A plot of energy versus internuclear distance for two hydrogen atoms.
32 1.7 ybridization: sp 3 Orbitals and the Structure of Methane - valence bond theory ydrogen molecule is simple, but complicate for tetravalent carbon atom arbon uses two kinds of orbitals (2s 2 2p 2 ) to form bonds but four identical bonds (tetrahedral). ow? ybridization (1931, Linus Pauling): mathematically explained how an s orbital and three p orbitals can combine (sp 3 orbitals)
33 ybridization 2s 2p y hybridization 4 identical tetrahedral sp 3 orbitals 2p x 2p z an sp 3 orbital
34 hybridized sp 3 orbitals are unsymmetrical: one of two lobes is much larger, and form stronger bonds than unhybridized s or p orbitals unsymmetrical sp 3 orbitals: two lobes of p-orbital have opposite signs hybridization + _ + _ + 2s 2p x an sp 3 orbital
35 bond angle (tetrahedral) o 110 pm bond length - bond strength in methane: 438 kj/mol (105 kcal/mol)
36 1.8 The Structure of Ethane -overlap of two sp 3 orbitals h.1 Structure and Bonding 154 pm o - bond strength in ethane: 420 kj/mol (100 kcal/mol) - bond strength in ethane: 376 kj/mol (90 kcal/mol)
37 1.9 ybridization: sp 2 Orbitals and the Structure of Ethylene 2 4 ; double bond, planar 133 pm pm o o
38 s + 2 (p) 3 (sp 2 ) p 3 identical sp 2 orbitals π-bond - bond strength in ethylene: 444 kj/mol (106 kcal/mol) = bond strength in ethylene: 611 kj/mol (146 kcal/mol) σ-bond
39
40 Practice Line-bond structure, hybridization of carbon? 2 O O O sp 2 Lewis structure Line-bond structure
41 Practice Line-bond structure, hybridization of carbon? 3 = 2 sp 3 sp 2
42 1.10 ybridization: sp Orbitals and the Structure of Acetylene 2 2 ; triple bond, linear 106 pm 120 pm 180 o
43 s + 1 (p) 2 (sp) h.1 Structure and Bonding 2 identical sp orbitals p p π-bond π-bond σ-bond - bond strength in acetylene: 552 kj/mol (132 kcal/mol) bond strength in acetylene: 835 kj/mol (200 kcal/mol)
44
45 omparison of / and / bonds Molecule Bond Bond Strength kj/mol kcal/mol Bond Length (pm) 4 (sp 3 ) (sp 3 )-(sp 3 ) (sp 3 ) (sp 2 )=(sp 2 ) = 2 (sp 2 ) (sp) (sp) (sp)
46 Summary of ybridizations tetrahedral: sp 3 hybrid planar: sp 2 hybrid 3 3 linear: sp hybrid 3 3
47 Practice Line-bond structure, hybridization of carbon? 3 sp 3 sp
48 Practice 2 == 2 ybrid, Geometry? sp sp 2 sp 2
49 1.11 ybridization of Other Atoms: Nitrogen and Oxygen covalent bonds of other elements can also be described by hybridization sp 3 N pm bond length N-: 449 kj/mol (107 kcal/mol) o N 3 N
50 sp 3 O O-: 498 kj/mol (119 kcal/mol) 95.8 pm bond length o 2 O O
51 Practice Line-bond structure, hybridizations? Formaldimine, 2 N sp 2 N N sp 2
52 Practice Geometry? 3 O 3 3 N P O 3 N 3 3 P
53 1.12 ybridization of arbon Species arbocation sp 2 planar
54 1.12 Molecular Orbital Theory key ideas Molecular orbitals describe regions of space in a molecule where electrons are most likely to be found, and they have a specific size, shape, and energy level. Molecular orbitals are formed by combining atomic orbitals. The number of MO's formed is the same as the number of atomic orbitals combined. bonding MO: lower in energy than the starting atomic orbitals antibonding MO: higher in energy nonbonding MO: same energy
55 antibonding MO (unfilled) energy 1s orbital 1s orbital bonding MO (filled)
56 Molecular orbital description = -bond Actually, p-orbitals have two lobes with opposite signs. nodal plane combine π-antibonding (π*) MO π-bonding (π*) MO
57 Work hemical Toxicity and Risk All chemicals are toxic to some extent, and the difference between help and harm is matter of degree. LD 50 value: the amount of a substance per kilogram body weight that is lethal to 50% of the test animals. The lower the value, the more toxic the substance. Substance Aflatoxin B1 Aspirin hloroform Ethyl alcohol Formaldehyde Sodium cyanide Sodium cyclamate LD 50 (g/kg) 4 x x
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