Presentation is loading. Please wait.

Presentation is loading. Please wait.

1. Structure and Bonding Why this chapter?

Similar presentations


Presentation on theme: "1. Structure and Bonding Why this chapter?"— Presentation transcript:

1 1. Structure and Bonding Why this chapter?
• Review ideas from general chemistry: atoms, bonds, molecular geometry

2 What is Organic Chemistry?
Living things are made of organic chemicals Proteins that make up hair DNA, controls genetic make-up Foods, medicines Examine structures below

3 Origins of Organic Chemistry
Foundations of organic chemistry from mid-1700’s. Compounds obtained from plants, animals hard to isolate, and purify. Compounds also decomposed more easily. Torben Bergman (1770) first to make distinction between organic and inorganic chemistry. It was thought that organic compounds must contain some “vital force” because they were from living sources.

4 Because of “Vital force”, it was thought that organic compounds could not be synthesized in laboratory like inorganic compounds. 1816, Chervrut showed that not to be the case, he could prepare soap from animal fat and an alkali 1828, Wholer showed that it was possible to convert inorganic salt ammonium cyanate into organic substance “urea”

5 90% of more than 30 million chemical compounds contain carbon.
• Organic chemistry is study of carbon compounds. • Why is it so special? 90% of more than 30 million chemical compounds contain carbon. Examination of carbon in periodic chart answers some of these questions. Carbon is group 4A element, it can share 4 valence electrons and form 4 covalent bonds.

6 1.1 Atomic Structure Structure of an atom
Positively charged nucleus (very dense, protons and neutrons) and small (10-15 m) Negatively charged electrons are in a cloud (10-10 m) around nucleus Diameter is about 2  m (200 picometers (pm)) [the unit angstrom (Å) is m = 100 pm]

7 X The Atomic Symbol A = Atomic mass Z = Atomic #
= # p + # n C = Charge = + or - values X A C Z # Z = Atomic # # p = # e # = Number of atoms in a formula.

8 X - - - - - - The Atomic Symbol 6 A = Atomic mass 12 Z = Atomic number
= # protons + # neutrons 6 6 - - X 12 - - 6 + + Z = Atomic number = # protons = # electrons - + + + + -

9 C - - - - - - 6 The Atomic Symbol A = Atomic mass 12 Z = Atomic number
= # protons + # neutrons - - C 12 - - 6 + + Z = Atomic number = # protons = # electrons - + + + + -

10 Na A = Atomic mass Z = Atomic # = p = 11 C = Charge = +1
The Atomic Symbol A = Atomic mass = p + n = 23 C = Charge = +1 11 12 23 1+ Na 11 # = 1 atom in formula. Z = Atomic # = p = 11 Sodium

11 Why is the atomic weight on the tables not a whole #?
47 Atomic number Name of the element Elemental Symbol Silver Ag 107.87 Atomic mass (weight) Atomic weight = The average, relative mass of an atom in an element.

12 - - - Isotopes of Hydrogen
Isotopes = Atoms of the same element but having different masses. 1 2 1 3 1 H H H - - - + + + Protium 99.99% Tritium Trace % Deuterium 0.01%

13 Average Atomic weight of Hydrogen
Isotopes of Hydrogen Isotopes = Atoms of the same element but having different masses. 1 2 1 3 1 H H H - - - + + + Average Atomic weight of Hydrogen = amu

14 Average Atomic weight of C= 12.011 amu
Isotopes of Carbon 12 13 14 C C C 6 6 6 - - - - - - - - + + + - + + + + + + + + + - - + + + + + + - - - - - - - 98.89% 1.11% Trace % Average Atomic weight of C= amu

15 C - - Radioactive Isotopes 6 14 H H-3 C-14 + + Nucleus is unstable
So falls apart (decays) Giving radioactive particles

16 Electronic arrangement
A new layer is added for each row or period in the table.

17 fill layers around nucleus
Electron arrangement 24 12 Mg Electrons fill layers around nucleus Low  High 32 18 8 2 Shells = Energy levels

18 IA IIA 1 H 4 2 He 7 3 Li 9 4 Be 2, 1 2, 2

19 IA IIA IIIA 1 H 11 5 B 7 3 Li 9 4 Be 2, 1 2, 2 2, 3

20 IIIA IVA VA 11 5 B 12 6 C 13 7 N 2, 3 2, 4 2, 5

21 H He Be Li Ne Mg Ar Na 1 4 2 9 7 4 3 24 40 12 18 IA IIA VIIIA 20 10
2, 1 2, 2 2, 8 24 12 Mg 40 18 Ar 23 11 Na 2, 8, 8 2, 8, 1 2, 8, 2

22 H Be Li B Mg Al Na 1 9 7 4 3 Valence electrons Where most chemical
Reactions occur. 1 H 2 3 9 4 Be 7 3 Li 11 5 B 2, 3 2, 1 2, 2 24 12 Mg 27 13 Al 23 11 Na 2, 8, 3 2, 8, 1 2, 8, 2

23 H He Be Li Ne Mg Ar Na Octet Rule 1 4 2 9 4 7 3 24 40 12 18 1 8 2 20
10 Ne 7 3 Li 2, 1 2, 2 2, 8 24 12 Mg 40 18 Ar 23 11 Na 2, 8, 8 2, 8, 1 2, 8, 2

24 The octet rule Atoms are most stable if they have a filled or empty outer layer of electrons. Except for H and He, a filled layer contains 8 electrons - an octet. Atoms gain, lose or share electrons to make a filled or empty outer layer. Atoms gain, lose or share electrons based on what is easiest.

25 1.2 Atomic Structure: Orbitals
Quantum mechanics: describes electron energies and locations by a wave equation Wave function solution of wave equation Each wave function is an orbital,

26 Orbitals A plot of y 2 describes where electron most likely to be
Electron cloud has no specific boundary so we show most probable area p (3) s (1) d (5) f (7) An f orbital 4 different kinds of orbitals for e-s s and p orbitals most important in organic & biochem

27 Orbitals and Shells Orbitals are grouped in shells of increasing size and energy Different shells contain different numbers and kinds of orbitals Each orbital can be occupied by two electrons

28 Orbitals and Shells 1st shell contains one s orbital, denoted 1s, holds 2 e-s 2nd shell contains one s orbital (2s) and three p orbitals (2p), 8 electrons 3rd shell contains an s orbital (3s), three p orbitals (3p), and five d orbitals (3d), 18 electrons

29 p-Orbitals In each shell there are three perpendicular p orbitals, px, py, and pz, of equal energy Lobes of a p orbital are separated by region of zero electron density, a node

30 1.3 Atomic Structure: Electron Configurations
Ground-state electron configuration (lowest energy arrangement) of an atom lists orbitals occupied by its electrons. Rules: 1. Lowest-energy orbitals fill first: 1s  2s  2p  3s  3p  4s  3d (Aufbau (“build-up”) principle) 2. Electrons act as if they were spinning around an axis. Electron spin can have only two orientations, up  and down . Only two electrons can occupy an orbital, and they must be of opposite spin (Pauli exclusion principle) to have unique wave equations 3. If two or more empty orbitals of equal energy are available, electrons occupy each with spins parallel until all orbitals have one electron (Hund's rule).

31 1.3 Atomic Structure: Electron Configurations
The Aufbau (“build-up”) principle Electrons fill from the low  high. fill n = 1 before n = 2 , fill s before p ... s p d f n = 4 Pauli Exclusion principle n = 3 Each subshell contains orbitals which can hold a max of 2 e-s e’s in same orbital must be of opposite spins: 1  & 1 . n = 2 n = 1 Hund’s Rule Electrons don’t share same orbital unless they need to. (i.e. no pairing until each orbital of the set has an e-)

32 Major trends in electron filling
5p 4d 4f 4s 4p 3d 3s 3p 2s 2p 1s Major trends in electron filling Exceptions: Fill 4s before 3d Fill 5s before 4 d Fill 5p before 4f This is why transition metals are assigned as B group elements.

33 Electron Configuration
1s __ 2s __ 2p __ __ __ 3s __ 3p __ __ __ 4s __ 3d __ __ __ __ __ 5s __ 4p __ __ __ 4d __ __ __ __ __ 5p __ __ __ 6s __ 5d __ __ __ __ __ 4f __ __ __ __ __ __ __ 7 3 Li 1s22s1 x y z x y z

34 Electron Configuration
1s __ 2s __ 2p __ __ __ 3s __ 3p __ __ __ 4s __ 3d __ __ __ __ __ 5s __ 4p __ __ __ 4d __ __ __ __ __ 5p __ __ __ 6s __ 5d __ __ __ __ __ 4f __ __ __ __ __ __ __ 16 8 O 1s22s22p4 x y z

35 Electron Configuration
1s __ 2s __ 2p __ __ __ 3s __ 3p __ __ __ 4s __ 3d __ __ __ __ __ 5s __ 4p __ __ __ 4d __ __ __ __ __ 5p __ __ __ 6s __ 5d __ __ __ __ __ 4f __ __ __ __ __ __ __ 30 Zn x y z 1s22s22p63s23p64s23d10 x y z [Ar] 4s23d10 [Ar] 3d104s2

36 Classification by sublevels
p H He d Li Be B C N O F Ne Na Mg Al Si P S Cl Ar K Ca Sc Hf Zr Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr Rb Sr Y Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe Cs Ba Ls Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn Fr Ra Ac Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu f Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr 3 - 25

37 Using the periodic table to find sublevels
1 H Li Na Cs Rb K Ba Fr He Be Mg Sr Ca Ra p 2 2 Tl Rn At Po Bi Pb In Xe I Te Sb Sn Ga Kr Br Se As Ge Al Ar Cl S P Si B Ne F O N C 3 d 3 4 3 Sc Ti V Cr Mn Fe Co Ni Cu Zn 4 5 4 Y Zr Nb Mo Tc Ru Rh Pd Ag Cd 5 5 Ls Hf Ta W Re Os Ir Pt Au Hg 6 6 6 Ac 7 4 Gd Cm Tb Bk Sm Pu Eu Am Nd U Pm Np Ce Th Pr Pa Yb No Lu Lr Er Fm Tm Md Dy Cf Ho Es f 5

38 Learning Check: 1s22s22px22py02pz0 1s22s22px12py12pz0
Which of the following represents the ground-state electronic configuration for carbon (atomic number 6)? 1s22s22px22py02pz0 1s22s22px12py12pz0 1s22s12px12py1 2pz1 1s12s22px12py1 2pz1 1s22s22px02py02pz2

39 Solution: 1s22s22px22py02pz0 1s22s22px12py12pz0 1s22s12px12py1 2pz1
Which of the following represents the ground-state electronic configuration for carbon (atomic number 6)? 1s22s22px22py02pz0 1s22s22px12py12pz0 1s22s12px12py1 2pz1 1s12s22px12py1 2pz1 1s22s22px02py02pz2

40 H He Be Li Ne Ar Mg Na Octet Rule 1 4 2 9 7 4 3 1s2 1s1 20 10 1s2, 2s2
1s2, 2s2 2p6 1s2, 2s1 40 18 Ar 24 12 Mg 23 11 Na 1s2, 2s2 2p6, 3s2 3p6 1s2, 2s2 2p6, 3s1 1s2, 2s2 2p6, 3s2 [Ne] 3s1

41 H Li Na K H Li Lewis Structures Show only Valence Electrons Na 1 7 3
23 11 Na K

42 He O F N Si S P Ca H Li C Na K Se B Ne Be Al Cl Ar Mg Kr Ga As Br Ge 1
Nonmetals Share e-s with other nonmetals 1 8 H He 2 3 4 5 6 7 C O Li B N F Ne Be Si Al S Cl Ar Mg P Na Metals give e-s to nonmetals Kr Ca Ga As Se Br K Ge

43 Common ions Representative Elements Transition Elements 1+ 4+ 4- H He
2+ 3+ 3- 2- 1- Transition Elements Li Be B C N O F Ne Variable Na Mg Al Si P S Cl Ar K Ca Hg Au Hf Ls Pt Ir Os Re W Ta Cd Ag Zr Y Pd Rh Ru Tc Mo Nb Ac Zn Cu Ti Sc Ni Co Fe Mn Cr V Gd Cm Tb Bk Sm Pu Eu Am Nd U Pm Np Ce Th Pr Pa Yb No Lu Lr Er Fm Tm Md Dy Cf Ho Es Ga Ge As Se Br Kr Rb Sr In Sn Sb Te I Xe Cs Ba Tl Pb Bi Po At Rn Fr Ra 12 12

44 Na + Cl Na+ + Cl _ Formation of NaCl e- moves from Metal  Nonmetal
Stable octets _ Na + Cl Na+ + Cl Metal Cation Nonmetal Anion + and - ions attract to form an ionic bond. 15 15

45 Ionic compounds NaCl sodium chloride Not individual molecules
Form crystal arrays Ions touch many others Formula represents the average ion ratio NaCl sodium chloride Na Cl Cl Na Cl Na 18 18

46 He O F N Si S P Ca H Li C Na K Se B Ne Be Al Cl Ar Mg Kr Ga As Br Ge 1
Nonmetals Share e-s with other nonmetals 1 8 H He 2 3 4 5 6 7 C O Li B N F Ne Be Si Al S Cl Ar Mg P Na Metals give e-s to nonmetals Kr Ca Ga As Se Br K Ge

47 Covalent Bonds H H H + Cl Cl + Cl O O O + N N N +

48 Covalent Bonds H H-H H2 Cl Cl-Cl Cl2 O O=O O2 N2 N N N

49 May modify rules to improve sound. ie - monoxide not monooxide.
Covalent Bonds CO Carbon monoxide C O C O CO2 Carbon dioxide O C O O=C=O May modify rules to improve sound. ie - monoxide not monooxide.

50 Naming covalent compounds
Review: Naming covalent compounds CO CO2 N2O5 SiO2 ICl3 P2O5 CCl4 carbon monoxide carbon dioxide dinitrogen pentoxide silicon dioxide iodine trichloride diphophorous pentoxide carbon tetrachloride May modify rules to improve the sound. Example - use monoxide not monooxide. 34 34

51 Shape CO2 O C O O=C=O Linear Shape (180o)

52 Shape BF3 F B F F F B (120o) Trigonal Planar

53 1.4 Development of Chemical Bonding Theory
Kekulé and Couper independently observed that carbon always has four bonds van't Hoff and Le Bel proposed that the four bonds of carbon have specific spatial directions Atoms surround carbon as corners of a tetrahedron

54 O=C=O Shape Linear Trigonal planar e-’s in 2 directions = 180o

55 Shape N H N H Pyramidal Trigonal Pyramid Configuration of Atoms
e-’s in 4 directions = 109.5o N H N H Pyramidal (109.5o) Tetrahedral Configuration of Electrons Trigonal Pyramid Configuration of Atoms

56 Tetrahedral electron-pair Geometries
Pyramidal Bent

57

58 Non-bonding electrons
Valence electrons not used in bonding are called nonbonding electrons, or lone-pair electrons Nitrogen atom in ammonia (NH3) Shares six valence electrons in three covalent bonds and remaining two valence electrons are nonbonding lone pair

59 Learning Check How many hydrogen atoms does phosphorus bond to in forming phosphine, PH? Hint: Draw the electron dot structure and then use Hydrogen to provided the needed bonded electrons

60 Solution How many hydrogen atoms does phosphorus bond to in forming phosphine, PH? H P H 3 P H H Hint: Draw the electron dot structure and then use Hydrogen to provided the needed bonded electrons

61 Learning Check How many hydrogen atoms does Carbon need in forming chloroform, CH?Cl C Cl H

62 Solution 3 C Cl H H C Cl H Cl C H H
How many hydrogen atoms does Carbon need in forming chloroform, CH?Cl 3 C Cl H H C Cl H Cl C H H

63 1.5 The Nature of Chemical Bonds: Valence Bond Theory
Covalent bond forms when two atoms approach each other closely so that a singly occupied orbital on one atom overlaps a singly occupied orbital on the other atom Two models to describe covalent bonding. Valence bond theory, Molecular orbital theory Valence Bond Theory: Electrons are paired in the overlapping orbitals and are attracted to nuclei of both atoms H–H bond results from the overlap of two singly occupied hydrogen 1s orbitals H-H bond is cylindrically symmetrical, sigma (s) bond

64 Bond Energy Reaction 2 H·  H2 releases 436 kJ/mol
Product has 436 kJ/mol less energy than two atoms: H–H has bond strength of 436 kJ/mol. (1 kJ = kcal; 1 kcal = kJ)

65 Bond Length Distance between nuclei that leads to maximum stability
If too close, they repel because both are positively charged If too far apart, bonding is weak

66 1.6 sp3 Orbitals and the Structure of Methane
Carbon has 4 valence electrons (2s2 2p2) In CH4, all C–H bonds are identical (tetrahedral) sp3 hybrid orbitals: s orbital and three p orbitals combine to form four equivalent, unsymmetrical, tetrahedral orbitals (sppp = sp3), Pauling (1931)

67 The Structure of Methane
sp3 orbitals on C overlap with 1s orbitals on 4 H atoms to form four identical C-H bonds Each C–H bond has a strength of 436 (438) kJ/mol and length of 109 pm Bond angle: each H–C–H is 109.5°, the tetrahedral angle.

68 1.7 sp3 Orbitals and the Structure of Ethane
Two C’s bond to each other by s overlap of an sp3 orbital from each Three sp3 orbitals on each C overlap with H 1s orbitals to form six C–H bonds C–H bond strength in ethane 423 kJ/mol C–C bond is 154 pm long and strength is 376 kJ/mol All bond angles of ethane are tetrahedral

69 1.8 sp2 Orbitals and the Structure of Ethylene
sp2 hybrid orbitals: 2s orbital combines with two 2p orbitals, giving 3 orbitals (spp = sp2). This results in a double bond. sp2 orbitals are in a plane with120° angles Remaining p orbital is perpendicular to the plane

70 Bonds From sp2 Hybrid Orbitals
Two sp2-hybridized orbitals overlap to form a s bond p orbitals overlap side-to-side to formation a pi () bond sp2–sp2 s bond and 2p–2p  bond result in sharing four electrons and formation of C-C double bond Electrons in the s bond are centered between nuclei Electrons in the  bond occupy regions are on either side of a line between nuclei

71 Structure of Ethylene H atoms form s bonds with four sp2 orbitals
H–C–H and H–C–C bond angles of about 120° C–C double bond in ethylene shorter and stronger than single bond in ethane Ethylene C=C bond length 134 pm (C–C 154 pm)

72 1.9 sp Orbitals and the Structure of Acetylene
C-C a triple bond sharing six electrons Carbon 2s orbital hybridizes with a single p orbital giving two sp hybrids two p orbitals remain unchanged sp orbitals are linear, 180° apart on x-axis Two p orbitals are perpendicular on the y-axis and the z- axis

73 Orbitals of Acetylene Two sp hybrid orbitals from each C form sp–sp s bond pz orbitals from each C form a pz–pz  bond by sideways overlap and py orbitals overlap similarly

74 Bonding in Acetylene Sharing of six electrons forms C ºC
Two sp orbitals form s bonds with hydrogens

75

76 1.10 Hybridization of Nitrogen and Oxygen
Elements other than C can have hybridized orbitals H–N–H bond angle in ammonia (NH3) 107.3° C-N-H bond angle is ° N’s orbitals (sppp) hybridize to form four sp3 orbitals One sp3 orbital is occupied by two nonbonding electrons, and three sp3 orbitals have one electron each, forming bonds to H and CH3.

77 1.11 Molecular Orbital Theory
A molecular orbital (MO): where electrons are most likely to be found (specific energy and general shape) in a molecule Additive combination (bonding) MO is lower in energy Subtractive combination (antibonding) MO is higher energy

78 Molecular Orbitals in Ethylene
The  bonding MO is from combining p orbital lobes with the same algebraic sign The  antibonding MO is from combining lobes with opposite signs Only bonding MO is occupied

79 1.12 Drawing Structures Drawing every bond in organic molecule can become tedious. Several shorthand methods have been developed to write structures. Condensed structures don’t have C-H or C-C single bonds shown. They are understood. e.g.

80 3 General Rules: 1) Carbon atoms aren’t usually shown. Instead a carbon atom is assumed to be at each intersection of two lines (bonds) and at the end of each line. 2) Hydrogen atoms bonded to carbon aren’t shown. 3) Atoms other than carbon and hydrogen are shown (See table 1.3).

81 Summary Organic chemistry – chemistry of carbon compounds
Atom: positively charged nucleus surrounded by negatively charged electrons Electronic structure of an atom described by wave equation Electrons occupy orbitals around the nucleus. Different orbitals have different energy levels and different shapes s orbitals are spherical, p orbitals are dumbbell-shaped Covalent bonds - electron pair is shared between atoms Valence bond theory - electron sharing occurs by overlap of two atomic orbitals Molecular orbital (MO) theory, - bonds result from combination of atomic orbitals to give molecular orbitals, which belong to the entire molecule

82 Summary (cont’d) Sigma (s) bonds - Circular cross-section and are formed by head-on interaction Pi () bonds – “dumbbell” shape from sideways interaction of p orbitals Carbon uses hybrid orbitals to form bonds in organic molecules. In single bonds with tetrahedral geometry, carbon has four sp3 hybrid orbitals In double bonds with planar geometry, carbon uses three equivalent sp2 hybrid orbitals and one unhybridized p orbital Carbon uses two equivalent sp hybrid orbitals to form a triple bond with linear geometry, with two unhybridized p orbitals Atoms such as nitrogen and oxygen hybridize to form strong, oriented bonds The nitrogen atom in ammonia and the oxygen atom in water are sp3-hybridized

83 Learning Check: Based on the octet rule, which line-bond structures is/are correct? A and B A and C B and D Only B Only C

84 Solution: Based on the octet rule, which line-bond structures is/are correct? A and B A and C B and D Only B Only C

85 Which of these statements concerning p-orbitals is false?
Learning Check: Which of these statements concerning p-orbitals is false? They consist of two equivalent lobes. They are absent from the first shell of atomic orbitals. They can form π bonds. They only participate in bonding on carbon atoms. They can hold a maximum of two electrons.

86 Which of these statements concerning p-orbitals is false?
Solution: Which of these statements concerning p-orbitals is false? They consist of two equivalent lobes. They are absent from the first shell of atomic orbitals. They can form π bonds. They only participate in bonding on carbon atoms. They can hold a maximum of two electrons.

87 Learning Check: s p sp sp2 sp3
What is the hybridization of the oxygen-bonded carbon atom in the following molecule? s p sp sp2 sp3

88 Solution: What is the hybridization of the oxygen-bonded carbon atom in the following molecule? s p sp sp2 sp3

89 What is the hybridization of nitrogen atom in trimethylamine :N(CH3)3?
Learning Check: What is the hybridization of nitrogen atom in trimethylamine :N(CH3)3? sp sp2 sp3 dsp2 unhybridized

90 What is the hybridization of nitrogen atom in trimethylamine :N(CH3)3?
Solution: What is the hybridization of nitrogen atom in trimethylamine :N(CH3)3? sp sp2 sp3 dsp2 unhybridized

91 Learning Check: s p sp sp2 sp3
What type of orbital is not used in constructing the molecule shown below? s p sp sp2 sp3

92 Solution: What type of orbital is not used in constructing the molecule shown below? s p sp sp2 sp3

93 What is the molecular formula of the molecule shown below?
Learning Check: What is the molecular formula of the molecule shown below? C8H12O C8H11O C8H7O C7H7O C7H10O

94 What is the molecular formula of the molecule shown below?
Solution: What is the molecular formula of the molecule shown below? C8H12O C8H11O C8H7O C7H7O C7H10O

95 Electrons cannot occupy an antibonding molecular orbital.
Learning Check: Electrons cannot occupy an antibonding molecular orbital. True False

96 Electrons cannot occupy an antibonding molecular orbital.
Solution: Electrons cannot occupy an antibonding molecular orbital. True False

97 Learning Check: How many hydrogen atoms are present in the naturally-occurring terpene α-terpinene shown below? 14 15 16 17 18

98 Solution: How many hydrogen atoms are present in the naturally-occurring terpene α-terpinene shown below? 14 15 16 17 18

99 Learning Check: Select the best condensed structural formula for the following molecule: (CH3)2CHCH2COHOHCOH CH3CH3CHCH2C(OH)2CHO (CH3)2CHCH2C(OH)2CHO (CH3)2CHCH2C(OH)2COH CH3CHCH3CH2C(OH)2CHO

100 Solution: Select the best condensed structural formula for the following molecule: (CH3)2CHCH2COHOHCOH CH3CH3CHCH2C(OH)2CHO (CH3)2CHCH2C(OH)2CHO (CH3)2CHCH2C(OH)2COH CH3CHCH3CH2C(OH)2CHO

101 Learning Check: What is the correct order of carbon-carbon bond lengths in ethane, ethylene and acetylene? ethane < ethylene < acetylene ethane < acetylene < ethylene ethylene < ethane < acetylene ethylene < acetylene < ethane acetylene < ethane < ethylene acetylene < ethylene < ethane

102 Solution: What is the correct order of carbon-carbon bond lengths in ethane, ethylene and acetylene? ethane < ethylene < acetylene ethane < acetylene < ethylene ethylene < ethane < acetylene ethylene < acetylene < ethane acetylene < ethane < ethylene acetylene < ethylene < ethane

103 Learning Check: What kind of orbital and with how many nodal planes is shown in the picture below? σ antibonding, two nodal planes σ antibonding, one nodal plane π bonding, one nodal plane π antibonding, two nodal planes π antibonding, one nodal plane

104 Solution: What kind of orbital and with how many nodal planes is shown in the picture below? σ antibonding, two nodal planes σ antibonding, one nodal plane π bonding, one nodal plane π antibonding, two nodal planes π antibonding, one nodal plane

105 What is the best description of bond length?
Learning Check: What is the best description of bond length? a distance between nuclei that yields the best orbital overlap a distance between nuclei that yields the smallest nuclear-nuclear repulsion a distance between nuclei that yields the smallest electron-electron repulsion a distance between nuclei that yields the largest electron-nuclei attraction a distance between nuclei that is a compromise of all electrostatic interactions

106 What is the best description of bond length?
Solution: What is the best description of bond length? a distance between nuclei that yields the best orbital overlap a distance between nuclei that yields the smallest nuclear-nuclear repulsion a distance between nuclei that yields the smallest electron-electron repulsion a distance between nuclei that yields the largest electron-nuclei attraction a distance between nuclei that is a compromise of all electrostatic interactions

107 What is the O-C-O bond angle in potassium carbonate, K2CO3?
Learning Check: What is the O-C-O bond angle in potassium carbonate, K2CO3? 60° 90° 109.5° 120° 180°

108 What is the O-C-O bond angle in potassium carbonate, K2CO3?
Solution: What is the O-C-O bond angle in potassium carbonate, K2CO3? 60° 90° 109.5° 120° 180°

109 What is the relative orientation of two π orbitals in acetylene?
Learning Check: What is the relative orientation of two π orbitals in acetylene? 60° 90° 109.5° 120° 180°

110 What is the relative orientation of two π orbitals in acetylene?
Solution: What is the relative orientation of two π orbitals in acetylene? 60° 90° 109.5° 120° 180°

111 In relation to σ bonds, which statement about π bonds is correct?
Learning Check: In relation to σ bonds, which statement about π bonds is correct? π bonds are weaker and π electrons have higher energy π bonds are stronger and π electrons have higher energy π bonds are weaker and π electrons have lower energy π bonds are stronger and π electrons have lower energy π bonds and σ bonds have equal strengths, and their electrons have the same energies

112 In relation to σ bonds, which statement about π bonds is correct?
Solution: In relation to σ bonds, which statement about π bonds is correct? π bonds are weaker and π electrons have higher energy π bonds are stronger and π electrons have higher energy π bonds are weaker and π electrons have lower energy π bonds are stronger and π electrons have lower energy π bonds and σ bonds have equal strengths, and their electrons have the same energies

113 Which of the orbitals shown below can be occupied in a hydrogen atom?
Learning Check: Which of the orbitals shown below can be occupied in a hydrogen atom? A B C D Any one of these

114 Which of the orbitals shown below can be occupied in a hydrogen atom?
Solution: Which of the orbitals shown below can be occupied in a hydrogen atom? A B C D Any one of these

115 Which statement about the antibonding orbital in H2 is false?
Learning Check: Which statement about the antibonding orbital in H2 is false? It is higher in energy than the corresponding bonding orbital It is a molecular orbital It has a nodal plane (a node for short) It can hold up to two electrons The two lobes have different charges

116 Which statement about the antibonding orbital in H2 is false?
Solution: Which statement about the antibonding orbital in H2 is false? It is higher in energy than the corresponding bonding orbital It is a molecular orbital It has a nodal plane (a node for short) It can hold up to two electrons The two lobes have different charges

117 What is the hybridization of nitrogen in :NCCH3?
Learning Check: What is the hybridization of nitrogen in :NCCH3? sp sp2 sp3 sp3d nitrogen is not hybridized

118 What is the hybridization of nitrogen in :NCCH3?
Solution: What is the hybridization of nitrogen in :NCCH3? sp sp2 sp3 sp3d nitrogen is not hybridized

119 Each shell (floor of the Hotel)
Hotel Model Each shell (floor of the Hotel) Has subshells (s,p,d,f) f n = 4 (4th floor) d p 32 e- s n = 3 (3rd floor) d 18 e- p s n = 2 (2nd floor) p 8 e- s n = 1 (1st floor) s 2 e-


Download ppt "1. Structure and Bonding Why this chapter?"

Similar presentations


Ads by Google