1. Properties of Carbon Element

2. Properties of Carbon Element
We have learnt carbon element as the basis of Organic Chemistry.
Why does carbon can form so many different compounds?
Carbon is found the second period group IVA of the periodic table.

3. EXCEPTIONS OF CARBON COMPUDS WHICH ARE NOT ORGANIC
oxides of carbon (CO2, CO)
carbonates,bicarbonates(NaHCO3,CaCO3)
cyanides (NaCN, etc)

4. Properties of Carbon Element
The Lewis structure for carbon shows 4 unpaired valence electrons.
To fulfill the octet rule, a carbon atom needs 4 more electrons.
A carbon atom may form 4 covalent bonds and is capable of forming long chains with single, double or triple bonds between carbon atoms.
These chains may be continuous (straight) or branched.
The 2 ends of a chain can bond together to form a ring.
Carbon compounds are divided into classes based on their chemical similarity.

5. Hydrocarbons
Hydrocarbons are compounds containing hydrogen and carbon. Hydrocarbons may have different numbers of bonds between carbon atoms.
The four hydrocarbon classes are:
alkane (single bond),
alkene, (double bond),
alkyne (triple bond),
aromatic (benzene ring).
Alkanes contain only single C-C bonds. They contain as many hydrogen atoms as possible, and are said to be saturated.
Hydrocarbons containing double or triple bonds are unsaturated.
A homologous series is series of compounds that differ by a constant increment. Aromatic hydrocarbons include a benzene ring- 6 carbon atoms with all the bonds alternating between a single and a double bond.

6. Properties of Carbon Element
Carbon is unique
It has 6 electrons in its outer shell arranges 1s22s2sp2
It has room for 4 bonds to 4 other atoms.
Carbon-to-carbon bonds can be single (A),
double (B), or
triple (C).
Note that in each example, each carbon atom has four dashes, which represent four bonding pairs of electrons, satisfying the octet rule.

7. HYDROCARBONS

8. Hydrocarbons
Alkanes
C C
Alkenes
C C
Alkynes
C C
Aromatics

9. Properties of Carbon Element
A)The carbon atom forms bonds in a tetrahedral structure with a bond angle of 109.5O.
(B) Carbon-to-carbon bond angles are 109.5O, so a chain of carbon atoms makes a zigzag pattern.
(C) The unbranched chain of carbon atoms is usually simplified in a way that looks like a straight chain, but it is actually a zigzag, as shown in (B).

10. Properties of Carbon Element
Carbon-to-carbon chains can be
(A) straight,
(B) branched, or
(C) in a closed ring.
(Some carbon bonds are drawn longer, but are actually the same length.)

11. Why does carbon can form so many different compounds?
There are now more than ten million organic compounds known by chemists.
Many more undoubtedly exist in nature, and organic chemists are continually creating (synthesizing) new ones.
Carbon is the only element that can form so many different compounds because each carbon atom can form four chemical bonds to other atoms, and because the carbon atom is just the right, small size to fit in comfortably as parts of very large molecules.
Having the atomic number 6, every carbon atom has a total of six electrons.
Two are in a completed inner shell, while the other four are valence electrons—outer electrons that are available for forming bonds with other atoms.

12. Why does carbon can form so many different compounds?
The carbon atom's four valence electrons can be shared by other atoms that have electrons to share, thus forming covalent (shared-electron) bonds.
They can even be shared by other carbon atoms, which in turn can share electrons with other carbon atoms and so on, forming long strings of carbon atoms, bonded to each other like links in a chain.
Silicon (Si), another element in group 4A of the periodic table, also has four valence electrons and can make large molecules called silicones, but its atoms are too large to fit together into as great a variety of molecules as carbon atoms can.

13. Why does carbon can form so many different compounds?
Carbon's ability to form long carbon-to-carbon chains is the first of five reasons that there can be so many different carbon compounds; a molecule that differs by even one atom is, of course, a molecule of a different compound.
The second reason for carbon's astounding compound-forming ability is that carbon atoms can bind to each other not only in straight chains, but in complex branchings, like the branches of a tree.
They can even join "head-to-tail" to make rings of carbon atoms.
There is practically no limit to the number or complexity of the branches or the number of rings that can be attached to them, and hence no limit to the number of different molecules that can be formed.
The third reason is that carbon atoms can share not only a single electron with another atom to form a single bond, but it can also share two or three electrons, forming a double or triple bond.
This makes for a huge number of possible bond combinations at different places, making a huge number of different possible molecules.
And a molecule that differs by even one atom or one bond position is a molecule of a different compound.

14. Why does carbon can form so many different compounds
The fourth reason is that the same collection of atoms and bonds, but in a different geometrical arrangement within the molecule, makes a molecule with a different shape and hence different properties.
These different molecules are called isomers.
The fifth reason is that all of the electrons that are not being used to bond carbon atoms together into chains and rings can be used to form bonds with atoms of several other elements.
The most common other element is hydrogen, which makes the family of compounds known as hydrocarbons.
But nitrogen, oxygen, phosphorus, sulfur, halogens, and several other kinds of atoms can also be attached as part of an organic molecule.
There is a huge number of ways in which they can be attached to the carbon-atom branches, and each variation makes a molecule of a different compound.

15. The Greater Stability of C-C Bonds
Since the average bond dissociation energy of C-C is greater than the average bond energies between different atoms.
Thus the energy released when carbon atom bonds to another carbon atom is greater than the energy released when the other atoms like B,N,O,Si,P and S bonds to each other.
Thus C-C bond is more stable than the others like B-B,N-N,
O-O,Si-Si,P-P and S-S.

16. Ability to Form Chains Between Their Atoms
The atoms closer to C in the periodic table are B,N,O,Si,P and S.
The ability of these atoms to bond each other to form chains is lower than C.
For examle Si can produce chains made of at most 11 atoms of it and N at most three atoms it.
Although the ability to form chains between their atoms for P and S is greater than Si and N but it is very much smaller compared to C.

17. Ability to Form Chains Between Their Atoms
The greater ability of carbon to form chains compared to atoms closer to it in the periodic table can be explained by two reasons:
The average bond dissociation energies of them is lower than that of carbon.
The electronegativity values B,Si and P lower than that of C.atoms.Thus the attraction forces between these atoms are smaller than that of carbon.This is also true when these atoms are bonded to the other atoms like hydrogen or halogens.

18. Electronegativity Electronegativity: Pauling scale
a measure of an atom’s attraction for the electrons it shares with another atom in a chemical bond
Pauling scale
generally increases left to right in a row
generally increases bottom to top in a column

19. Greater Bonding Capacity of C compared to N and O
The electronegativity values of N and O are greater than that of C. But their bonding capacities are smaller than that of C since they have lower number of unpaired electrons.
Lewis Dot Diagrams of Selected Elements

20. Summary…
Compared to C atom B,Si,P,N and O atoms can not be expected to form greater number of compounds and unbrached and branched chains and cyclic compounds.
Carbon compounds are more stable than Si4,P4,O3,S8 and B4 molecules.

21. Electron Configuration of Elements

22. Lewis Dot Structures… Gilbert N. Lewis Valence shell:
the outermost occupied electron shell of an atom
Valence electrons:
electrons in the valence shell of an atom; these electrons are used to form chemical bonds and in chemical reactions
Lewis dot structure:
the symbol of an element represents the nucleus and all inner shell electrons
dots represent valence electrons

23. Lewis Dot Structures
Table 1.4 Lewis Dot Structures for Elements 1-18

24. Lewis Model of Bonding…
Atoms bond together so that each atom acquires an electron configuration the same as that of the noble gas nearest it in atomic number
an atom that gains electrons becomes an anion
an atom that loses electrons becomes a cation
the attraction of anions and cations leads to the formation of ionic solids
an atom may share electrons with one or more atoms to complete its valence shell; a chemical bond formed by sharing electrons is called a covalent bond
bonds may be partially ionic or partially covalent; these bonds are called polar covalent bonds

25. Covalent Bonds! The simplest covalent bond is that in H2
the single electrons from each atom combine to form an electron pair
the shared pair functions in two ways
simultaneously; it is shared by the two atoms and fills the valence shell of each atom
The number of shared pairs
one shared pair forms a single bond
two shared pairs form a double bond
three shared pairs form a triple bond

26. Hydrogen Molecule Formation
Potential Energy Diagram - Attraction vs. Repulsion
Energy (KJ/mol)
balanced attraction
& repulsion
no interaction
increased
attraction
The change in potential energy during the formation of hydrogen molecule. The minimum energy, at 0.74 angstrom, represents the equilibrium bond distance.
The energy at this point, -426 kJ/mol, corresponds to the energy change for formation of the H – H bond.
Potential energy is based on the position of an object. Low potential energy = high stability.
increased
repulsion
- 436
0.74 A
H – H distance
(internuclear distance)
Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 318

27. Lewis Structures! To write a Lewis structure
determine the number of valence electrons
determine the arrangement of atoms
connect the atoms by single bonds
arrange the remaining electrons so that each atom has a complete valence shell
show a bonding pair of electrons as a single line
show a nonbonding pair of electrons as a pair of dots
in a single bond atoms share one pair of electrons, in a double bond they share two pairs of electrons, and in a triple bond they share three pairs of electrons

28. Table of Lewis Structures!
In neutral molecules
hydrogen has one bond
carbon has 4 bonds and no lone pairs
nitrogen has 3 bonds and 1 lone pair
oxygen has 2 bonds and 2 lone pairs
halogens have 1 bond and 3 lone pairs

29. Resonance!
In chemistry, resonance is a way of describing delocalized electrons within certain molecules or polyatomic ions where the bonding cannot be expressed by one single Lewis formula.
A molecule or ion with such delocalized electrons is represented by several structures called resonance structures.

30. Tautomerization
Tautomerization usually involves the movement of a hydrogen atom between a different location on the molecule, resulting in two or more molecular structures.
These structures are called tautomers, which exist in dynamic equilibrium with each other.
Enol form Keto form

31. Molecular Geometry and Bonding Theories

32. CH4 C H C H molecular molecular structural formula shape formula
ball-and-stick
model
tetrahedral
shape of
methane
tetrahedron

33. Methane & Carbon Tetrachloride
molecular
formula
structural
formula
molecular
shape
ball-and-stick
model C H H
109.5o C
CH4
The molecular geometry is predicted by first writing the Lewis structure, then using the VSEPR model to determine the electron-domain geometry, and finally focusing on the atoms themselves to describe the molecular structure.
space-filling model C Cl
CCl4

34. Molecular Geometry Trigonal planar Linear Tetrahedral Bent
Trigonal pyramidal
H2O CH4 AsCl3 AsF5 BeH2 BF3 CO2

35. N H .. .. C H O .. H H CH4, methane NH3, ammonia H2O, water 107o

36. Molecular Shapes Three atoms (AB2) Four atoms (AB3) B A Linear (180o)
Bent
Trigonal planar (120o)
Trigonal pyramidal B A
linear
trigonal planar
Five atoms (AB4) B A
Tetrahedral (109.47o)
tetrahedral
Bailar, Moeller, Kleinberg, Guss, Castellion, Metz, Chemistry, 1984, page 313.

37. Bonding and Shape of Molecules
Number
of Bonds
Number of
Unshared Pairs
Covalent
Structure
Shape
Examples 2 3 4 1 2
-Be-
Linear
Trigonal planar
Tetrahedral
Pyramidal
Bent
BeCl2
BF3
CH4, SiCl4
NH3, PCl3
H2O, H2S, SCl2 B C N : O :

38. Molecular Shapes AB2 Linear AB3 Trigonal planar AB2E Angular or Bent
Tetrahedral
AB3E
Trigonal
pyramidal

39. Molecular Polarity Molecular Structure
Courtesy Christy Johannesson

40. + - Dipole Moment H Cl Direction of the polar bond in a molecule.
Arrow points toward the more electronegative atom.
H Cl + -

41. Determining Molecular Polarity
Depends on:
dipole moments
molecular shape
Courtesy Christy Johannesson

42. Determining Molecular Polarity
Nonpolar Molecules
Dipole moments are symmetrical and cancel out.
BF3 F B
Courtesy Christy Johannesson

43. Determining Molecular Polarity
Polar Molecules
Dipole moments are asymmetrical and don’t cancel .
H2O H O
net
dipole
moment
Courtesy Christy Johannesson

44. Determining Molecular Polarity
Therefore, polar molecules have...
asymmetrical shape (lone pairs) or
asymmetrical atoms
CHCl3 H Cl
net
dipole
moment
Courtesy Christy Johannesson

45. Dipole Moment Nonpolar m = Q r Polar C O O O H H .. Bond dipoles
In H2O the bond dipoles are also equal in
magnitude but do not exactly oppose each
other. The molecule has a nonzero overall
dipole moment. C O O ..
Overall dipole moment = 0 O
Bond dipoles
Nonpolar H H
The overall dipole moment of a molecule
is the sum of its bond dipoles. In CO2 the
bond dipoles are equal in magnitude but
exactly opposite each other. The overall
dipole moment is zero.
Overall dipole moment
m = Q r
Dipole moment, m
Coulomb’s law
Polar
Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 315

46. Polar and Nonpolar Molecules.. .. .. F O N H Cl H H H B H H F F
Polar
Polar
Nonpolar
Polar Cl Cl
Students often confuse electron-domain geometry with molecular geometry.
You must stress that the molecular geometry is a consequence of the electron domain geometry.
The best arrangement of a given number of electron domains is the one that minimizes the repulsions among them. Cl H C C Cl H Cl H
Nonpolar
Polar
A molecule has a zero dipole moment because their dipoles cancel one another.

47. Debye, Physicist Peter Debye’nin soyadindan geliyor.

48. Formation of BeH2 using pure s and p orbitals
Be = 1s22s2 H Be
BeH2 H s p
No overlap = no bond!
atomic orbitals
atomic orbitals
The formation of BeH2 using hybridized orbitals Be H s p
atomic orbitals Be H
hybrid orbitals Be s p Be
BeH2 sp p
All hybridized bonds have equal strength and have orbitals with identical energies.

49. Hybridization - The Blending of Orbitals+ =
Poodle +
Cocker Spaniel =
Cockapoo + =
s orbital +
p orbital =
sp orbital

50. sp hybrid orbitals shown together
Ground-state Be atom 1s 2s 2p
Be atom with one electron “promoted” sp
hybrid orbitals
Energy 1s sp 2p
Be atom of BeH2 orbital diagram px py pz
n = 1
n = 2 s
two sp hybrid orbitals
s orbital
p orbital
hybridize H Be
sp hybrid orbitals shown together
(large lobes only)

51. sp Animation

52. sp2 hybrid orbitals shown together
Ground-state B atom 2s 2p 2s 2p
B atom with one electron “promoted”
sp2
hybrid orbitals
Energy
sp2 2p px py pz s
B atom of BH3 orbital diagram
p orbitals H B
three sps hybrid orbitals
sp2 hybrid orbitals shown together
(large lobes only)
hybridize
s orbital

53. Sp2 Animation

54. sp3 Hybrid Orbitals Carbon 1s22s22p2 Carbon could only make two bonds
if no hybridization occurs. However,
carbon can make four equivalent bonds. B A
sp3
hybrid orbitals
Energy px py pz
sp3 s
C atom of CH4 orbital diagram
Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 321

55. Sp3 Animation

56. Multiple Bonds C2H4, ethene 2s 2p 2s 2p sp2 2p H C
promote
hybridize
2s p s p sp p
C2H4, ethene C H
one s bond and one p bond H C s H C
Two lobes of
one p bond
Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page

57. Multiple Bonds C2H4, ethene (ethylene) one s bond and one p bond C
promote
hybridize
2s p s p sp p p
C2H4, ethene (ethylene) C H H
sp2
one s bond and one p bond H C s H C
Two lobes of
one p bond
Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page

58. Sigma and pi Bonds Animation

59. p bond
Internuclear axis p p

60. p bond Two atomic p orbitals Pi (p) Bond 3D view of Pi (p) Bond
(pi) Bond – overlap of two p orbitals oriented perpendicular
to the line connecting the nuclei.

61. Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 326

62. Orbital Picture of Ethylene
H s
H s 1 px px
sp2
sp2   
C C 
sp2
sp2
sp2
H s
sp2
H s  

63. Ethylene Animation

64. Bonding in Formaldehyde

65. s bonds in Benzene H H C C H C C H C C H H C6H6 = benzene
Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 329

66. 2p atomic orbitals in Benzene
Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 329

67. bonds and p bonds H H C C H C C H C C H H
Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 329

68. s bonds in Benzene H C H C H C H C H C H C H C H C
Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 329

69. s bonds H C H C H C H C H C H C H C H C
Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 329

70. Ethyne sp Hybridization

71. Ethyne sp Hybridization

72. Orbital Picture of Acetylene1 px py px py   2
H s sp sp sp sp
H s
C C 

73. Acetylene Animation

74. Acetic Acid, CH3COOH H O H C C O H H Number of electron domains 4 3 4
Trigonal
planar
Electron-domain geometry
Tetrahedral
Tetrahedral
Predicted bond angles
109.5o
120o
109.5o
Hybridization of central atom
sp3
sp2
sp3
Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 314

75. VSEPR Theory Valence Shell Electron Pair Repulsion Theory
Electron pairs orient themselves in order to minimize repulsive forces.
Courtesy Christy Johannesson

76. The Shapes of Some Simple ABn Molecules.. .. ..
H2O F B O S F N H B
Linear
Bent
Trigonal
planar
Trigonal
pyramidal
SF6
Students often confuse electron-domain geometry with molecular geometry.
You must stress that the molecular geometry is a consequence of the electron domain geometry.
The best arrangement of a given number of electron domains is the one that minimizes the repulsions among them.
Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 305

77. Lone pairs repel more strongly than bonding pairs!!!
VSEPR Theory
Types of e- Pairs
Bonding pairs - form bonds
Lone pairs - nonbonding electrons
Lone pairs repel more strongly than bonding pairs!!!
Courtesy Christy Johannesson

78. VSEPR Theory Lone pairs reduce the bond angle between atoms.
Courtesy Christy Johannesson

79. Determining Molecular Shape
Draw the Lewis Diagram.
Tally up e- pairs on central atom.
double/triple bonds = ONE pair
Shape is determined by the # of bonding pairs and lone pairs.
Know the 5 common shapes
& their bond angles!
Courtesy Christy Johannesson

80. Common Molecular Shapes
2 total
2 bond
0 lone B A
BeH2
LINEAR
180°
Courtesy Christy Johannesson

81. Common Molecular Shapes
3 total
3 bond
0 lone B A
BF3
TRIGONAL PLANAR
120°
Courtesy Christy Johannesson

82. Common Molecular Shapes
4 total
4 bond
0 lone B A
CH4
TETRAHEDRAL
109.5°
Courtesy Christy Johannesson

83. Common Molecular Shapes
4 total
3 bond
1 lone
NH3
TRIGONAL PYRAMIDAL
107°
Courtesy Christy Johannesson

84. Common Molecular Shapes
4 total
2 bond
2 lone
H2O
BENT
104.5°
Courtesy Christy Johannesson

85. O C O 180° Examples LINEAR 2 total 2 bond 0 lone CO2
Courtesy Christy Johannesson

86. F P F F 107° Examples TRIGONAL PYRAMIDAL 4 total 3 bond 1 lone PF3
Courtesy Christy Johannesson

87. No Loners Animation

88. Loners Animation

90. References