1. Chapter 1 Structure and Bonding

2. Learning Objectives Atomic structure: The nucleus
Atomic structure: Orbitals
Atomic structure: Electron configurations
Development of chemical bonding theory
Describing chemical bonds: Valence bond theory
sp3 hybrid orbitals and the structure of methane
sp3 hybrid orbitals and the structure of ethane
sp2 hybrid orbitals and the structure of ethylene
sp hybrid orbitals and the structure of acetylene
Hybridization in nitrogen and oxygen
Describing chemical bonds: Molecular orbital theory
Drawing chemical structures

3. What is Organic Chemistry?
The branch of chemistry, originally limited to substances found only in living organisms, but now dealing with the compounds of carbon, hence the name organic.
Living things are made of organic chemicals
Proteins that make up muscles and hair
DNA, controls genetic make-up
Foods, medicines
Examine structures to the right
Vioxx is used in the treatment of pain; osteoarthritis; rheumatoid arthritis; migraine
Penicillins are used in the treatment or prevention of many different bacterial infections.

4. Origins of Organic Chemistry
Organic chemistry is study of carbon compounds. Why is it so special?
90% of more than 50 million chemical compounds contain carbon.
Examination of element carbon in periodic table answers some of these questions.
Carbon is group 4A element, it can share 4 valence electrons and form 4 covalent bonds. Only element that has the ability to form immense diversity of compounds
To under stand what are valence electrons, we need to learn about the atomic structure of atoms/elements
To under stand why carbon has 4 valence electrons, we need to learn about the atomic structure of elements

5. 1.1 Atomic Structure Structure of an atom
Positively charged nucleus (very dense, contains protons and neutrons) and very small in size (diameter about10-15 m)
Negatively charged electrons are in a cloud (diameter about m) around nucleus
Diameter of atom is about 2  m (200 picometers (pm)) or 2 Å [the unit angstrom (Å) is m = 100 pm]
1/10th of nanometer is one A or 10 A = 1nanometer.

6. Atomic Number and Atomic Mass
The atomic number (Z) is the number of protons in the atom's nucleus
The mass number (A) is the number of protons plus neutrons
All the atoms of a given element have the same atomic number
Isotopes are atoms of the same element that have same number of protons but different numbers of neutrons and therefore different mass numbers
The atomic mass (atomic weight) of an element is the weighted average mass in atomic mass units (amu) of an element’s naturally occurring isotopes
Example Carbon atomic mass = (weighted average of C-12 and C-13 isotopes, 98.89% x 12 amu % x amu)
Bromine = (weighted average of Br-79 and Br-81)
1 amu = × 10-24 grams; Avogadro number 6.022×1023, 1g hydrogen has Avagadro number of atoms, 12g of carbon has Avagadro number of atoms; 79.9g of bromine has Avagadro #.
Carbon atomic mass = amu
(weighted average of C-12 and C-13 isotopes, 98.89% x 12 amu % x amu)
Bromine atomic mass =
(weighted average of Br-79 and Br-81 isotopes, 50.69% x 79 amu % x 81 amu)

7. 1.2 Atomic Structure: Orbitals Shapes of Atomic Orbitals for Electrons
Four different kinds of orbitals
Denoted s, p, d, and f
s and p orbitals most important in organic and biological chemistry
s orbitals: spherical, nucleus at center
p orbitals: dumbbell-shaped, nucleus at middle
d orbitals: elongated dumbbell-shaped, nucleus at center
Four are clover leaf shape and one is different shape, dumbbell with a ring shape. The f orbitals have more complex shape and will not be covered here.

8. Orbitals and Shells
Orbitals are grouped in shells of increasing size and energy (shell, n = 1,2,3 …)
Different shells contain different numbers and kinds of orbitals (sub shells s, p, d, f)
Each orbital can be occupied by two electrons (with opposite spin)
First shell contains one s orbital, denoted 1s, holds only 2 electrons- (1 sub shell)
Second shell contains one s orbital (2s) and three p orbitals (2p), total 8 electrons- (2 sub shells)
Third shell contains one s orbital (3s), three p orbitals (3p), and five d orbitals (3d), total18 electrons- (3 sub shells)
4th shell contains one s orbital (4s), three p orbitals (4p), five d orbitals (4d) and seven f orbitals (4f) 32 electrons

9. 1.3 Atomic Structure: Electron Configurations
Ground-state electron configuration (lowest energy arrangement) Rules:
1. Lowest-energy orbitals fill first: 1s  2s  2p  3s  3p  4s  3d
(Aufbau principle)
2. 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)
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)

10. Order of orbitals (filling) in multi-electron atom
Look for n+l value and if it is same for two, the one with lower n value comes first (lower energy)
1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s

11. Worked Example
Give the ground-state electron configuration (EC) for sulfur
Solution:
Atomic number of sulfur is 16
Number of electrons = 16
Orbital diagram
In a more concise way it can be written as
1s2 2s2 2p6 3s2 3p4 (EC) or
[Ne] 3s2 3p4 (condensed or core EC)
[Ne]
1s22s22p63s23p4 The large numbers represent the energy level.   
The letters represent the sublevel.   
The superscript numbers indicate the number of electrons in the sublevel.

12. Worked Example
How many electrons does sodium have in its outermost electron shell?
Solution:
Elements of the periodic table are organized into groups based on the number of outer-shell electrons each element has
Using the periodic table we locate the group of the element, sodium
Sodium - Group 1A
Has one electron in its outermost shell
1s2 2s2 2p6 3s2 or [Ne] 3s1
Atomic structure: Electron configurations

14. 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

15. Valence shell: Atom’s outermost shell
Atoms form bonds because the resulting compound is more stable than the separate atoms
Valence shell: Atom’s outermost shell
Impart special stability to the noble gas elements
Ionic bonds - Ions held together by a electrostatic attraction (NaCl)
Formed as a result of electron transfers
Covalent bond: Formed by sharing of electrons (H2 , CH4)
Organic compounds have covalent bonds from sharing electrons
Development of chemical bonding theory

16. Molecule: Neutral collection of atoms held together by covalent bonds
Lewis structures (electron dot) show valence electrons of an atom as dots
Hydrogen has one dot
Carbon has four dots
Kekule structures (line-bond structures) have a line drawn between two atoms indicating a 2 electron covalent bond.
Stable molecule results at completed shell, octet (eight dots) for main-group atoms (two for hydrogen)
Molecule: Neutral collection of atoms held together by covalent bonds
Groups 1, 2 (1A, 2A) and 13 to 18 (3A to 8A) constitute the main group. Sometimes group 12 is also included in the main group.

17. Atoms with one, two, or three valence electrons form one, two, or three bonds.
Atoms with four or more valence electrons form as many bonds as they need electrons to fill the s and p levels of their valence shells to reach a stable octet.
Carbon has four valence electrons (2s2 2p2), forming four bonds (CH4).
Nitrogen has five valence electrons (2s2 2p3) but forms only three bonds (NH3).
Oxygen has six valence electrons (2s2 2p4) but forms two bonds (H2O)

19. 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 or lone pair

20. Worked Example
Draw a molecule of chloroform, CHCl3, using solid, wedged, and dashed lines to show its tetrahedral geometry
Solution:
Development of chemical bonding theory

21. 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, called sigma (s) bond
In Valence-Bond Theory, electrons of two atoms begin to occupy the same space.
This is called “overlap” of orbitals.
The sharing of space between two electrons of opposite spin results in a covalent bond.

22. 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.

23. Bond Length
Distance between nuclei of two atoms that leads to maximum stability
If too close, they repel because both are positively charged
If too far apart, bonding is weak
1000 pm = 1 nm 0r 10 A

24. 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)

25. 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 438 kJ/mol and length of 109 pm
Bond angle: each H–C–H is 109.5°, the tetrahedral angle.
All the bonds are sigma bonds

26. sp3 Orbitals and the Structure of Ethane
Two C’s bond to each other by s overlap of an sp3 orbital from each
Remaining three sp3 orbitals on each C overlap with H 1s orbitals to form six C–H bonds
C–H bond strength in ethane 421 kJ/mol
C–C bond is 154 pm long and strength is 377 kJ/mol
All bond angles of ethane are close to tetrahedral (109.5°)

27. sp2 Orbitals and the Structure of Ethylene
sp2 hybrid orbitals: Derived by combination of an s atomic orbital with two p atomic orbitals
sp2 orbitals are in a plane with an angle of 120°from each other
Remaining p orbital is perpendicular to the plane
sp2 hybrid orbitals and the structure of ethylene

28. sp2 Orbitals and the Structure of Ethylene
Two sp2 hybridized orbitals on each carbon overlap to form a σ bond
p orbitals on each C interact by sideways overlap to form a pi () bond
sp2–sp2 σ bond and 2p–2p  bond result in sharing four electrons and formation of C=C double bond
Electrons in the σ bond are centered between nuclei
Electrons in the  bond occupy regions on either side of a line between nuclei
sp2 hybrid orbitals and the structure of ethylene

29. Structure of Ethylene H atoms form s bonds with four sp2 orbitals
H–C–H and H–C–C bond angles are of about 120°
C=C double bond in ethylene is shorter and stronger than single bond (C-C) in ethane
Ethylene C=C bond length is 134 pm (C–C 154 pm)
Vinylic C-H bond energy 464 kJ/mole C=C 620 kJ/mole
Caetylinic C-H 556 kJ/mol C=C 836 kJ/mole

30. Worked Example
Draw electron-dot and line-bond structures of formaldehyde
Indicate the hybridization of the carbon orbitals
Solution:
Two hydrogens, one carbon, and one oxygen can combine in one way
The orbitals are sp2-hybridized
sp2 hybrid orbitals and the structure of ethylene

31. sp Orbitals and the Structure of Acetylene
Carbon can form 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
sp hybrid orbitals and the structure of acetylene

32. Figure 1.15 - sp Hybridization
sp hybrid orbitals and the structure of acetylene

33. Orbitals of Acetylene
Two sp hybrid orbitals from each C form sp–sp σ bond
pz orbitals from each C form a pz–pz  bond by sideways overlap
py orbitals overlap sideways to form py–py  bond
sp hybrid orbitals and the structure of acetylene

34. 1.9 sp Orbitals and the Structure of Acetylene
C=C a triple bond sharing six electrons
Carbon: sp hybrids
sp orbitals are linear, 180 apart on x-axis
Two p orbitals are perpendicular

35. C=C
1 kcal = 4.18 kJ 1pm = m

36. Worked Example Draw a line-bond structure for propyne, CH3CºCH
Indicate the hybridization of the orbitals on each carbon
Predict a value for each bond angle
Solution:
sp hybrid orbitals and the structure of acetylene

37. Worked Example C3-H bonds are σ bonds C1-H bond is a σ bond
Overlap of an sp3 orbital of carbon-3 with s orbital of hydrogens
C1-H bond is a σ bond
Overlap of an sp orbital of carbon-1 with an s orbital of hydrogen
C2-C3 bond is a σ bond
Overlap of an sp orbital of carbon-2 with an sp3 orbital of carbon 3
Three C1-C2 bonds (one σ and 2 )
Bond angle
Between three carbon atoms and H–C1≡C2 is 180°)
Between hydrogen and the sp3-hybridized carbon is 109°
sp hybrid orbitals and the structure of acetylene

38. 1.10 Hybridization of Nitrogen and Oxygen
Elements other than C can have hybridized orbitals
N’s orbitals (sppp) hybridize to form four sp3 orbitals
One sp3 orbital is occupied by two nonbonding electrons (lone pair electrons), and three sp3 orbitals have one electron each, forming bonds to H and CH3.

39. Hybridization of Nitrogen and Oxygen
Oxygen atom in methanol can be described as sp3-hybridized
C–O–H bond angle in methanol is 108.5
Two sp3 hybrid orbitals on oxygen are occupied by four nonbonding electrons (two lone pairs)
Hybridization of nitrogen, oxygen, phosphorus, and sulfur

40. Worked Example
Identify all nonbonding lone pairs of electrons in the oxygen atom in dimethyl ether, CH3–O–CH3
What is its expected geometry
Solution:
The sp3-hybridized oxygen atom has tetrahedral geometry
Hybridization of nitrogen, oxygen, phosphorus, and sulfur

41. How to predict the hybridization of the central atom?
You can find the hybridization of the atom by finding its steric number:
Steric Number = # of atoms bonded + # lone pairs on that atom.
Steric Number
Geometry
Hybridization
Examples 2
linear sp
BeCl2, C2H2
BF3, C2H4 3
trigonal planar
sp2 4
tetrahedral
sp3
CH4, NH3, H2O
Phosphorus atomic # 15 = 1s22s22p63s23p3 Sulfur atomic # 16 = 1s22s22p63s23p4
PCl5 5
trigonal bipyramidal
sp3d 6
octahedral
sp3d2
SF6
If steric number comes to be 4, there is sp3 hybridization in the atom.
If steric number comes to be 3, there is sp2hybridization in the atom
If steric number comes to be 2, there is sp hybridization in the atom

42. 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
Molecular Orbitals differ from atomic orbitals because they represent the entire molecule, not a single atom.
Whenever two atomic orbitals overlap, two molecular orbitals are formed: one bonding, one antibonding.
Bonding MO is lower in energy
Antibonding MO is higher energy

43. MO Diagram
An energy-level diagram, or MO diagram shows how orbitals from atoms combine to give the molecule.
In H2 the two electrons go into the bonding molecular orbital (lower in energy).
Bond order = ½(# of bonding e-s – # of antibonding e-s)
= ½(2 – 0) = 1 bond
Larger bond order means greater bond strength.

44. 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

45. 1.12 Drawing Structures
Drawing every bond in organic molecule can become tedious.
Condensed structures: Condensed structures don’t have C-H or C-C single bonds shown. They are understood.
e.g.

46. Rules for Drawing Skeletal Structures
Carbon atoms aren’t usually shown
Carbon atom is assumed to be at each intersection of two lines (bonds) and at the end of each line
Hydrogen atoms bonded to carbon aren’t shown
Atoms other than carbon and hydrogen are shown
Drawing chemical structures

48. Worked Example
How many hydrogens are bonded to each carbon in the following compound
Give the molecular formula of the substance
Drawing chemical structures

49. Worked Example
Solution:
Drawing chemical structures

50. Summary Organic chemistry – chemistry of carbon compounds
Atom: positively charged nucleus surrounded by negatively charged electrons
Electronic structure of an atom
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

51. Summary (cont’d)
Sigma (s) bonds - Circular cross-section and are formed by head-on interaction of orbitals
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