1. CHEM
160 √

2. Text: Organic Chemistry, 6th ed. Brown, Foote, Iverson
Also: Study Guide and Problems & models are strongly recommended!
Consider other aids: Organic As a Second Language
OWL
Different organic texts
Others?
Mid Exam: , 2016
lab-135

3. Chapter 1 COVALENT BONDING & SHAPES OF MOLECULES
1.1 Electronic Structure of Atoms
1.2 Lewis Model (Octet Rule, Formal Charge)
1.3 Functional Groups
1.4 Bonding Angles and Shapes of Molecules
1.5 Polar and Nonpolar Molecules
1.6 Resonance
1.7 Quantum or Wave Mechanics
1.8 Molecular Orbital & Valence Bond Theory, Covalent Bonds, Hybridization (sp3, sp2, sp)
SUMMARY and OVERVIEW 1-11
Aug 22, 2002 (1)
Jan 21(14) lecture 1 1

4. forms strong double and triple bonds (to C or other atoms)
Organic Chemistry: the study of compounds of carbon
Li Be B C N O F Ne
WHY is CARBON SPECIAL?
CARBON - is small, intermediate electronegativity
forms strong bonds with itself/other atoms
C-C kcal/mole
2X as strong as: N-N, O-O, Si-Si
forms strong double and triple bonds (to C or other atoms)
Over 10 million structures identified; ~1000 new/day!

5. 4/18/2018
Dr Seemal Jelani

6. C Cu Carbon Copper Symbols All elements have their own unique symbol.
It can consist of a single capital letter, or a capital letter and one or two lower case letters. C
Carbon Cu
Copper
4/18/2018
Dr Seemal Jelani

7. Common Elements and Symbols
4/18/2018
Dr Seemal Jelani

8. motion of e’s particle & wave like
Recall the Structure of an ATOM:
nucleus
(neutrons + protons(+))
electrons(-)
1833 to 1
mass:
quantum mechanics:
motion of e’s particle & wave like
electrons confined to regions of space: shells (principle energy levels)

9. Valence Electrons
The number of valence electrons an atom has may also appear in a square.
Valence electrons are the electrons in the outer energy level of an atom.
These are the electrons that are transferred or shared when atoms bond together.
4/18/2018
Dr Seemal Jelani

10. Electronic Structure of Atoms
each shell can hold 2n2 electrons

11. Ground state electronic configuration (atoms or molecules)
Aufbau Principle:
fill orbitals lowest to highest energy
Pauli Exclusion Principle:
2 electrons per orbital, spins paired
Hund’s Rule:
degenerate orbitals, 1 electron in each
then create a pair

12. writing electronic configuration
Valence
Shell
closed
shell
Hund’s
Rule 3
(s)
next level
Pauli
&
Lewis
(p) E 2
(s) 1
(s)
Li Be B C N O F Ne
H He
writing electronic configuration

13. Electronic Configuration
No. of e's
electronic configuration
1 H s1
2 He s2
3 Li s s1
4 Be …….
5 B ………
6 C s2 2s2 2p2
7 N etc.
What is the electronic configuration of OXYGEN?
1s2, 2s2, 2p4

14. Lewis Structures (Gilbert N. Lewis)
Valence shell:
the outermost electron shell
Valence electrons:
electrons in valence shell
electrons used to bonds
Lewis structure:
atom symbol = nucleus + inner e-
dots represent valence electrons

15. Lewis Structures
Table 1.4 Lewis Structures

16. Bonding extremes - IONIC or COVALENT
ionic bonds ?
- loss or gain of valence electrons
covalent bonds ?
-share electrons to fill shells or

17. Electronegativity
Electronegativity: a measure of an atom’s attraction for the electrons it shares with another atom
Pauling scale
increases left to right in a row
increases bottom to top in a column
Jan 12, ‘04 start here (~11 slides per lecture)

18. Electronegativity Table 1.6 Classification of Bonds
electronegativity difference bond type
less than 0.5
covalent
H3C-H + -
0.5 to 1.9
polar covalent
Na+ -Cl
greater than 1.9
ionic

19. COVALENT BONDING & SHAPES OF MOLECULES
Chapter 1
COVALENT BONDING & SHAPES OF MOLECULES
1.1 Electronic Structure of Atoms
1.2 Lewis Model (Octet Rule, Formal Charge)
1.3 Functional Groups
1.4 Bonding Angles and Shapes of Molecules
1.5 Polar and Nonpolar Molecules
1.6 Resonance
1.6 Quantum or Wave Mechanics
1.7 Molecular Orbital & Valence Bond Theory, Covalent Bonds, Hybridization (sp3, sp2, sp)
SUMMARY and OVERVIEW 1-11
Aug 22, 2002 (1)
Jan 21(14) lecture 1 1

20. Lewis (electron dot) structures covalent molecules and ions
When discussing the physical and chemical properties of an element, chemists often focus on the electrons in the outermost shell of the atom because these electrons are involved in the formation of chemical bonds and in chemical reactions.
Carbon, for example, with the ground-state electron
configuration 1s22s 22p 2 has four outer-shell electrons. Outer-shell electrons are called valence electrons
The energy level in which they are found is called the valence shell.

21. To show the outermost electrons of an atom, we commonly use a representation called a Lewis dot structure, after the American chemist Gilbert N. Lewis (1875–1946) who devised this notation.
A Lewis dot structure shows the symbol of the element surrounded by a number of dots equal to the number of electrons in the outer shell of an atom of that element. In Lewis dot structures, the atomic symbol represents the core; that is, the nucleus and all inner shell electrons.
Duplet
Octet
Show how the loss of an electron from a sodium atom leads to a stable octet.

22. Shells
The probability of finding an electron in various regions of space relative to nucleus
The energy of electrons in the shells is quantized.
(specific values of energy are possible)
The shells only occur at quantized energies in which three important effects balance each other
Electrostatic attraction (electrons have to the nucleus and that draws them toward the nucleus)

23. Electrostatic repulsion between electrons
Wavelike nature of an electron that prefers to be delocalized,
Thereby spreading the electron density away from the nucleus.
DELOCALIZATION
Term that describes the spreading of electron density over a larger volume of space

24. Distribution of electrons in shells
ES are identified by the Principle quantum numbers
1,2,3……………
Each shell can hold upto 2n2 electrons where n is the number of the shell.
Calculate the electrons in the shells
First shell x(1)2 =2
2x(2)2 = 2x(2x2)= 2x4=8
Electrons with lowest energy strongly hold by positively charged protons

26. Formation of chemical bonds
Anion
Cation
Ionic interaction
Covalent bonds
Electronegativity (A measure of the force of an atom’s
attraction for electrons.)
Electronegativity and Chemical Bonds (Polar and non-polar)

27. Formation of Ions
Ions are formed by the transfer of electrons from the valence shell of an atom of lower electronegativity to the valence shell of an atom of higher electronegativity.
Example: Na and F
Covalent Bonds
A covalent bond is a chemical bond formed between atoms by the sharing of one or more pairs of electrons to give a noble gas configuration at each atom. The simplest example of a covalent bond occurs in the hydrogen molecule(non-polar) and HCl (polar)

28. Bond length
The distance between nuclei in a covalent bond in Picometers
(pm; 1 pm=10-12)
Ångstrom A0
1A0 = 10-10

29. Bond dipole moment (m)
A measure of the polarity of a covalent bond. It is the product of the charge on either atom of a polar covalent bond times the distance between the nuclei.
and is given the symbol µ (Greek mu).
The SI unit for a dipole moment is the coulomb meter, but
they are commonly reported instead in a derived unit called the Debye (D: 1 D = 3.34 X C . m).
Bond dipole moment is
defined as “The product of the charge, e (either the δ- or δ + because each is the same in absolute magnitude), on one of its atoms X the distance, d, separating the two atoms.

30. p=qd
where d is the displacement vector pointing from the negative charge to the positive charge.
Thus, the electric dipole moment vector p points from the negative charge to the positive charge.

31. Lewis (electron dot) structures
Formal charge
(HNCO HCN HNC )

32. Charge on an atom/molecule is formal charge, i.e. H3O+, HO-
To assign Formal Charge:
1. Write correct Lewis structure
2. Assign each atom: all non-bonding e's half the shared e’s
3. Compare this number to
valence e-s in neutral unbonded atom.
August 27, 2002 (2)

33. number of valence electrons
Formal Charge on an atom
Formal Charge -
number of valence electrons =
unshared electrons +
1/2 of shared electrons
August 27, 2002 (2)

34. e’s belong to the first N =+1 -1
valence of N = 5
e’s belong to the first N =
[2 + 1/2(6)] = 5
5 - 5 = 0 charge
e’s belong to the 2nd N =
[1/2(8)] = 4
5 - 4 = +1 charge
e’s belong to the 3rd N =
[4 + 1/2(4)] = 6
5 - 6 = -1 charge

35. COVALENT BONDING & SHAPES OF MOLECULES
Chapter 1
COVALENT BONDING & SHAPES OF MOLECULES
1.1 Electronic Structure of Atoms
1.2 Lewis Model (Octet Rule, Formal Charge)
1.3 Functional Groups
1.4 Bonding Angles and Shapes of Molecules
1.5 Polar and Nonpolar Molecules
1.7 Molecular Orbital & Valence Bond Theory, Covalent Bonds, Hybridization (sp3, sp2, sp)
SUMMARY and OVERVIEW 1-11
Aug 22, 2002 (1)
Jan 21(14) lecture 1

36. Functional Groups: atom(s) bonded to C having characteristic properties
Determine:
reactions
properties
basis of nomenclature & classification
End of lecture 1 8/23/01 started functional groups

37. Functional Groups: atom(s) bonded to C having characteristic properties
here
methyl,1o, 2o, 3o
see Hs & Cs
alcohols: hydroxyl group
1o, 2o or 3o by
Hs on N
amines:amino group

38. Carboxylic Amide - (carbonyl + amine -C(O)N-)
FUNCTIONAL GROUPS
alcohols: (hydroxyl group, sp3-> C-O-H)
1o, 2o, 3o
amines:
1o, 2o, 3o
Aldehyde / Ketone (carbonyl group C=O)
Carboxylic Acid (carbonyl + hydroxyl group -CO2H)
Carboxylic Ester - (carbonyl + alcohol -CO2R)
Carboxylic Amide - (carbonyl + amine -C(O)N-)
Others ethers, halides, etc.
Formulas: complete, condensed and/or line

39. VSEPR - electrons (e-) in bonds/orbitals repel each other
 4 groups/bonds repulsion yields a tetrahedral shape ~109.5o
 3 bonds repulsion yields a trigonal shape ~120o
 2 bonds repulsion yields a linear shape 180o
eg: CH4 or NH3, HCO2H, CO2

40. and Nonpolar Molecules
Polar if:
(1) has polar bonds
(2) if the arrangement is “irregular” + - + -

45. 3 sp2 and p orbitals on axis
Orbitals for sp2 hybridization
Sept 3, 2002 lecture 3
3 sp2 and p orbitals on axis
hybrid

46. 3 sp2 and p orbitals on axis

48. sp sp py sp HYBRIDIZATION pz s two sp hybrid orbitals px

49. End of Chapter 1  2