1. Introduction & Review: Orbitals, Bonds, Structures & Acidity
Highland Hall Biochem Block: Handout #1
Introduction & Review: Orbitals, Bonds, Structures & Acidity
Adapted from Organic Chemistry, L. G. Wade, Jr.

2. Particle Location Mass (amu)
Atomic Structure
3 components
Particle Location Mass (amu)
proton nucleus ~ 1
neutron nucleus ~ 1
electron surrounds ~ 1/1800
the nucleus
Electrons determine an atom’s reactivity
Recall that opposite charges attract each other, same charges repel…

3. Atomic Structure atomic number = number of protons
atomic weight ≈ # of protons + # of neutrons
typically #(protons) = #(neutrons)
Isotope: atoms with different number of neutrons than protons

4. Electronic Structure
Atomic orbital : space occupied by an electron associated with a nucleus
Electron density : statistical probability of finding the electron at a given location (density maps)
Orbital shells described by principal quantum number (n)
As n ↑, energy ↑, and (# of e-) ↑
Atomic orbitals are defined by three quantum numbers: n, ℓ, mℓ

5. Quantum Numbers
n : orbital shell, determines size + energy of an orbital
aka : principle quantum number
ℓ : orbital subshell, determines shape of an orbital
ranges from 0 to n-1, for a given n.
mℓ : magnetic number, determines spatial orientation of an orbital, an integer ranging from - ℓ, … , 0 , …, +ℓ
ms : spin of an electron (-½, +½) ℓ 1 2 3 4
orbital type

6. 1s Atomic Orbital e- density highest at nucleus
Spherically symmetrical
(Same in all three dimensions)

7. In-class exercise
Which is a better model for a 1s orbital – a basketball or an avocado?
Can you think of another object/food that might be an even better model than the two above?

8. 2s Atomic Orbital e- density highest at nucleus
spherically symmetrical
Has a node:
region of zero
e- density
Chapter 1

9. 2p Atomic Orbitals node (plane) at nucleus directional degenerate:
2px, 2pY, 2pz
degenerate:
Chapter 1

10. The Aufbau “building up” principle
Distribute the e- so that the energy is as low as possible
# of e- in an atom = its atomic number
Each added electron will enter orbitals in order of increasing energy
Any given orbital cannot take more than 2 e- (__________ _________ Principle)

11. Electronic Configurations
=>   
Hund’s rule:

12. Partial Periodic Table
Valence electrons:

13. Table: electronic configurations
=>

14. Octet Rule Filled shells of electrons are particularly stable
Example : noble gases
definition 1 : atoms transfer or share e- in such a way as to attain a filled shell
definition 2 : atoms undergo bonding by gaining, losing, or sharing e- so as to acquire the electron configuration arrangement of a noble gas.
For atoms in 2nd row of the periodic table,
a filled shell = 8 valence e-
“octet rule”

15. Ionic Bond Formation
Ionic bonding: electrons are transferred from one atom to another to attain noble gas configuration.
The ionic bond is the result of attraction between oppositely charged ions.

16. Covalent Bond Formation
Covalent bonding: electron pair is shared between 2 nuclei.
e- not transferred
Polar covalent bond: electron pair is unequally shared

17. Lewis Structures A way to symbolize covalent bonds in molecules
Bonding electrons : symbolized by a line
Nonbonding electrons/ lone pairs: symbolized by dots

18. Valence :

20. Multiple Bonding
Single bond : sharing of one pair of e- between two atoms
Double bond : sharing of two pairs of e- between two atoms
Triple bond : sharing of three pairs of e- between two atoms
Carbon has a valence of 4
=>

21. Multiple Bonding Double bond example : ethylene and formaldehyde
Triple bond example : acetylene
=>
Chapter 1

22. In class exercise: Drawing Lewis Structures
Chapter 1

23. Covalent Bond Types
Non-polar : bond between atoms in which electrons are shared equally
(C―C, H―H, C―H)
Polar : bond between atoms in which electrons are shared unequally
(O―H, C―Cl)
Ionic bonds are simply the most extreme case of a polar bond.
To predict polarity of a bond, one exploits electronegativities.
=>

24. Electronegativity
Electronegativity : ability of a bonded atom in a molecule to attract electrons.
Atoms with higher electronegativities pull the shared electrons more strongly.
Electronegativity increases
from left to right and from
bottom to top.

25. Electronegativity and Bond Polarity
The larger the difference in electronegativities between two atoms the more polar the covalent bond.
If DE.N. > 2, bond is ionic
To quantify how polar a bond is, calculate the dipole moment.
Electronegativity values:

26. Dipole Moment (μ) Dipole moment : amount of charge separation.
Charge separation shown by electrostatic potential map
Red = δ+ region
Blue = δ- region
Chapter 1

27. Structural Formulas CH3CHOHCH2CHFCH3 Full structural formula
(no lone pairs shown)
Line-angle formula
stick figure
ends of lines = carbon
assume enough H’s to give each carbon 4 bonds
show all heteroatoms
Condensed structural formula
not shown in actual bonding order
CH3CHOHCH2CHFCH3

28. Chapter 1

29. Drawing molecules in 3-D using the line-wedge formula

30. In class exercise: Make these molecules
What is the relationship between molecules A and B?
Between B and C?

31. BrØnsted-Lowry Acids and Bases
Acids can donate a proton (H+) in water
Bases can accept a proton (H+) in water
These reactions form conjugate acid-base pairs
H2O
acid base
H2O
acid base

32. Conjugate Acid-base Pairs
Conjugate acid : acid formed upon protonation of a base
Conjugate base : base formed upon the deprotonation of an acid
H2O
acid base
conjugate base conjugate acid
H2O
acid base
conjugate base conjugate acid

33. Acid and Base Strength Acid dissociation constant, Ka
Ka is the approximate strength of the acid.
Base dissociation constant, Kb
For conjugate pairs, (Ka)(Kb) = Kw
Spontaneous acid-base reactions proceed from stronger to weaker.
pKa 4.74
pKb 3.36
pKb 9.26
pKa 10.64

34. Acid Strength
Example : Water

35. Typical pKa values
pKa α 1/ acid strength

36. In class exercise: predicting acid base reactions which reactions are favorable?