1. Why chemistry, why again?
Because every topic we study in biology leads back to and requires an understanding of chemistry.

2. You are what you eat! Energy and Building blocks for
You are made up of biomolecules!
Your biomolecules are build from food you eat
Your body breaks them down and rearranges them to give you:
Energy and
Building blocks for
proteins,
carbohydrates,
lipids and
Nucleic acids (DNA, RNA)

3. What do you eat?

4. Food – a mixture of biomolecules
Carbohydrates (sugar and starch)
Protein
Lipids – (fats, oils, cholesterol)
Vitamins
Minerals (Ca, PO4-2, Se, Fe, Ni, Zn, etc)
H2O
and
Energy

5. Let’s review the basics - atoms
element
Carbohydrate
Lipids
Protein
Nucleic acids C  H O N P * S
**
* phospholipids
** some proteins

6. Atomic Structure
Protons
Electrons
Neutrons

7. Valence Electrons Anyone?
What are they?
How do I figure out how many an atom has?
How do I figure out how many an ion has?

8. The Ultimate Graphic Organizer

9. Lewis Dot Structures Anyone?

10. The top 18 Lewis Dot Structures

11. Molecular
Forces
Intramolecular
Forces
Intermolecular
Forces
Ionic Bonds
Covalent
Bonds
Dispersion
Forces
Dipole
Forces
Hydrogen
Bonds

12. Ionic and Covalent Bonding

13. Which Bond – When? it’s all about Electronegativity
Definition – the “grabbiness” of an atom for an electron;
the more grabby, the greater the EN value
ex, F = 4.0; Na = 0.9; C = 2.5; H = 2.1; O = 3.5
Use EN difference to determine bond type
Ionic bonds – EN difference ≥ 1.5
Covalent Bonds - EN Difference < 1.5

15. Covalent Bond Polarity
Nonpolar Covalent bonds – equal sharing
0.0 < Electronegativity difference ≤ 0.4
Polar Covalent Bond – unequal sharing
0.4 < Electronegativity difference < 1.5
Line becomes arrow from positive to negative
Insulin

17. Practice – Which Bond Type?
Chemical Bond
EN difference
I, PC, NPC
H-F
|2.1 – 3.98|= 1.88 I
C-H
|2.55 – 2.1|= 0.45
NPC
Cl-O
K-Cl
O-H
C-O
S-O
Br-Br

18. Molecular Polarity

19. Molecular Polarity examples
CH4
CH3Cl
HCl
NH3
H2O

20. Molecular
Forces
Intramolecular
Forces
Inter molecular
Forces
Ionic Bonds
Covalent
Bonds
Dispersion
Forces
Hydrogen
Bonds
Dipole
Forces

21. Intermolecular Forces
Bond type
Dissociation energy (kcal),[2][3]
Covalent
400
Hydrogen bonds
12–16
Dipole–dipole
0.5–2
Dispersion Forces
<1

22. Dispersion Weakest intermolecular force
Caused by the motion of electrons
More electrons per molecule, more attraction between molecules; i.e. halogens

23. Intermolecular Forces - Dispersion
Electrons momentarily gang up on one side of the molecule
Dispersion forces influence boiling point in nonpolar covalently bonded compounds. The greater the number of electrons in the molecule, the stronger the dispersion forces and hence the more energy requested to cause a change of phase (liquid to gas). Dispersion forces are the primary intermolecular attraction forces in nonpolar molecules.
e.g. bromine is a liquid, whereas iodine is a solid at STP

24. Intermolecular Forces - Dispersion
“Dispersion forces. Dispersion forces arise when a normally nonpolar atom becomes momentarily polar due to an uneven distribution of electrons, leading to an instantaneous dipole that induces a shift of electrons in a neighboring nonpolar atom. Dispersion forces are weak but can be important when other types of interactions are either missing or minimal (part (d) of Figure 18.6, “Tertiary Protein Structure Interactions”). This is the case with fibroin, the major protein in silk, in which a high proportion of amino acids in the protein have nonpolar side chains. The term hydrophobic interaction is often misused as a synonym for dispersion forces. Hydrophobic interactions arise because water molecules engage in hydrogen bonding with other water molecules (or groups in proteins capable of hydrogen bonding). Because nonpolar groups cannot engage in hydrogen bonding, the protein folds in such a way that these groups are buried in the interior part of the protein structure, minimizing their contact with water.”

25. Intermolecular Forces – Dipole Interactions
attraction between polar molecules
defines the behavior of many biological compounds

26. Intermolecular Forces – Hydrogen Bonds
Strongest of the intermolecular forces
Only molecules with hydrogen in them
BIG role in living organisms!!!
Helical protein structures are stabilized by intrachain hydrogen bonding between the carbonyl oxygen atom of one amino acid and the amide hydrogen atom four amino acids up the chain (located on the next turn of the helix). This figure represents a right-handed α-helix.
“Hydrogen bond patterns in beta-sheets. A four-stranded beta-sheet is drawn schematically which contains three antiparallel and one parallel strand. Hydrogen bonds are indicated with red lines (antiparallel strands) and blue lines (parallel strands) connecting the hydrogen and receptor oxygen. “

27. Hydrogen Bonds

28. Intermolecular Forces
Hydrogen Bonding

29. Hydrophilic - Hydrophobic

30. Water – remember your chemistry?
H2O
What’s the H, what’s the O?
Why the 2?
Hydrogen Bonding – a big deal!
solubility (water is the universal solvent)
cohesion
adhesion
heat retention

31. Amazing Water Properties
Surface Tension
Capillary Action
Density Differences
Solubility

32. Water Properties – Surface Tension

33. Water Properties- Capillary Action
425 ft
Paper Towels and Redwood Trees

34. Water Properties - Density

35. Water Properties - Solubility
Life’s necessity

36. Molecular
Forces
Intramolecular
Forces
Intermolecular
Forces
Ionic Bonds
Covalent
Bonds
Dispersion
Forces
Dipole
Forces
Hydrogen
Bonds