1. Biochemistry: The Chemical Basis of Life Mr. Nichols PHHS

2. Gathering Information  # of Protons: Atomic Number, easy!  # of Electrons: Atoms by definition have no charge so the number of protons is too the number of electrons.  # of Neutrons: Atomic weight is the total number of atoms in the nucleus.  Example: Krypton’s mass number is 83.80, round to 84 for the neutron calculation. So….  For Krypton: 84=(number of protons)+(number of neutrons), SOLVE!??!?!  Practice, solve for: K, C, O and Br. (On your notes)

4. Octet Rule A valence electron is an electron associated with an atom that can participate in the formation of a chemical bond (Covalent mostly). “OCT”- Meaning 8

5. Where are valence electrons? Outer most orbital, highest overall energy (most potential energy)

6. Finding the Number of Valence Electrons

7. Electronegativity:HOW BAD AN ATOM WANTS DAT ELECTRON(Z)!!!!

9. Structure of Atoms and Bonding (Covalence)

10. Energy Get your mind right, prepare yourself!

11. Types of Energy

12. In General

13. Chemical Equilibrium, Exo- endothermic

14. Bond Energies

15. Try it out! Calculate the bond energies and determine whether the reaction is exothermic or endothermic.

16. Potential Energy (E p )- An object gains or loses its ability to store energy based on it’s position.

18. More Examples!

19. Potential Energy in Atoms!

20. Kinetic Energy ( E k )and Thermal Energy (T) Kinetic energy on a molecular basis are due to the fact that atoms/molecules are in constant motion, this is also referred to as Thermal Energy: How much thermal energy an object has equates to its temperature.

21. 1 st Law of Thermodynamics

22. Entropy (ΔS) and the 2 nd Law of Thermodynamics

23. Entropy

24. Entropy Examples

25. Another Entropic Example

26. Molecular Chaos and Order

27. Chaos Vs. Order=Potential Energy

28. Gibbs Free Energy Equation ΔG = ΔH – TΔS  ΔG-Free Energy (The energy associated with a chemical reaction that can be used to do work)  ΔH-Enthalpy The change ΔH is positive in endothermic reactions, and negative in heat- releasing exothermic processes.  T-Temperature in Kelvin (273+Degrees Celsius) Please Note: To get from Fahrenheit to Celsius you must subtract 32 from the Fahrenheit temperature, multiply that number by 5 then divide by 9= Temp in Celsius.  ΔS-Entropy of a system.

29. Used to predict the spontaneity of a reaction!  Chemical reactions are spontaneous if they proceed on their own, without any continuous external influence such as added energy. Two factors determine whether a reaction is or isn’t spontaneous.  1.) Reactions tend to be spontaneous if the products have a lower potential energy than the reactants.  2.) Reactions tend to be spontaneous when the products molecules are less ordered than the reactant molecules.

30. More Detail!!! The value of G o for a reaction measures the difference between the free energies of the reactants and products.

31. Exothermic and Endothermic with (G)

32. Try these Problems! Calculate ΔG Problem 1  T=298.15 K  Δ S=.1087 kJ/K  Δ H=28.05 kJ  Problem 2  T=773.15 K  Δ S=.19875 kJ/K  Δ H=-92.22 kJ

33. Importance and Properties of Carbon This will require outside of class mastery and “memorization of functional groups.”

34. Why Carbon? Carbon offers 4 valence electrons which allows for robust bonding options. Carbon can be bonded into linkages or rings and is very stable when bonded with nearly all elements.

35. Functional Groups + Carbon=Awesome (aka Life) “Ketone”

36. Six Main Functional Groups Commonly Attached to Carbon

37. All orbitals are not created equal! Fun Fact: Electrons are extremely small, whereas protons and neutrons are roughly the same size. (Neutrons slightly larger)

38. S, P, D, F Orbitals http://www.youtube.com/watch?v=sMt5 Dcex0kg

39. Draw the molecule and the dipole directionality. Then state polar or non- polar  CCl 4 C2H4C2H4  SO 2  NH 3 H2SH2S

40. Polarity  Definition: Describes how equally bonding electrons are shared between atoms. (atoms and bonds can be described as being polar vs. non-polar.

41. Dipole  A dipole when you have a positive end and a negative end on the same molecule. Remember the more negative end is where the more electronegative element is.

42. Charting the Dipole (Molecular Geometry)

43. Dipole-Dipole Interactions (Very Weak)

44. Hydrogen Bonds: Most common, most important to Biology.

45. Hydrogen Bonds Continued

46. Ionic Bonds: A type of chemical bond formed through an electrostatic attraction between two oppositely charged ions.

47. Ion Formation.

48. Properties of Water- 1.) Solvent

49. 2.) Cohesion & Adhesion: Not exclusive to water.

50. Surface Tension: Not exclusive to water.

51. Acid-Base Reactions and pH

53. Key Idea

54. Examples

56. What exactly is pH and what are acids and/or bases?  Fun fact: Measuring pH is quite mathematic, so much so that pH stands for the negative logarithm of the hydrogen ion concentration.

57. The very basics, the essentials!

58. Challenge Problem 1!!!! Consider the following equation: 4NH 3 +3O 2  2N 2 + 6H 2 O 1.) Draw each of the molecules (reactants and products), state whether it’s polar or not and if it has a dipole or not (draw dipole directionality arrows). 2.) Calculate the bond energy of the reactants and products. Determine whether the reaction is exothermic or endothermic, give an explanation.

59. Challenge Problem 2 PROVE THAT THIS GRAPH IS CORRECT.

60. Challenge Problem 3 1.) 2.) 4.) 3.)