1. Biochemistry: The Chemical Basis of Life
Dr. Nichols Coronado HS

2. 1.2- An Element's Properties Depends on the Structure of its Atoms
1.1-Matter Consists of chemical elements in pure form and in combination called compounds
1.2- An Element's Properties Depends on the Structure of its Atoms
1.3-The Formation and Function of Molecules Depend on Chemical Bonding Btw. Atoms

3. 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)

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

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

7. Finding the Number of Valence Electrons

8. Electronegativity: The strength of an atoms pull for an electron.

10. Structure of Atoms and Bonding (Covalence)

11. Get your mind right, prepare yourself!
Energy
Get your mind right, prepare yourself!

12. Types of Energy

13. In General

14. Chemical Equilibrium, Exo-endothermic

15. Bond Energies

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

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

19. More Examples!

20. Potential Energy in Atoms!

21. Kinetic Energy (Ek)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.

22. 1st Law of Thermodynamics

23. 1.4- Chemical Reactions Make and Break Chemical Bonds
1.5- Polar Covalent Bonds in Water Molecules result in hydrogen bonding
1.4- Chemical Reactions Make and Break Chemical Bonds

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

25. Entropy

26. Entropy Examples

27. Another Entropic Example

28. Molecular Chaos and Order

29. Chaos Vs. Order=Potential Energy

30. 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.

31. 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.

32. More Detail!!!
The value of Go for a reaction measures the difference between the free energies of the reactants and products.

33. Exothermic and Endothermic with (G)

34. Try these Problems! Calculate ΔG
T= K
ΔS=.1087 kJ/K
ΔH=28.05 kJ
Problem 2
T= K
ΔS= kJ/K
ΔH= kJ

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

36. 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.

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

38. Six Main Functional Groups Commonly Attached to Carbon

39. Draw the molecule and the dipole directionality
Draw the molecule and the dipole directionality. Then state polar or non-polar
CCl4
C2H4
SO2
NH3
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. 1.7- Acidic and Basic Conditions Affect Living Organisms

48. Ion Formation.

49. 1.6- Four Emergent Properties of Water Contribute to Earth's suitability for life.

50. Properties of Water- 1.) Solvent

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

52. Surface Tension: Not exclusive to water.

53. Acid-Base Reactions and pH

55. Key Idea

56. Examples

58. 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.

59. The very basics, the essentials!

60. Challenge Problem 1!!!! Consider the following equation:
4NH3 +3O2  2N2 + 6H2O
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.

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