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Unformatted text preview: BIS103 (002) Abel Biological Sciences
(CRN 65050) Spring Quarter 2009 Tuesdays/Thursdays 1:40 3:00 p.m. (198 Young Hall) Lectures will be recorded and posted on class website (podcast). Dr. Steffen Abel, Professor Phone: 752-5549 Email: email@example.com Office Hours: Fridays (9:30 - 11:30 am), Location: 210 Asmundson Hall (Office); for larger groups of students, 242 Asmundson Hall (Conference Room). Other appointments are available upon request. Rebecca Shipman Email: firstname.lastname@example.org Office Hours: Wednesdays (11 am - 1 pm) Katrina Edgar Email: email@example.com Office Hours: Mondays (1-3 pm) Location: Asmundson Hall (Conference Room Annex, 2. Floor). One of these three textbooks is REQUIRED: Biochemistry by Garrett & Grisham Principles of Biochemistry by Lehninger, Nelson & Cox Fundamentals of Biochemistry by Voet, Voet & Pratt Instructor: Teaching Assistants: Text: ... any other decent biochemistry text will do ... Booklet: A collection of visual aids specifically prepared for this class is available for purchase at Campus Books, which you need to bring to every lecture. However, this booklet does not substitute for one of the required textbooks! MyUCDavis (BIS103-002) Go to The "Real" BIS103 Website (http://www.plantsciences.ucdavis.edu/bis103) Your final grade will be calculated based on your performance in three examinations: Midterm 1 (33.3%), Midterm 2 (33.3%), and the final test (33.3 %). Parts of the final exam will be cumulative. Regrade requests for midterms must be submitted in writing within one week after return of the tests and will only be considered if you have used permanent ink during the exam. Exams will be given at the assigned times only! No early finals will be given! Exceptions may be granted to students with documented verification of personal loss or sickness and if the instructor was contacted before the exam. MIDTERM 1 MIDTERM 2 FINAL EXAM Tuesday, April 21 (in class) Tuesday, May 19 (in class) Saturday, June 6, 10:30 am - 12:30 pm Class Websites: Exams: How to "study"? Attend lectures Review lectures (your notes, podcast, slides) Stick your nose into a text!!! Office hours (read first) Attend review sessions Do NOT fall behind! Class Context
General and Organic Chemistry Thermodynamics, redox reactions, nucleophiles, electrophiles, major organic compound classes, chemistry of carbonyls BIS 102 "Structural" Biochemistry (carbohydrates, proteins, lipids, nucleic acids, enzymology) BIS 103 "Functional" Biochemistry (intermediary or primary metabolism) Nutrition, Microbiology, Human Physiology, Neurobiology, Exercise Biology, Plant Biology, Pharmacology, Cell Biology, etc. Lecture 1 Topics Why is there a need for metabolism? How is chemical energy stored? Directionality of metabolism
Gibb's Free Energy (G) Reduction Potential (E) Lecture 1 Topics Why is there a need for metabolism? How is chemical energy stored? Directionality of metabolism
Gibb's Free Energy (G) Reduction Potential (E) Fuel or Food + O2 CO2 + H2O + U
One Reaction John Candy's Metabolism Many Reactions Fuel Energy Action Food Energy Regeneration Reproduction Carbon Chemistry Chemical (Potential) Energy Fossil Fuels Hydrogen U Exhaust (CO2, H2O) U = q (Heat) + work (PV) Chemical (Potential) Energy Food H "trapped" as ATP
Intermediates CO2, H2O U = q (Heat) + work (PV)
In living cells, P and V are assumed to be constant U = q (Heat) + "chemical work" = H (Enthalpy) What is the purpose of metabolism? Provides energy (motion, transport, syntheses, heat) Provides building blocks for regeneration/growth Energy conversion compatible with C-based life All life forms are likely based on carbon chemistry Lecture 1 Topics Why is there a need for metabolism? How is chemical energy stored? Directionality of metabolism
Gibb's Free Energy (G) Reduction Potential (E) How is energy stored? In chemical bonds!
Bond Enthalpies (kJ/mol) 0 C, O, H, N, S, C -C - 500 C -H C -N C -S C -O O=O N-H S-H H -O Photosynthesis H-O-H Biological Oxidation - 1,000
Metabolic Energy Light Energy - 1,500 O=C=O Table 2: The Bonds of Life (Enthalpies of Chemical Bonds)
Bond kJ/mol Bond kJ/mol Bond kJ/mol Bonds kJ/mol C-H C-S C-N C-C C-O H-H H-O N-H N-N N-O S-H S-S S-O O-O P-O 413 259 293 348 358 436 463 391 163 201 339 266 146 599 N=N N=O S=S S=O O=O 418 607 418 523 495 NN 941 C=N C=C C=O 615 614 799 CN CC CO 891 839 1,072 O=C=O 1,598 H-O-H 926 p. 2 Bond Energy ~ Bond Strength
The strength of a chemical bond depends on: relative electronegativities (affinity for electrons) distance of electrons from nuclei number of electrons shared nuclear charge +H X Y
H X Y Energy required to break a bond (+H) Energy released during bond formation (H) Table 1: Electronegativities of Biologically Important Elements Hydrogen (H) Carbon Nitrogen Oxygen 436 kJ/mol 463 kJ/mol 2.2 2.6 3.0 3.4 H O (P) (S) Phosphorus Sulfur (C) (N) (O) H H H H H O + 436 kJ/mol + 463 kJ/mol p. 2 A + 2B 3C + D + 2E H of a reaction is the sum of bond energies consumed during breaking of all reactant bonds and of bond energies released during formation of all product bonds. If H > 0: Endothermic reaction 0 If H < 0: Exothermic reaction 0 Example: Ethanol combustion/oxidation
C2H5OH + 3O2
HH H-C-C-O-H HH O=O O=O O=O O=C=O 2CO2 + 3H2O
O=C=O O-H-O O-H-O O-H-O Cracking all bonds (+H) (kJ/mol) 3,234 4,719 1,485 Forming all new bonds (H) (kJ/mol) 3,196 5,974 2,778 H = 1,255 kJ/mol (exothermic) Other Examples
C6H12O6 + 6O2
Glucose 6CO2 + 6H2O H = 2,639 kJ/mol 2H2 + O2 2H2O H = 485 kJ/mol H H H O H O H H O O H H Good molecules for storing energy:
carbon polymers, hydrocarbons, hydrogen "Fuel Values" of some foods and fuels (kJ/g)
Carbohydrates Proteins Lipids 17 17 38 Coal Crude Oil Hydrogen 32 45 142 Wood 18 H H H H Metabolic Oxidation -C--C--C--CH H H H Carbohydrates Proteins Lipids "Harvest" hydrogen and transport as a "biologically safe" form NAD--H + H+ Most H (ATP) O=O H 2O CO2 FOOD "High Energy" Substrates Cellular Macromolecules Catabolism
H (ATP) Motion Transport Heat "Low Energy" Products Waste H (ATP) Anabolism Intermediates Precursor Molecules Lecture 1 Topics Why is there a need for metabolism? How is chemical energy stored? Directionality of metabolism
Gibb's Free Energy (G) Reduction Potential (E) Gibbs Free Energy (G) G = H TS
If G = 0: If G < 0: If G > 0: System is at equilibrium Exergonic ("downhill" process) Endergonic ("uphill" process) Process can be driven by H, +S, or both
Example: C6H12O6 + 6O2 7 Molecules 6CO2 + 6H2O 12 Molecules H: 2,639 kJ mol-1 S: positive (more disorder, fragmentation) G: 2,840 kJ mol-1 See p. 5 Glucose + 6O2 Glucose + 6O2
John Candy's Metabolism One Step 6CO2 + 6H2O 6CO2 + 6H2O Alternative way to calculate G
Reversible Reactions Reactants [R]
G: Products [P] Measure for the displacement of a reaction from its equilibrium (EQ)
Reaction is (already) at EQ Not (yet) at EQ: Forward reaction is favored Not (yet) at EQ: Reverse reaction is favored If G = 0: If G < 0: If G > 0: See p. 5 R P G = RTlnQ RTlnKeq Q = [P]initial/[R]initial Keq = [P]eq/[R]eq RTlnKeq = Go' G = Go' + RTlnQ Go' determined at "Standard Conditions" o 25 oC (298 K) 1 M (or 1 atm ) of R and P; or equimolar concentrations pH 7 (10-7 M H+) 55.5 M H2O 1 mM Mg2+ (if part of the reaction) ' See p. 5 Question: Under standard conditions, how far are conditions equimolar concentrations from equilibrium (Go') ? RRR ? PPP At thermodynamic equilibrium If Keq = 1 If Keq > 1 If Keq < 1 Go' = 0 Go' < 0 Go' > 0 RRR R RRRR PPP PPPPP PP See p. 5 Go' = RT lnKeq
Go' (kJmol-1) - 17.1 - 11.4 - 5.7 0 5.7 11.4 17.1 Keq 1,000 100 10 1 0.1 0.01 0.001 Go' = ~1,300 kJ mol-1 Go' = 2,840 kJ mol-1 Keq ~ 10220 Keq ~ 10500 C2H5OH + 3O2 C6H12O6 + 6O2 2CO2 + 3H2O 6CO2 + 6H2O Go': useful as a general guide to predict direction of a process (reaction) However, real conditions in cells often differ from standard conditions! Temperature: 37 oC (310 K) Cellular or initial concentrations of R and P
(they are unlikely to be at equilibrium) G: G = Go' + RTlnQ Go' = RTlnKeq Q = [P]initial/[R]initial See p. 5 G = RT lnQ RT lnKeq = RT ln Q/Keq If Q = Keq Reaction is already at EQ
(initial P/R ratio equals P/R ratio at equilibrium) G = 0 If Q < Keq Reaction is NOT yet at EQ
(relatively more R or less P, P P formation favored) G < 0 If Q > Keq Reaction is NOT yet at EQ
(relatively less R or more P, P R formation favored) G > 0 See p. 5 Reduction Potential (E)
... yet another way to express G... Metabolic reactions are often redox reactions, involving transfer of electrons from a donor to an acceptor 1. Direct combination with oxygen (X X + O=O 2. Transfer of the hydride anion (H--H, or H- + H+) 3. Transfer of hydrogen (H, or e- + H+) 4. Direct transfer (e-) 2O X) Redox reactions can be written as two "half-reactions"
(by convention: each is written in the direction of the reduction!) convention reduction Aoxidized + Breduced
1. Aoxidized + e2. Boxidized + e- Areduced + Boxidized
Areduced Breduced Electronegativities can predict direction of e- transfer See p. 6 Question: Which "half-reaction" has the higher affinity for electrons at standard conditions ?
Reference Electrode: H+ + eH2 "half-reaction" (1M each) (Eo= 0.00 V) Test Electrode:
By convention: A+ + eA (Eo'= ??? V) "half-reaction" (1M each) If e- flow from reference to test ("test" is stronger e- acceptor): Eo > 0 (+V) If e- flow from test to reference ("test" is weaker e- acceptor): Eo < 0 (- V) See p. 6 Table 3: Standard Reduction Potentials
Half-reaction (written as reduction by convention) Excited (Chlorophyll a)2* Acetate + 2H+ + 2eNAD+ + 2H+ + 2ePyruvate + 2H+ + 2e2H+ + 2eAcetaldehyde + H2O NADH + H+ Lactate Eo (V) ~ - 1.00 - 0.58 - 0.32 - 0.18 e- flow H2 (at standard conditions, 1M each, pH 0) NO2- + H2O 2H2O (Chlorophyll a)2 0.00 + 0.42 + 0.82 + 1.10 NO3- + 2H+ +2eO2 + 4H+ + 4e- (Chlorophyll a)2.+ + e- See p. 3 Reduction Potential E for each "Half-Reaction" (Nernst Equation) 1. Aoxidized + eAreduced E= Eo' RT [e- Acceptor, Aox] + nF ln [e- Donor, Ared] RT [e- Acceptor, Box] + nF ln [e- Donor, Bred] 2. Boxidized + e- Breduced E= Eo' E of Redox Reaction (Aox + Bred Ared + Box) E = EOxidant (A) EReductant (B) Relationship between E and G G = n F E E = G/nF Go' = n F Eo' Eo' = Go'/nF See p. 6 ...
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This note was uploaded on 09/06/2009 for the course BIS 103 taught by Professor Abel during the Spring '08 term at UC Davis.
- Spring '08