CH08a_2_ - Metabolism Overview Overview Definitions&...

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Unformatted text preview: Metabolism Overview Overview Definitions & Perspectives Catabolism & Anabolism Kinetic Energy & Potential Energy Two Laws of Thermod namics Th dy Free Energy: G & G Metabolism Overview Overview Spontaneous vs. NonNonSpontaneous reactions Exergonic vs. Endergonic ATP & Cellular Work ATP-Coupling ATP ATP Cycle Bio 230, Summer 2010, Ch8a, Page 1 Metabolism Metabolism is a series of ordered, ordered, controlled chemical reactions within \ among cells Manages the flow of material and energy within the cell to th avoid threatening deficiencies and useless surpluses Metabolic Chart Bio 230, Summer 2010, Ch8a, Page 2 Metabolism Uses enzymes to control synthesis and breakdown of substances Metabolism = Catabolism Catabolism & Anabolism Two Forms of Metabolism Catabolism: breaking down of complex molecules to less complex molecules to less complex complex ones to yield energy for cellular work Downhill; “spontaneous” Example: Cellular respiration O2+ glucose CO2 + H2O + energy Bio 230, Summer 2010, Ch8a, Page 3 Two Forms of Metabolism Anabolism: using energy to assemble complex molecules from assemble complex molecules from simpler simpler ones Uphill; “non-spontaneous” “non Example: protein synthesis from amino acids AA1 + AA2 + energy protein + H2O Two Forms of Energy Energy does work on matter, thereby changing the way it is thereby changing the way it is arranged arranged Kinetic Energy is evident as matter is changing Examples: a moving object; an ongoing chemical reaction; light Bio 230, Summer 2010, Ch8a, Page 4 Two Forms of Energy Potential Energy is stored and gy available for conversion into kinetic energy Examples: person at top of slide, car battery lid Cellular Examples Kinetic Energy Muscle contraction Light hitting chloroplasts Ions moving through channel Bio 230, Summer 2010, Ch8a, Page 5 Cellular Examples Potential Energy Storage Glucose Starch Fat Water Potential - concentration gradients Mechanical Energy More Potential Energy Less Stable Kinetic Energy Less Potential Energy More Stable Fig. 8-5a 8Bio 230, Summer 2010, Ch8a, Page 6 Transport Energy More Potential Energy Less Stable Kinetic Energy Less Potential Energy More Stable Fig. 8-5b 8- Chemical Energy More Potential Energy Less Stable Kinetic Energy Less Potential Energy More Stable Fig. 8-5c 8Bio 230, Summer 2010, Ch8a, Page 7 Laws of Thermodynamics First Law: Energy can be interconverted, interconverted, but total energy in the universe is constant Second Law: Each energy transfer makes the universe bit transfer makes the universe a bit more more disordered; increases entropy (S) Laws of Thermodynamics Example: light hitting plant cell No net energy is lost or gained by the universe Some light changed into chemical energy stored in glucose (decreases S) Bio 230, Summer 2010, Ch8a, Page 8 Laws of Thermodynamics Example: light hitting plant cell Source (the sun) becomes more disordered (increases S) All light energy eventually dissipated as heat (increases S) Net entropy of universe increases a bit Potential & Kinetic Energy Fig. 8-5 8- Mechan. Transport Chemical High P.E. K.E. Low P.E. Bio 230, Summer 2010, Ch8a, Page 9 Free Energy From 1st and 2nd Laws of Thermodynamics, we can define Thermodynamics, we can define Free Free Energy (G) as: The portion of a system’s total potential energy that can perform work wor When temperature is uniform within system Free Energy Units kcal or kcal/mole 1,000 cal = 1 kcal = 1 Cal 1 cal = 4.32 joules cal 4.32 joules Bio 230, Summer 2010, Ch8a, Page 10 Spontaneous Reactions G is High (100) K.E. G is G is Low (50) Fig. 8-5 8- Free Energy In spontaneous reactions, energy is released G decreases, Gf < Gb f = final, b = beginning Therefore G is negative G = (Gf - Gb) < 0 Example: Gb = 100, Gf = 50 G = 50 - 100 = -50 Bio 230, Summer 2010, Ch8a, Page 11 NonNon-Spontaneous Reactions G is Low (50) Work G is + G is High (100) Fig. 8-5 (modified) 8- Free Energy In non-spontaneous reactions, nonenergy is absorbed abso G increases, Gf > Gb f = final, b = beginning Therefore G is positive G = (Gf - Gb) > 0 Example: Gb = 50, Gf = 100 G = 100 - 50 = +50 Bio 230, Summer 2010, Ch8a, Page 12 Free Energy G = H - TS G is free energy; able to do work H is total energy = enthalpy S is entropy (disorder) T is temperature (oK = oC + 273) is temperature 273) Some of a system’s total energy is “locked up” as disorder; unable to do work Spontaneous Reactions G = H - TS (-) = (-) - (+) Exergonic Energy is released from chemical bonds (H) and from increased disorder (S) Bio 230, Summer 2010, Ch8a, Page 13 NonNon-Spontaneous Reactions G = H - TS (+) = (+) - (-) Endergonic Energy is absorbed into chemical bonds (H) and into decreased disorder (S) Example: Cellular Respiration C6H12O6 6CO2 + 6H2O G = -686 kcal/mol G H T S -686 = -673-(300)(0.043) = -673 -13 673- Bio 230, Summer 2010, Ch8a, Page 14 Interpretation H = -673 kCal/mole released from from energy stored in chemical bonds -TS = -13 kCal/mole released from energy stored as from energy stored as molecular molecular orderliness Example: Photosynthesis 6CO2 + 6H2O C6H12O6 G = +686 kcal/mol G H T S +686 = +673-(300)(-0.043) = +673 +13 +673-(300)(- Bio 230, Summer 2010, Ch8a, Page 15 Interpretation H = +673 kCal/mole of energy energy stored into chemical bonds -TS = +13 kCal/mole of energy stored as increased energy stored as increased molecular molecular orderliness Exergonic and Endergonic Reactions Exergonic, G < 0 A B Endergonic, G > 0 B A At equilibrium, G = 0, forward and reverse rates are same A B Bio 230, Summer 2010, Ch8a, Page 16 Terms Terms and Units T erm G H S T -T S D efin itio n U n its " free" en erg y , en erg y K ca l/m o l a v a ila b le t o d o w o rk to t o ta l en erg y sto red in K ca l/m o l ch em ica l b o n d s e n tro p y , d iso rd erlin ess K ca l/m o l o K o f m o lecu le t em p era tu re; a v era g e o K sp eed o f m o lecu les K ca l/m o l e n erg y lo st to m o lecu la r o d iso rd erlin ess ( K x K ca l/m o l o K ) ch a n g e = (v a lu e a t en d – Sam e as G , H or S v a lu e a t b eg in n in g fo r G , H or S Final: Final: 6CO2 + 6H2O G H T S 314 = 527-(300)(0.710) = 527 - 213 527Begin: C6H12O6 G H T S 1000 = 1200-(300)(0.667) = 1200 - 200 1200Change (Final-Begin) (FinalG H T S -686 = -673-(300)(0.043) = -673 -13 673Bio 230, Summer 2010, Ch8a, Page 17 Final: C6H12O6 G H T S 1000 = 1200-(300)(0.667) = 1200 - 200 1200Begin: 6CO2 + 6H2O G H T S 314 = 527-(300)(0.710) = 527 - 213 527Change (Final-Begin) (FinalG H T S +686 = +673-(300)(-0.043) = +673 +13 +673-(300)(- Spontaneous Reactions S is Low G is High (100) K.E. G = -50 50 S is High G is Low (50) Fig. 8-5 8- S = ?? Bio 230, Summer 2010, Ch8a, Page 18 Most Spontaneous Reactions G = H - TS (-) = (-) - (+) Exergonic Energy is released from chemical bonds (H) and from increased disorder (S) NonNon-Spontaneous Reactions S is High G is Low (50) Work G = +50 +50 S is Low G is High (100) Fig. 8-5 (modified) 8- S = ?? Bio 230, Summer 2010, Ch8a, Page 19 Most Non-Spontaneous NonReactions G = H - TS (+) = (+) - (-) Endergonic Energy is absorbed into chemical bonds (H) and into decreased disorder (S) Coupled Coupled Reactions Endergonic reactions require energy to be added for them to occur To drive endergonic reactions, cells couple them to exergonic cells reactions Bio 230, Summer 2010, Ch8a, Page 20 G= -7.3 kcal/mol (exergonic) ATP ADP+Pi Fig. 8-10 8- (endergonic) Glutamic Acid + NH3 Glutamine ATP ATP ADP + Pi G = +3.4 G = -7 . 3 Glutamic Acid Glutamine + G = -3.9 + NH3 + ATP ADP + Pi ATP ADP+Pi Fig. 8-10: 8Energy Coupling Mechanism Bio 230, Summer 2010, Ch8a, Page 21 ATP ADP+Pi Fig. 8-10: 8Energy Coupling ATP and Cellular Work Cells do three kinds of work: Mechanical: movement, contraction, beating Transport: pump molecules against concentration gradient Chemical: drive endergonic reactions; e.g. building protein from amino acids Bio 230, Summer 2010, Ch8a, Page 22 Fig. 8-11: Cellular Work 8- ATP and Cellular Work Cells use ATP to drive most cellular work Energy stored in PO4 -PO4 bonds ATP ADP + Pi G = - 7.3 kcal/mol Adenine H2O P Ribose ATP P P P PO4s ADP Bio 230, Summer 2010, Ch8a, Page 23 P + P Pi Fig. 8-12: The ATP Cycle 8- Exergonic Reactions Endergonic Reactions Bio 230, Summer 2010, Ch8a, Page 24 ...
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This note was uploaded on 11/28/2011 for the course BIOL 230 taught by Professor J.breckler during the Spring '11 term at S.F. State.

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