Reaction, endergonic equilibrium

Both of these reactions are followed by exergonic reactions. The equilibrium of the reaction malate oxaloacetate (step 8) lies in favor of malate formation, so at equilibrium the concentration of oxaloacetate will be low. The next reaction in the cycle (oxaloacetate + acetyl-CoA — citrate) (step 1) is, however, exergonic, and the oxaloacetate is removed to condense with acetyl-CoA. Similarly, the conversion of citrate to isocitrate is endergonic, and at equilibrium the reaction favors the formation of citrate. The next reaction in the cycle (isocitrate—>2-oxoglutarate) is exergonic, and so the isocitrate is removed thus allowing this reaction to proceed. 

FIGURE 8.11 Energy profile of an endergonie reaction. Endergonic reactions are "uphill" processes in terms of energy, and the reactant is favored at equilibrium. 

Reactions for which the sign of AG° is negative are described as exergonic those for which AG° is positive are endergonic. Exergonic reactions have an equilibrium constant greater than 1 endergonic reactions have equilibrium constants less than 1. 

If AG is equal to 0, the process is at equilibrium, and there is no net flow either in the forward or reverse direction. When AG = 0, A.S = H/T, and the enthalpic and entropic changes are exactly balanced. Any process with a nonzero AG proceeds spontaneously to a final state of lower free energy. If AG is negative, the process proceeds spontaneously in the direction written. If AG is positive, the reaction or process proceeds spontaneously in the reverse direction. (The sign and value of AG do not allow us to determine how fast the process will go.) If the process has a negative AG, it is said to be exergonic, whereas processes with positive AG values are endergonic. 

AGr Gibbs free-cnergy change The energy difference between reactants and products. When AG° is negative, the reaction is exergonic, has a favorable equilibrium constant, and can occur spontaneously. When AGC is positive, the reaction is endergonic, has an unfavorable equilibrium constant, and cannot occur spontaneously. 

Enantiotopic (NMR), 455 Endergonic. 153 Endergonic reaction, Hammond postulate and, 197-198 Endo stereochemistry, Diels-Alder reaction and, 495 Endothermic, 154 -ene, alkene name ending, 176 Energy difference, equilibrium position and, 122... 

Every chemical reaction reaches after a time a state of equilibrium in which the forward and back reactions proceed at the same speed. The law of mass action describes the concentrations of the educts (A, B) and products (C, D) in equilibrium. The equilibrium constant K is directly related to the change in free enthalpy G involved in the reaction (see p.l6) under standard conditions (AG° = - R T In K). For any given concentrations, the lower equation applies. At AG < 0, the reaction proceeds spontaneously for as long as it takes for equilibrium to be reached (i.e., until AG = 0). At AG > 0, a spontaneous reaction is no longer possible (endergonic case see p.l6). In biochemistry, AG is usually related to pH 7, and this is indicated by the prime symbol (AG° or AG ). 

The second NH2 group of the later urea molecule is provided by aspartate, which condenses with citrulline into argininosucci-nate. ATP is cleaved into AMP and diphosphate (PPi) for this endergonic reaction. To shift the equilibrium of the reaction to the side of the product, diphosphate is removed from the equilibrium by hydrolysis. 

AG, is directly associated with the direction in which a particular chemical reaction can proceed. If AG < 0 for a given set of conditions of a particular reaction, then the reaction will proceed spontaneously in the indicated direction until equilibrium is reached. Conversely, if AG is positive, then energy will be needed to shift the reaction further from its equilibrium condition. See Helmholtz Energy Endergonic Exergonic Enthalpy Entropy Thermodynamics Biochemical Thermodynamics... 

This simple model predicts that the observed kinetics is determined by the rate of fragmentation only when the reverse process is much slower than counterdiffusion (i.e. when k f under activation control. On the other hand, for an endergonic fragmentation it is expected that k f k and /Cobs = The reaction now is described as a pre-equilibrium... 

Similarly, conjugated dehydrogenation of ethylbenzene is performed [6], At low temperature this is an endergonic reaction requiring high temperatures (800-900 °C), at which equilibrium output is high ... 

This base-catalyzed aldol addition is an equilibrium reaction, and all steps of this reaction are reversible. The free enthalpy of reaction AG " of such aldol reactions is close to zero. In fact, AG° is negative only if there are many H atoms among the substituents R1, R2, and R3 of the two reacting components (structures in Figure 13.44, bottom). Otherwise, the formation of the aldol adduct is endergonic because of the destabilization due to the van-der-Waals repulsion between these substituents. A base-catalyzed aldol addition between two ketones, therefore, is never observed. 

Endergonic reaction (Section 8.3) A reaction for which AG° is positive the products are higher in energy than the reactants and the reaction favors the reactants at equilibrium. 

The A G° for the reaction is +5.7 kcal mofi, whereas the A G in the cell is -0.3 kcal mokf Calculate the ratio of reactants to products under equilibrium and intracellular conditions. Using your results, explain how the reaction can be endergonic under standard conditions and exergonic under intracellular conditions. 

Under standard conditions, A G° = -RT n [product]/[reactants]. Substituting +5.7 kcal mok for A G° and solving for [products]/[reactants] yields 7 x iQ-s. in other words, the forward reaction does not take place to a significant extent. Under intracellular conditions, A G is 0.3 kcal mokk If one uses the equation A G = A G° + i rin [product]/ [reactants] and solves for [products]/[reactants], the ratio is 3.7 x io-5. Thus, a reaction that is endergonic under standard conditions can be converted into an exergonic reaction by maintaining the [products]/[reactants] ratio below the equilibrium value. This conversion is usually attained by using the products in another coupled reaction as soon as they are formed. 

Ion-radical pairs may undergo back electron transfer [22-24] (A et in Scheme 1) in competition with the desired follow-up reactions, which limits the yields of the final reaction products. In fact, the more endergonic the initial electron-transfer step (A et in Scheme 1) is, the faster is the corresponding (exergonic) back electron transfer, which shifts the ET equilibrium towards the precursor (EDA) complex. However, even in such cases of highly endergonic electron transfer, a net chemical reaction to the final products may still be observed if the rate of the follow-up reaction is competitive with that of the back electron transfer in Scheme 1 [25]. This... 

To extend the empirical tunneling expressions above to consider endergonic reactions, we can assume for simplicity that the forward and reverse electron transfers can be related by a temperature and Boltzman constant-dependent equilibrium constant. 

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