Cells get energy from the following reaction: ATP ⇌ ADP + Pi ΔG° = -33 kJ/mol at 37 °C Calculate the value of the equilibrium constant for the reaction at this temperature.
Added by Sebastian P.
Step 1
Given: Delta G = -33 kJ/mol Convert -33 kJ/mol to Joules/mol: -33 kJ/mol * 1000 J/1 kJ = -33,000 J/mol Show more…
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Having the reactions at 37°C: 1) A + B ⇌ Z ΔGº = 23.8 kJ mol-1 2) ATP ⇌ ADP + phosphate ΔGº = -31.0 kJ mol-1 Calculate the equilibrium constant for the coupled reaction: A + B + ATP ⇌ Z + ADP + phosphate
Adi S.
ATP + H2O → ADP + Pi ΔG°' = -30.5 kJ mol-1 The final ΔG°' for the coupled reaction is -30.5 kJ mol-1. The reaction that is coupled is: A + C ↔ B + M + R ΔG°' = 30.5 kJ mol-1 U + RK ↔ C ΔG°' = 61 kJ mol-1 KU + R ↔ C + A ΔG°' = -61 kJ mol-1 BA + C ↔ K + M ΔG°' = -30.5 kJ mol-1
How many moles of ATP must be converted to ADP by the reaction $\mathrm{ATP}(a q)+\mathrm{H}_{2} \mathrm{O} \longrightarrow \mathrm{ADP}(a q)+\mathrm{HPO}_{4}^{2-}(a q)+2 \mathrm{H}^{+}(a q) \quad \Delta G^{\circ}=-31 \mathrm{~kJ}$ to bring about a nonspontaneous biochemical reaction in which $\Delta G^{\circ}=+372 \mathrm{~kJ} ?$
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