Thermochemistry: Biology-Related Examples

Thermochemistry: Biology-Related Examples See Lectures 16-18 for thermodynamics. See Lecture 18 notes for ATP coupled reactions and enthalpies of hydrogen bonds. Recall from Lecture 17 that the ∆H° of ATP hydrolysis is -24 kJ/mol of ATP. NH2 NH2 N C N O O C O HC C O P O P O P O CH2 N O O HC CH N O O O N CH HC OH CH O + 2H2O (l) (aq) P O O P O CH2 N O O O C C N HC HC CH HC OH C CH N +HPO4–2 (aq) + H3O+ CH OH OH ATP (aq) ADP (aq) ∆H° = -24 kJ/mol ATP Example 1) from page 3 of Lecture 18 notes: ATP-coupled reactions in the body Many biological reactions are non-spontaneous, meaning they require energy to proceed in the forward direction. The hydrolysis of ATP (ATP → ADP), a spontaneous process, can be coupled to a nonspontaneous reaction to drive the non-spontaneous reaction forward. ATP ADP The resulting ∆Gº of the coupled reaction is the sum of the individual ∆Gº values. First, let’s calculate the ∆Gº for ATP hydrolysis at 310 K (body temperature) given that ∆H° = -24 kJ/mol (from Lecture #17) and ∆S° = +22 J/K•mol. ∆Gº = ∆Hº - T∆Sº ∆Gº = -24 kJ/mol – 310 K (0.022 kJ/mol) ∆Gº = -31 kJ/mol of ATP hydrolyzed Note: the calculated free energies are under standard conditions. This is an approximation since these molecules are NOT under standard conditions in cells. Consider the following ATP-coupled reaction: the conversion of glucose to glucose-6-P. Why this reaction is important: Adding a phosphate (P) group to glucose gives the glucose a negative charge, which prevents the glucose molecule from diffusing back out of the cell through the “greasy” cell membrane. - O O P O OH ATP H2C HO C H CH CH H C HO - O ADP H2C HO CH O CH HC OH HO OH C H H C O HC OH OH ∆Gº = +17 kJ/mol for glucose to glucose-6-P ∆Gº = -31 kJ/mol for ATP hydrolysis An enzyme couples the glucose to glucose-6-P reaction to ATP hydrolysis. The net change in free energy = 17 kJ/mol + (-31 kJ/mol) = -14 kJ/mol If ATP hydrolysis is spontaneous, why is it not occurring unregulated in the cell? KINETICS! A reaction can be thermodynamically spontaneous, but kinetically very very slow. Example 2) from page 5 of Lecture 18 notes: Enthalpy and hydrogen bonds A hydrogen bond is an electrostatic interaction between a hydrogen atom in a polar bond (typically a H-F, H-O or H-N bond) and a “hydrogen-bond acceptor”, a strongly electronegative atom. X H δ+ Y where X = O, N, F And Y is the hydrogen bond acceptor: N, O, or F δ- The H-bond acceptor (Y) atom must be small, highly electronegative atom with a lone pair of electrons available for bonding. For example, hydrogen bonds form between water molecules: O H H H O H Mean bond enthalpies of hydrogen-bonds (H-bonds): H-bonds are the strongest type of intermolecular interaction. However, H-bonds are weaker than covalent or ionic bonds. mean bond enthalpy (in kJ/mol) OH---O H-bond 20 H-O covalent bond 463 OH---N H-bond 29 NH---N H-bond 14 H-N covalent bond 388 Hydrogen bonding can be intermolecular (as in the water molecules above) or intramolecular. Intramolecular hydrogen bonds in proteins are required for a protein’s 3-dimensional shape. Hydrogen bonds are also important in DNA structure and function. Hydrogen bonding in DNA Hydrogen bonding binds together complementary strands of DNA to form a double helix. H HC N H O N H N N N C C C C N N H N N H O HC CH C C Guanine CH N C C N CH C N O H H CH3 C C C N N CH N Cytosine H 3 H-bonds between C and G Adenine O Thymine 2 H-bonds between A and T The lower bond enthalpies of hydrogen bonds compared to covalent bonds facilitate the separation of DNA strands during DNA replication.