1. Themodynamics

2. Metabolism = ‘change’ Refers to all the chemical reactions that change or transform matter and energy in cells Metabolic Pathway = a sequential series of chemical reactions in living cells energy in Anabolic Pathway

3. How can we define energy Energy = the capacity to do work Kinetic – energy due to movement (e.g. flowing water) Potential – stored energy (e.g. A boulder perched on top of a hill) Chemical energy is potential energy stored up in the bonds of a molecule.

4. Types of Energy Mechanical energy is the energy found in the motion of objects. Chemical is the energy contained within the bonds that hold atoms together to form molecules Radiant the source of energy for all living things, derived from the sun All these types of energy may be found in actualized form such as kinetic energy of a falling stone, or in potential form energy stored in objects at rest.

5. Thermodynamics – the science that studies the transfer and transformation of thermal energy (heat)

6. First Law of Thermodynamics Take a guess – this law concerns the amount of energy in the universe The total amount of energy in the universe is constant. Energy cannot be created or destroyed, but it can be transformed from one type into another and transferred from one object to another.

7. Second Law of Thermodynamics During any process, the universe tends toward disorder (i.e., entrophy increases). entropy – a measure of randomness / disorder greater entropy = greater disorder metabolise glucose into CO 2 and H 2 O C 6 H 12 O 6 + O 2  CO 2 + H 2 O + energy

8. Thermodynamics In living systems, the primary potential energy source is chemical energy Bond energy is a measure of the stability of the covalent bond between atoms and is measured in kilojoules(kJ) It is the amount of energy required to break one mole of bonds between two atoms and is also equal to the amount of energy released when that bond is formed. E.g. the C – C bond requires 346kJ/mol. to break it apart, C = O bond requires 799kJ/mol to break

9. Bond Energies Measured in kJ / mol (kilojoules per mole) Double bonds require more energy to break than single bonds The greater the bond energy, the more stable the bond. Bond TypeAverage Bond Energy (kJ/mol) H-H436 C-H411 O-H459 N-H391 C-C346 C-O359 C=O799 O=O494

10. Bond Energies and Thermodynamics During chemical reactions, the reactants molecules must first be broken apart and then the product molecules must be formed. When energy or heat is applied to the reactant molecule, the bonds will absorb the energy until they weaken and break apart. If a chemical reaction releases more heat than it uses, it is called an exothermic reaction If more heat is absorbed into the system than is actually released at the end of the reaction, it is referred to as an endothermic reaction.

11. Bond Energies exothermic reactions – energy released endothermic reactions – energy absorbed energy is required to break bonds energy is released to form bonds breaking > forming breaking < forming endothermic exothermic

12. Potential Energy Diagram EXOTHERMIC REACTION

13. Potential Energy Diagram

14. Gibbs Free Energy We can measure the amount of energy actually available to break and subsequently form other chemical bonds = free energy of the molecule Free energy  G) = energy that can do work Note: heat is “useless” as it dissipates  G = G final – G initial

15. Gibbs Free Energy  G = G final – G initial Change in free energy can be used to predict whether a chemical reaction is spontaneous or not When:  G is negative – reaction can proceed spontaneously = exergonic reaction, products contain less free energy (lower bond energies) than reactants Greater disorder

16. Gibbs Free Energy  G = G final – G initial Change in free energy can be used to predict whether a chemical reaction is spontaneous or not When:  G is positive – reaction is non-spontaneous Endergonic rx - requires an input of energy Products contain more free energy (higher bond energies) than reactants Lower disorder in system

17. Gibbs Free Energy Exergonic reaction – spontaneous, -  G Endergonic reaction – not spontaneous, +  G C 6 H 12 O 6 + 6O 2  6CO 2 + 6H 2 O  G = -2870 kJ/mol 6CO 2 + 6H 2 O  C 6 H 12 O 6 + 6O 2  G = +2870 kJ/mol

18. Potential Energy Diagram GG

19. Equilibrium Equilibrium reactions convert back and forth with minimal energy. For equilibrium reactions:  G = 0

21. Driving Endergonic Reactions Couple to an exergonic reaction Example: 1. Flow of a solute down its concentration gradient 2. ATP hydrolysis 3. Exergonic redox reactions Exergonic reaction provides the energy required for the endergonic reaction

22. Adenosine Triphosphate ATP is the primary free energy source for cells ATPADP + P i + energy H 2 O  G = -31 kJ/mol in the cell,  G is closer to -54 kJ/mol ATPase

23. ATP Molecule Free energy is used to do 3 main kinds of work 1. Mechanical work such as the contraction of muscles cells, the flow of cytoplasm or the movement of chromosomes during cell reproduction 2. Transport work such as the pumping of substances across membranes against the concentration gradient 3. Chemical work such as synthesizing complex molecules from simpler atoms The immediate source of energy that powers these cellular processes is ATP

24. ATP – adenosine triphosphate Composed of the nitrogen base adenine bonded to the 5-carbon sugar, ribose which is in turn bonded to 3- phosphate groups The phosphate tail is unstable, and the bonds between the last phosphate groups can be broken When the bond is broken, a molecule of inorganic phosphate is removed from ATP and it becomes ADP

25. Adenosine Triphosphate (ATP)

26. The ATP Molecule

27. The ATP/ADP Cycle