Metabolism

1. Metabolism Energy and the Laws of Thermodynamics

2. Metabolism • energy is the ability to do work – including growth and repair, reproduction, finding food, fleeing from predators, learning about biology • organisms continually capture, store, convert and use energy • reactions in cells usually occur in a step-by-step process or metabolic pathways • metabolism is the sum of all the chemical reactions that occur in an organism anabolism + catabolism = metabolism • anabolic pathways require an input of energy • catabolic reactions release energy

3. Forms of Energy • Potential energy • Chemical bonds • Concentration gradients • Electrical potential • Gravitational potential • Kinetic energy • Heat = molecular motion • Mechanical = moving molecules past each other • Electrical = moving charged particles • Falling = gravitational kinetic

4. Forms of Energy • Energy is being converted between these two forms all the time, including during metabolic reactions • Example: burning a marshmallow in a campfire • a marshmallow contains a lot of chemical potential energy (stored in the bonds of the sugar molecules) – during combustion, the sugar molecules are changed into CO2 molecules, and the potential energy is converted into kinetic energy (heat and motion of molecules) as well as light energy • In living cells, much of the work that is done is achieved by transforming chemical or electrical potential energy into kinetic energy • e.g.: molecules highly concentrated on one side of a membrane have potential energy due to the gradient

5. Bond energy • Where does the energy come from when a marshmallow burns? • Bond energy is a measure of the stability of a covalent bond, and there are many covalent bonds in glucose • Measured as the minimum amount of energy required to break (or form) one mole of bonds between two types of atoms • Different bonds have different bond energies • Bond energy is also known as enthalpy(designated as H)

6. Average bond energies

7. Describing energy in a cell • All activities, both in living things and in the non-living world involve changes in energy - thermodynamics is the study of these changes • There are two laws of thermodynamics which apply to a system and its surroundings • System: any defined environment or boundary within which energy transformations are taking place • The Universe, the Earth, an ecosystem, an organism, a cell, or reactant molecules in a test tube

8. First Law of Thermodynamics • There is a finite and constant amount of energy in the universe. Energy cannot be created or destroyed, but can be changed in form. • Also known as the Law of Conservation of Energy • If an object or process gains energy, there must be an object somewhere else in the universe that loses an equivalent amount of energy. • When a chemical reaction occurs and releases energy, some of the energy can be converted into mechanical energy (as in the contraction of muscle fibres), and the rest can be transformed into other forms of energy such as heat

9. Energy Changes in Chemical Reactions • During a chemical reaction, some bonds are broken (in the reactants) and some new bonds are formed (in the products) • Breaking bonds: energy is required because the electrons must be pulled away from the nucleus of atoms • Forming bonds: energy is released because the electrons now orbit at a closer distance to their nuclei • Activation energy (Ea) – this is the minimum amount of energy required to break bonds of the reactants and start a chemical reaction – the Ea is reached at the transition state or when the reactants form an activated complex

10. Potential energy diagrams • To describe the change in energy during a chemical reaction, we can use a potential energy diagram (Ep)

11. Estimating Energy Changes During a Chemical Reaction • We can calculate the energy changes that occur in a chemical change using these simple steps. • Write the balanced chemical equation. • Examine the bond arrangements of the reactants and products. • Calculate the net energy change. Try these steps for the following chemical changes: • Combustion of methane • Decomposition of water • Combustion of butanoic acid

12. Second Law of Thermodynamics • During any energy transfer or process, the universe tends toward an increase in disorder. • No energy conversion is completely efficient - all energy conversions ultimately result in the production of heat • Heat cannot be converted into any useful form – it is considered to be the final state of energy • Consider a car engine – what energy transformations are there?

13. Second Law of Thermodynamics • During any process or transformation of energy, there is an overall increase in disorder • Which is more likely: • A stack of books falls over or a set of books arrange themselves to form a tidy stack? • A car when left alone rusts and falls apart or a set of car parts arrange themselves into a functioning car?

14. Entropy • Entropy is the measure of disorder in a system, or the amount of randomness • e.g.: Your room is likely high in entropy • You could decrease the entropy by putting some energy into cleaning it up and making it less random and more ordered

15. Entropy • Entropy in a chemical system will increase when: • Solid reactants are converted to liquid or gaseous products • Fewer moles of reactant are converted into a greater number of moles of products • Complex reactant molecules react to form a greater number of less complex molecules • Solutes move from an area of high concentration to an area of low concentration until they are distributed evenly

16. Describing the energy of a system • The 2nd law states that all energy transformations are inefficient, so some is lost as heat to the environment. • The portion of energy still available to do work in a system is called free energy (or Gibbs free energy) • Since chemical reactions are breaking bonds and forming new ones, the change in free energy is described by the equation: ∆G = ∆ H - T ∆ S • This can be used to predict the spontaneity of a chemical reaction

17. ∆G = ∆ H - T ∆ S H = enthalpy – the total energy in a system ∆ H = the change in the total energy in a system Endothermic examples (increase in H): • small sugar molecules join together to make disaccharides • amino acids join together to make dipeptides Exothermic examples (decrease in H): • hydrolysis of maltose into two glucose molecules • combustion of the fatty acid butanoic acid

18. ∆G = ∆ H - T ∆ S If ∆H is negative (decrease in bond energy), this means ∆G will most likely be negative. These reactions tend to be spontaneous. If ∆H is positive (increase in bond energy), this means ∆G will be positive. These reactions tend to be non-spontaneous.

19. ∆G = ∆ H - T ∆ S ∆ S = change in entropy If products are more ordered, there is a decrease in entropy. • formation of glucose during photosynthesis • active transport that causes a concentration gradient If products are less ordered, there is an increase in entropy. • combustion of glucose during cellular respiration • diffusion of molecules along a concentration gradient

20. ∆G = ∆ H - T ∆ S If ∆S increases, this decreases ∆G. (Happens spontaneously.) If ∆S decreases, this increases ∆G. (Unlikely to happen without continuous energy input – non-spontaneous)

21. ∆G = ∆ H - T ∆ S T = temperature of the system If reactants have a higher temperature, this makes it more likely that ∆G will be negative. If reactants have a lower temperature, this makes it less likely that ∆G will be negative.

22. Spontaneity in chemical reactions • Sign of ΔG determines reaction direction: • ΔG < 0 (negative ΔG): favourable, spontaneous, exergonic • Process goes left to right, as written. • ΔG > 0 (positive ΔG): unfavourable, not spontaneous, endergonic • Process goes right to left, reverse direction from what is written. • ΔG = 0: Process is AT EQUILIBRIUM • no net reaction in either direction

23. Spontaneity in chemical reactions • If ∆G is positive, the products contain more free energy than the reactants – this is because either the bond energy (H) is higher or the disorder (S) in the system is lower • Endergonic – reactants must absorb free energy from the surroundings • Often anabolic • If ∆G is negative, the products contain less free energy than the reactants - - either bond energy is lower, or disorder is higher or both • Exergonic – these tend to occur spontaneously and release excess free energy • Often catabolic • If ΔG = 0 the process is at equilibirum • no net reaction in either direction

24. J. W. Gibbs

25. Coupling Reactions • In living cells endergonic reactions are not spontaneous but are needed to synthesize any macromolecules • For example, anabolic reactions include creating proteins for a muscle cell, or cellulose fibers for a cell wall – these are highly unlikely to occur on their own – they have a ∆G value which is positive • If they are paired together with exergonic reactions, and the overall ∆G value is negative, then both reactions will tend to occur spontaneously. • This “coupling” of reactions allows many endergonic reactions in a cell to occur – all that is needed is a constant supply of free energy – the molecule that does this is ATP

26. Free energy and enzymes • Use this web applet to help you explain how enzymes are able to increase the rate of reaction for their substrates Free energy and enzymes

27. ATP: a central role in energy cycling + Stored chemical energy is released in catabolic (exergonic) reactions to make ATP ATP is used in energy requiring (endergonic) reactions like muscle movement

28. This bond is easy to break and requires energy! Adenosine triphosphate (ATP) Overall there is a net release of energy H2O Hydrolysis of ATP Formation of these new bonds releases energy H H Inorganic phosphate (Pi) Adenosine diphosphate (ADP)

29. ATP: a central role in metabolism • ATP is notthe highest energy molecule • provides intermediate energy • ATP hydrolysis releases energy – overall it is exergonic • phosphate groups require low energy to break • new bonds formed release more energy than the energy required to break the bond • Phosphorylation by ATP (attaching a phosphate) increases the free energy of other molecules

30. Formation of ATP • Substrate-level phosphorylation • occurs during catalysis of a substrate molecule – energy released is used directly to form high-energy phosphate bonds in ATP • Chemiosmotic synthesis • occurs indirectly, using a concentration gradient across the inner membrane of the mitochondrion and a special enzyme called ATP synthase

31. Use of ATP in cells • Cells use ATP to drive endergonic reactions – but they are spontaneous – their products possess more free energy than their reactants • If the breakage of the high-energy terminal phosphate bond releases more energy than the other reaction consumes, they can be coupled to provide a net release of energy (-∆G)

32. Oxidation – Reduction reactions • Any chemical reaction involves a transfer of electrons • If one molecule gains electrons, another loses them • Reduction: gain of electrons • Oxidation: loss of electrons • LEO the lion says GER • This happens during cell respiration – electrons gained by the oxidation of glucose are transferred to energy carrier molecules and then dropped off at the electron transport chain NADH ⇌ NAD+ + 2e- + H+

33. Energy Carriers • The oxidation of food molecules in cells is controlled – the free energy is released in small steps – rather than all at once like a burning marshmallow • This energy is transferred to energy carrier molecules like the coenzyme nicotinamide adenine dinucleotide (NAD+) • At various points in cellular respiration, enzymes called dehydrogenases remove two hydrogen atoms from a substrate molecule, and transfer two high energy electrons and one of the protons to NAD+ - the other proton is released into the cytoplasm • This reduced form of the energy carrier (or electron shuttle) is NADH and is a very efficient transfer of energy from the food molecule