BIOCHEMISTRY A GMS BI 555/755 Course manager: Konstantin Kandror

1. BIOCHEMISTRY A GMS BI 555/755 Course manager: Konstantin Kandror See Courseinfo (www.bumc.bu.edu)Every Monday and Wednesday, 10:00 to 12:00 Room R103 Every other Friday, 11:30 to 12:30 Room R103 (Computer room L1110 on September 11) GMS BI 555/755 Lecture 1.

2. GMS BI 555/755: BiochemistryResources • Instructor for lectures 1-3: Joseph Zaia, X8-6762, jzaia@bu.edu • Course notes, print them out BEFORE the lecture. • Biochemistry, 6th ed. Berg, Tymoczko, Stryer (hard copy) • Engines, Energy and Entropy John Fenn. • National Center for Biotechnology Information (on-line text books) • www.ncbi.nih.gov • Stryer 5th ed • Molecular Cell Biology Lodish et al. • Molecular Biology of the Cell, 4th ed, Alberts et al. • Many other texts • Wikipedia for some basic concepts • Discussion groups GMS BI 555/755 Lecture 1.

3. GMS BI 555/755 Lecture 1: . Bioenergetics: Molecular Interactions and Thermodynamics within Biological systems • Readings: Biochemistry, Berg et al, 6th ed, pp 6-17. • Supplementary: Molecular Cell Biology, Lodish, el al, Chap. 2, “Chemical Foundations” on-line at www.ncbi.nih.gov • Bioenergetics a. First Law of Thermodynamics b. Second law of Thermodynamics c. Free energy d. Flow of energy in living systems • Non-covalent interactions a. Covalent bonds b. Hydrogen bonding c. Ionic interactions d. Van der Waal’s forces e. Hydrophobic effect f. Protein specificity • Weak acids and bases in biological systems. GMS BI 555/755 Lecture 1.

4. The cell as a system: what principles underlie it’s organization and energetics and, by extension, life itself? Sir Arthur Eddington in The Nature of the Physical World: “The second law of thermodynamics holds, I think, the supreme position among the laws of Nature. If someone points out to you that your pet theory of the universe is in disagreement with Maxwell’s equations - then so much the worse for Maxwell’s equations. If it is found to be contradicted by observation, well, these experimentalists do bungle things sometimes. But if your theory is found to be against the second law of thermodynamics I can give you no hope; there is nothing for it but to collapse in deepest humiliation.” GMS BI 555/755 Lecture 1.

5. The Universe as a thermodynamic system: energy/matter is neither created nor destroyed but converts between different forms. • System: matter within a defined region of space A system is open if it can exchange energy and matter with its surroundings, closed if it can exchange energy but not matter, and isolated if it can exchange neither energy nor matter. GMS BI 555/755 Lecture 1. Bioenergetics

6. Types of energy: Potential and Kinetic • Kinetic energy: • Thermal energy - energy of motion, heat energy. Thermal regulation in warm blooded animals used for maintenance of ideal temperature for enzymatic reactions; cells do not use heat to do work (exception of thermophilic bacteria) • Radiant energy – energy of light (photons) may be converted to heat energy. May be used to generate chemical bond energy in photosynthesis. • Potential energy: stored energy • Energy stored in chemical bonds: (i.e. C6H12O6,)used to drive metabolism • Energy of concentration gradients sequestered by biological membranes. Ions, nutrients, waste, protons, etc. maintained in compartment specific concentrations. • Electric potential: all cells maintain a gradient of electrical charge across their plasma membrane. GMS BI 555/755 Lecture 1. Bioenergetics

7. Energy as work: directed energy that has a focused effect. • Lifting a weight • Transporting a nutrient across a membrane • Synthesizing a protein Energy as heat: transferred to surroundings to cause random molecular motions due to increase in temperature. Dissipation. GMS BI 555/755 Lecture 1. Bioenergetics

8. Biochemical thermodynamics • 1st Law: The amount of energy in a system plus surroundings is constant. • Energy cannot be created or destroyed but can be converted between different forms. Mass is a type of energy. • Any energy released in the formation of chemical bonds must be used to break other chemical bonds, converted to heat or light, or stored in another form • 2nd Law: The amount of entropy (randomness) in a system plus surroundings always increases. • The universe is expanding. The amount of disorder in the universe resulting from a (bio) chemical process always increases. • In order to create biological order, an organism must expend energy, but the net entropy in the universe increases. GMS BI 555/755 Lecture 1. Bioenergetics

9. Metabolism: An organism must continuously maintain order with respect to its environment. It must create order within itself at the expense of that of its environment (waste). It must free itself, through metabolism, from the disorder (decay) it must create in order to exist (1930s→1950s). Gibbs free energy (Gibbs’ Law, 1870’s): ΔG = ΔH-TΔS H = enthalpy (chemical bond energy) S = entropy (randomness) ΔG <0 for spontaneous processes ΔG >0 reverse reaction spontaneous ΔG = 0 both forward and reverse reactions occur at equal rates • Gibbs free energy explains the driving force behind (biological and non-biological) chemical change. • ΔS proportional to heat transferred from system to surroundings • ΔS inversely proportional to temperature J. Willard Gibbs Chemical Physicist Thermodynamically, the amount of energy capable of doing work during a chemical reaction is measured quantitatively by the change in the Gibbs free energy. GMS BI 555/755 Lecture 1. Bioenergetics

10. ΔG = ΔH - TΔS The total entropy change is given by the expression Substituting equation 1 into equation 2 yields Multiplying by -T gives The function -TΔS has units of energy and is referred to as free energy or Gibbs free energy, after Josiah Willard Gibbs, who developed this function in 1878: The free-energy change, ΔG, will be used throughout to describe the energetics of biochemical reactions. Recall that the Second Law of Thermodynamics states that, for a reaction to be spontaneous, the entropy of the universe must increase. Examination of equation 3 shows that the total entropy will increase if and only if Rearranging gives TΔSsystem > ΔH, or entropy will increase if and only if GMS BI 555/755 Lecture 1. Bioenergetics

11. ΔG = ΔH-TΔS Mixing of two solutes by removing a barrier ΔH ΔS ΔG 2H2 + O2 → 2H2O ΔH ΔS ΔG GMS BI 555/755 Lecture 1.

12. ΔG = ΔH-TΔS Entropy (S): randomness or disorder in a system. Entropy increases as a system becomes more disordered and decreases as it becomes more ordered. Chambers separated by permeable membrane ΔH = 0 ΔS = >0 Entropically favorable because glucose is allowed to diffuse over the maximum available volume achieving an increase in randomness H2O H2O 0.5 M glucose H2O 1 M glucose H2O 0.5 M glucose Entropically unfavorable: increase in order required to maintain homeostasis of nutrients inside cell and waste outside cell. nutrients A cell ΔH = 0 ΔS = <0 nutrients waste waste GMS BI 555/755 Lecture 1. Bioenergetics

13. Formation of the DNA double helix GCTAATTA CGATTAAT GCTAATTA CGATTAAT • An exothermic reaction (ΔH < 0) that increases entropy (ΔS > 0) occurs spontaneously (ΔG < 0). • An endothermic reaction (ΔH > 0) will occur spontaneously if ΔS increases enough so that the T ΔS term can overcome the positive ΔH. • If the conversion of reactants into products results in no change in free energy (ΔG = 0), then the system is at equilibrium; that is, any conversion of reactants to products is balanced by an equal conversion of products to reactants. ΔS <0 (more ordered) ΔH <0 (heat released) ΔG <0 (spontaneous process, heat released through formation of multiple H-bonds) GMS BI 555/755 Lecture 1. Bioenergetics

14. ΔG = ΔH-TΔS • Enthalpy ΔH = negative when the reactants have more chemical bond energy than the products. Excess energy is given off, usually as heat. • Combustion is an exothermic process: • 2C8H18 + 25O2→ 16CO2 + 18H2O + energy • Metabolism is an exothermic process: • C6H12O6 + 6O2→ 6CO2 + 6H2O + energy • ΔH = positive when the products have more chemical bond energy than the reactants. When enthalpy is positive, energy must be added to chemical bonds for the reaction to occur. GMS BI 555/755 Lecture 1. Bioenergetics

15. What does ΔG tell us about a biochemical reaction? • A reaction occurs spontaneously only if ΔG <0 (exothermic or exergonic) • A system is at equilibrium when ΔG = 0 (steady state) • A reaction in which ΔG>0 requires input of energy (endothermic or endergonic) • ΔG = Gproducts – Greacants, it does not depend on the reaction path (i.e. combustion vs metabolism) • ΔG provides no information on the reaction rate Activation energy ΔG Greactants Gproducts GMS BI 555/755 Lecture 1.

16. Parameters affecting ΔG of a reaction: temperature, pressure, concentrations of reactants and products Most biological reactions differ from standard conditions, particularly in the concentrations of reactants. However, we can estimate free-energy changes for different temperatures and initial concentrations, using the equation where R is the gas constant of 1.987 cal/(degree · mol), T is the temperature (in degrees Kelvin), and Q is the initial ratio of products to reactants. For the interconversion of glyceraldehyde 3-phosphate (G3P) and dihydroxyacetone phosphate (DHAP) Q = [DHAP]/[G3P] and ΔG°′ = −1840 cal/mol. Equation 2-8 for ΔG then becomes • ΔG may be calculated for any set of concentrations of DHAP and G3P. • If the initial concentrations of both DHAP and G3P are 1 M, then ΔG = ΔG°′ =  −1840 cal/mol, because RT ln 1 = 0. • If [DHAP] = 0.1 M and [G3P] = 0.001 M, then Q = 0.1/0.001 = 100, and GMS BI 555/755 Lecture 1. Molecular Cell Biology, Lodish et al.

17. An Unfavorable Chemical Reaction Can Proceed If It Is Coupled with an Energetically Favorable Reaction Many chemical reactions in cells are energetically unfavorable (ΔG > 0) and will not proceed spontaneously. One example is the synthesis of small peptides (e.g., glycylalanine) or proteins from amino acids. Cells are able to carry out a reaction that has a positive ΔG by coupling it to a reaction that has a negative ΔG of larger magnitude, so that the sum of the two reactions has a negative ΔG. Suppose that the reaction has a ΔG°′ of +5 kcal/mol and that the reaction has a ΔG°′ of −10 kcal/mol. In the absence of the second reaction, there would be much more A than B at equilibrium. The occurrence of the second process, by which X becomes Y + Z, changes that outcome: because it is such a favorable reaction, it will pull the first process toward the formation of B and the consumption of A. The ΔG°′ of the overall reaction will be the sum of the ΔG°′ values of each of the two partial reactions: The overall reaction releases energy. In cells, energetically unfavorable reactions of the type A ⇌ B + X are often coupled to the hydrolysis of the compound adenosine triphosphate (ATP), a reaction with a negative change in free energy (ΔG°′ = −7.3 kcal/mol), so that the overall reaction has a negative ΔG°′. GMS BI 555/755 Lecture 1. Molecular Cell Biology, Lodish et al.

18. Ion transport using Na+ gradients Lodish et al Molecular Cell Biology chap 15. GMS BI 555/755 Lecture 1.

19. Flow of energy in living systems (controlled oxidation and reduction) Diagram demonstrating the flow of energy in living organisms. Arrows point in the direction in which energy flows. We focus only on the most common processes and do not include less ubiquitous ones, such as bioluminescence. (Adapted from D.A. Harris, Bioenergetics at a glance, Blackwell Science, Oxford [1995].) GMS BI 555/755 Lecture 1. Bioenergetics

20. The energy Required to Break Some Important Covalent Bonds Found in Biological Molecules GMS BI 555/755 Lecture 1.

21. Chemical Bonds: Covalent Types of covalent bonds in biochemistry: proteins Amino acids Covalent bonding involves sharing of a pair of electrons in the form of orbital overlap Peptide: amide linked GMS BI 555/755 Lecture 1. Non-covalent bonding

22. Chemical Bonds: Covalent Types of covalent bonds in biochemistry: carbohydrates Monosaccharides (polyhydroxy aldehydes/ketones) Glycosidic bond: Ether linkage GMS BI 555/755 Lecture 1. Non-covalent bonding

23. Non-covalent interactions Hydrogen bonding: + 5 kcal/mol • H-bond formation releases a small amount of energy relative to the typical energies of a covalent bond. • Average kinetic energies of molecules at 25C is 0.6 kcal/mol • H-bonds are weak and transient in nature • In combination, H-bond interactions become significant for holding molecules together non-covalently • H-bonds underly the unique chemical properties of water (solvent of life). • Much higher boiling point than similarly sized non-hydrogen bonding molecules • Water is very dense; solid state has lower density than liquid state (ice floats) • Polar protic nature enabled evolution of biomolecules Structure of ice, showing H-bonds GMS BI 555/755 Lecture 1. Non-covalent bonding

24. Non-covalent interactions Hydrogen bonding: general form 0.26-0.31 nm (twice covalent bond length) Donors, N-H and O-H bonds in biomolecules Acceptors: O: and N: lone pairs. Numerous interactions possible in biomolecules. GMS BI 555/755 Lecture 1. Non-covalent bonding

25. Hydrogen bond formation in proteins and nucleic acids releases 1-2 kcal/mol, much lower than for bond formation among water molecules. The reason for this is that in order for a new protein-water H-bond to form, an H-bond with water must break. Water-free microenvironments enable formation of H-bonds between polar groups GMS BI 555/755 Lecture 1. Non-covalent bonding

26. Ionic interactions Electronegativity values of main-group elements in the periodic table. Atoms located to the upper right tend to have high electronegativity, fluorine being the most electronegative. Elements with low electronegativity values, such as the metals lithium, sodium, and potassium, are often called electropositive. The electronegativities of several atoms abundant in biological molecules differ enough that they form polar covalent bonds (e.g., O—H, N—H) or ionic bonds (e.g., Na+Cl−). Because the inert gases (He, Ne, etc.) have complete outer shells of electrons, they neither attract nor donate electrons, rarely form covalent bonds, and have no electronegativity values. GMS BI 555/755 Lecture 1.

27. q1 q2 r Hδ+ Oδ- Hδ+ Ionic interactions E = (kq1q2)/Dr2 (Coulomb’s law) k = proportionality constant D = dielectric constant = 80 for water r = distance • Atoms with large electronegativity differences do not share electrons by forming covalent bonds. Rather, electrons are completely transferred, forming a positive and a negative ion. X▪ + Y▪ → X+ + Y- Na▪ + Cl▪ → Na+ + Cl- • Many ionic compounds are readily soluble in water because of release of energy when they become hydrated (energy of hydration); they become surrounded by a shell of water molecules (hydration sphere) Hδ+ H2O solvation Cl- K+▪▪▪▪▪▪▪▪▪▪Oδ- Hδ+ (KCL lattice) GMS BI 555/755 Lecture 1.

28. Ionic interactions • Energetics of solvation of ionic compounds is a balance between the energy of the ionic lattice versus those of the hydrated ions Na▪ + Cl▪ → Na+ + Cl- Readily soluble, energy of solvation greater than lattice energy Ca3(PO4)2→ Insoluble due to high lattice energy • Ions must lose their hydration shell as they pass through ion transport channels in biological membranes. The energetics of these processes influences ion selective transport. • (Mg2+), six water molecules held tightly in place by electrostatic interactions between the two positive charges on the ion and the partial negative charge on the oxygen of each water molecule. GMS BI 555/755 Lecture 1. Non-covalent bonding

29. Ionic interactions: H-bonds revisited • Polarity and H-bonding capability (donating and accepting) of water makes it an excellent solvent for polar molecules • It has a high (80) dielectric constant (D) and weakens electrostatic interactions and H-bonds between other molecules by competing for interactions • Water free microenvironments exist in biological systems to enable strong interactions between polar molecules δ- δ+ δ- δ+ GMS BI 555/755 Lecture 1.

30. δ- δ- δ+ δ+ Van der Waal’s forces (aka London dispersion forces) • Two atoms approaching one another induce transient dipoles resulting in a weak, non-specific attractive force. • Occur in all molecules, polar and non-polar • Responsible for cohesion of non-polar liquids (alkanes). Stronger interactions (ionic, H-bonding) override van der Waal’s attractions. • Van der Waal’s contact: when the repulsive forces of a pair of atom electron clouds is balanced by van der Waal’s attractive forces. GMS BI 555/755 Lecture 1. Non-covalent bonding

31. Van der Waal’s forces • Van der Waal’s energy is ~1kcal/mol, weaker than H-bonds, only slightly higher than thermal energy • Van der Waal’s forces become significant for large molecules in which domains with complementary shapes allow for many contacts • Thus they are important driving forces behind the energetics of intra- and inter-molecular protein interactions: • Protein folding • Multi-protein complexes (molecular machines) • Enzyme-substrate binding GMS BI 555/755 Lecture 1. Non-covalent bonding

32. Hydrophobic interactions • Non polar molecules lack polar and/or ionic groups and do not form ionic or H-bonding interactions with water molecules. • Water must form an ordered cage surrounding a hydrophobic molecule. Thus there is an energetic cost to the dissolution of a hydrophobic molecule in water causing hydrophobic molecules to remain associated with one another (a droplet) or with a surface (vessel wall) rather than go into solution. • Hydrophobic interactions are thus driven by entropic effects. There is a strong tendency of hydrophobic molecules to interact with one another rather than become solvated in water. GMS BI 555/755 Lecture 1. Non-covalent bonding

33. Non-covalent interactions give rise to protein specificity The binding of a hypothetical pair of proteins by two ionic bonds, one hydrogen bond, and a large combination of hydrophobic and van der Waals interactions. The structural complementarity of the surfaces of the two molecules gives rise to this particular combination of weak bonds and hence to the specificity of binding between the molecules. GMS BI 555/755 Lecture 1.

34. A spectrum of macromolecular non-covalent interactions (binding partners) gives rise to biological activity Multiple weak bonds stabilize specific associations between large molecules.(Left) In this hypothetical complex, seven noncovalent bonds bind the two protein molecules A and B together, forming a stable complex. (Right) Because only four noncovalent bonds can form between proteins A and C, this interaction may be too weak for the A-C complex to exist in cells. GMS BI 555/755 Lecture 1. Non-covalent bonding

35. Ionization of water GMS BI 555/755 Lecture 1.

36. Definition of pH and pK The pH of a solution is a measure of its concentration of H+. The pH is defined as The ionization equilibrium of a weak acid is given by The apparent equilibrium constant Ka for this ionization is The pKa of an acid is defined as Inspection of equation 4 shows that the pKa of an acid is the pH at which it is half dissociated, when [A-]=[HA]. GMS BI 555/755 Lecture 1.

37. The Henderson-Hasselbalch Equation Relates pH and Keq of an Acid-Base System Many molecules used by cells have multiple acidic or basic groups, each of which can release or take up a proton. In the laboratory, it is often essential to know the precise state of dissociation of each of these groups at various pH values. The dissociation of an acid group HA, such as acetic acid (CH3COOH), is described by The equilibrium constant Ka for this reaction is By taking the logarithm of both sides and rearranging the result, we can derive a very useful relation between the equilibrium constant and pH as follows: or Substituting pH for −log [H+] and pKa for −log Ka, we have Henderson-Hasselbalch equation: pKa of any acid is equal to the pH at which half the molecules are dissociated and half are neutral (undissociated). GMS BI 555/755 Lecture 1.

38. Buffers: weak acids that resist pH changes near their pKa The titration curve of acetic acid (CH3COOH). The pKa for the dissociation of acetic acid to hydrogen and acetate ions is 4.75. At this pH, half the acid molecules are dissociated. Because pH is measured on a logarithmic scale, the solution changes from 91 percent CH3COOH at pH 3.75 to 9 percent CH3COOH at pH 5.75. The acid has maximum buffering capacity in this pH range. GMS BI 555/755 Lecture 1.

39. The titration curve of phosphoric acid (H3PO4). This biologically ubiquitous molecule has three hydrogen atoms that dissociate at different pH values; thus, phosphoric acid has three pKa values, as noted on the graph. The shaded areas denote the pH ranges — within one pH unit of the three pKa values — where the buffering capacity of phosphoric acid is maximum. In these regions the addition of acid (or base) will cause the least change in the pH. Phosphodiester bond (pKa ~ 3) GMS BI 555/755 Lecture 1.

40. pH mediated DNA denaturation • As pH is increased, the extent of deprotonatoin of the N-1 nitrogen atom of Guanine increases. At pH > 9.7, the N-1 proton is no longer there to participate in base pairing through H-bonding • At pH <5, some H-bond acceptors become protonated and unable to participate in H-bonding, disrupting the structure of DNA. GMS BI 555/755 Lecture 1.

41. Proteins in solution pKa ~2 pKa ~9 GMS BI 555/755 Lecture 1.