09/09/09GMS BI 555/755 Lecture 1.1 BIOCHEMISTRY A GMS BI 555/755 Course manager: Konstantin Kandror...

09/09/09 GMS BI 555/755 Lecture 1. 1

BIOCHEMISTRY A GMS BI 555/755Course 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)

09/09/09 GMS BI 555/755 Lecture 1. 2

GMS BI 555/755: BiochemistryResources

• Instructor for lectures 1-3: Joseph Zaia, X8-6762, [email protected]

• 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

09/09/09 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.gov1. Bioenergetics

a. First Law of Thermodynamicsb. Second law of Thermodynamicsc. Free energyd. Flow of energy in living systems

2. Non-covalent interactionsa. Covalent bondsb. Hydrogen bondingc. Ionic interactionsd. Van der Waal’s forcese. Hydrophobic effectf. Protein specificity

3. Weak acids and bases in biological systems.

09/09/09 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.”

09/09/09 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 ifit can exchange energy andmatter with its surroundings,closed if it can exchangeenergy but not matter, andisolated if it can exchangeneither energy nor matter.

Bioenergetics

09/09/09 GMS BI 555/755 Lecture 1. 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.

Bioenergetics

09/09/09 GMS BI 555/755 Lecture 1. 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.

Bioenergetics

09/09/09 GMS BI 555/755 Lecture 1. 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.

Bioenergetics

09/09/09 GMS BI 555/755 Lecture 1. 9

J. Willard GibbsChemical Physicist

Gibbs free energy (Gibbs’ Law, 1870’s):ΔG = ΔH-TΔSH = enthalpy (chemical bond energy)S = entropy (randomness)ΔG <0 for spontaneous processesΔG >0 reverse reaction spontaneousΔG = 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

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).

Bioenergetics

Thermodynamically, the amount of energy capable of doing work during a chemical reaction is measured quantitatively by the change in the Gibbs free energy.

09/09/09 GMS BI 555/755 Lecture 1. 10

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

ΔG = ΔH - TΔS

Bioenergetics

09/09/09 GMS BI 555/755 Lecture 1. 11

ΔG = ΔH-TΔS

Mixing of two solutes by removing a barrierΔHΔSΔG

2H2 + O2 → 2H2O

ΔHΔSΔG

09/09/09 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

H2O1 M glucose

H2O H2O0.5 M glucose

H2O0.5 M glucose

ΔH = 0ΔS = >0

Entropically favorable because glucose is allowed to diffuse over the maximum available volume achieving an increase in randomness

nutrients

waste

nutrients

waste

ΔH = 0ΔS = <0

Entropically unfavorable: increase in order required to maintain homeostasis of nutrients inside cell and waste outside cell.

A cell

Bioenergetics

09/09/09 GMS BI 555/755 Lecture 1. 13

Formation of the DNA double helix

CG

AT

TA

AT

GC

TAA

TTA

CG

AT

TA

AT

GC

TAA

TT

A

ΔS <0 (more ordered)ΔH <0 (heat released)ΔG <0 (spontaneous process, heat

released through formation of multiple H-bonds)

Bioenergetics

• 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.

09/09/09 GMS BI 555/755 Lecture 1. 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.

Bioenergetics

09/09/09 GMS BI 555/755 Lecture 1. 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

Gproducts

Greactants

Activation energy

ΔG

09/09/09 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

Molecular Cell Biology, Lodish et al.

09/09/09 GMS BI 555/755 Lecture 1. 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°′.

Molecular Cell Biology, Lodish et al.

Lodish et al Molecular Cell Biology chap 15.

Ion transport using Na+ gradients

09/09/09 18GMS BI 555/755 Lecture 1.

09/09/09 GMS BI 555/755 Lecture 1. 19

Flow of energy in living systems

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].)

Bioenergetics

(controlled oxidation and reduction)

09/09/09 GMS BI 555/755 Lecture 1. 20

The energy Required to Break Some Important Covalent Bonds Found in Biological Molecules

Type of Bond Energy (kcal/mol) Type of Bond Energy (kcal/mol)

SINGLE BOND DOUBLE BOND

O—H 110 C=O 170

H—H 104 C=N 147

P—O 100 C=C 146

C—H 99 P=O 120

C—O 84

C—C 83 TRIPLE BOND

S—H 81 C≡O 195

C—N 70

C—S 62

N—O 53

S—S 51

* Note that double and triple bonds are stronger than single bonds.

09/09/09 GMS BI 555/755 Lecture 1. 21

Chemical Bonds: Covalent

Types of covalent bonds in biochemistry: proteins

H2N

HN

NH

OH

R1

O R2

O R3

O

H2N

H2N

H2N

OH

R1

O R2

O R3

O

OH

OH

Amino acids

Peptide: amide linked

Non-covalent bonding

Covalent bonding involves sharing of a pair of electrons in the form of orbital overlap

09/09/09 GMS BI 555/755 Lecture 1. 22

O

OH

HOOH

O

OH

O

HONHAc

O

OH

O

OH

OHO

OH

O

HONHAc

O

OH

O

OH

OHO

OH

O

HONHAc

OH

OH

HO O

HONHAc

OH

OHO

OH

HOOH

OH

OH

Gal GlcNac

Polylactosamine

Chemical Bonds: Covalent

Types of covalent bonds in biochemistry: carbohydrates

Monosaccharides(polyhydroxy aldehydes/ketones)

Glycosidic bond: Ether linkage


09/09/09 GMS BI 555/755 Lecture 1. 23

Non-covalent interactions

Hydrogen bonding:H

O

HH

O

H

H

O

H

H

O

H

+ 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 25C 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

09/09/09 GMS BI 555/755 Lecture 1. 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 biomoleculesAcceptors: O: and N: lone pairs. Numerous interactions possible in biomolecules.


09/09/09 GMS BI 555/755 Lecture 1. 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

09/09/09 GMS BI 555/755 Lecture 1. 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.

09/09/09 GMS BI 555/755 Lecture 1. 27

Ionic interactions

• 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)

K+▪▪▪▪▪▪▪▪▪▪Oδ-

Hδ+

Hδ+

E = (kq1q2)/Dr2 (Coulomb’s law)k = proportionality constantD = dielectric constant = 80 for waterr = distance

q1 q2r

H2Osolvation Oδ-

Hδ+

Hδ+

Cl-

(KCL lattice)

09/09/09 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.

09/09/09 GMS BI 555/755 Lecture 1. 29

Ionic interactions: H-bonds revisited

δ-

δ+ δ-

δ+

• 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

09/09/09 GMS BI 555/755 Lecture 1. 30

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.

δ+ δ+δ-δ-


09/09/09 GMS BI 555/755 Lecture 1. 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


09/09/09 GMS BI 555/755 Lecture 1. 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.


09/09/09 GMS BI 555/755 Lecture 1. 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.

09/09/09 GMS BI 555/755 Lecture 1. 34

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.


A spectrum of macromolecular non-covalent interactions (binding partners) gives rise to biological activity

09/09/09 GMS BI 555/755 Lecture 1. 35

Ionization of water

Water dissociates into hydronium (H3O+) and hydroxyl (OH-) ions. For simplicity, we refer to the hydronium ion

as a hydrogen ion (H+) and write the equilibrium as

The equilibrium constant Keq of this dissociation is given by

in which the terms in brackets denote molar concentrations. Because the concentration of water (55.5 M) is changed little by ionization, expression 1 can be simplified to give

in which Kw is the ion product of water. At 25°C, Kw is 1.0 × 10-14.

Note that the concentrations of H+ and OH- are reciprocally related. If the concentration of H+ is high, then the concentration of OH- must be low, and vice versa. For example, if [H+] = 10-2 M, then [OH-] = 10-12 M

09/09/09 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].

09/09/09 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).

09/09/09 GMS BI 555/755 Lecture 1. 38

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.

Buffers: weak acids that resist pH changes near their pKa

09/09/09 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)

09/09/09 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.

NH

N

N

N

O

NH2

pKa = 9.7N

N

N

N

O

NH2

09/09/09 GMS BI 555/755 Lecture 1. 41

Proteins in solution

pKa ~2 pKa ~9

09/09/09GMS BI 555/755 Lecture 1.1 BIOCHEMISTRY A GMS BI 555/755 Course manager: Konstantin Kandror...

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