BCH 5045 Graduate Survey of Biochemistry - Environmental ...hort.ifas.ufl.edu/faculty/guy/bch5045/Lecture Files/Lecture 5.pdf · Bioenergetics and Biochemical Reaction Types

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BCH 5045

Graduate Survey of Biochemistry

Instructor: Charles GuyProducer: Ron ThomasDirector: Glen Graham

Lecture 5Slide sets available at:

http://hort.ifas.ufl.edu/teach/guyweb/bch5045/index.html

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• LEHNINGER• PRINCIPLES OF BIOCHEMISTRY

• Fifth Edition

David L. Nelson and Michael M. Cox

© 2008 W. H. Freeman and Company

CHAPTER 13Bioenergetics and Biochemical

Reaction Types



Gibbs equation

∆G = ∆H – T∆S

Define free-energy, G; entropy, S and H, enthalpy

For A + B C + D, then

Keq = [C][D]/[A][B], thus

∆G = ∆G° + RTln[C][D]/[A][B] or ∆G° = -RTlnKeq

State the 1st and 2nd

Laws of Thermodynamics



The chemistry of ATP and other nucleoside triphosphates is key and centrally important to their function and tightly related to the nature of the phosphoanhydridebond between the phosphate groups.

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Presentation Notes

FIGURE 1–25 Adenosine triphosphate (ATP) provides energy. Here, each P represents a phosphoryl group. The removal of the terminal phosphoryl group (shaded pink) of ATP, by breakage of a phosphoanhydride bond to generate adenosine diphosphate (ADP) and inorganic phosphate ion (HPO42–), is highly exergonic, and this reaction is coupled to many endergonic reactions in the cell (as in the example in Fig. 1–26b). ATP also provides energy for many cellular processes by undergoing cleavage that releases the two terminal phosphates as inorganic pyrophosphate (H2P2O72– ), often abbreviated PPi.



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FIGURE 1–26a Energy coupling in mechanical and chemical processes. (a) The downward motion of an object releases potential energy that can do mechanical work. The potential energy made available by spontaneous downward motion, an exergonic process (pink), can be coupled to the endergonic upward movement of another object (blue).



Requires energy input Can provide energy A coupled reaction that is energetically favorable with a product that can provide an accessible store of energy

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FIGURE 1–26b Energy coupling in mechanical and chemical processes. (b) In reaction 1, the formation of glucose 6-phosphate from glucose and inorganic phosphate (Pi) yields a product of higher energy than the two reactants. For this endergonic reaction, G is positive. In reaction 2, the exergonic breakdown of adenosine triphosphate (ATP) has a large, negative free-energy change (G2). The third reaction is the sum of reactions 1 and 2, and the free-energy change, G3, is the arithmetic sum of G1 and G2. Because G3 is negative, the overall reaction is exergonic and proceeds spontaneously.



Typical model reaction free energy transformation coordinate

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FIGURE 1–27 Energy changes during a chemical reaction. An activation barrier, representing the transition state (see Chapter 6), must be overcome in the conversion of reactants (A) into products (B), even though the products are more stable than the reactants, as indicated by a large, negative free-energy change (G). The energy required to overcome the activation barrier is the activation energy (G‡). Enzymes catalyze reactions by lowering the activation barrier. They bind the transition-state intermediates tightly, and the binding energy of this interaction effectively reduces the activation energy from G‡uncat (blue curve) to G‡cat (red curve). (Note that activation energy is not related to free-energy change, G.)



A stylized depiction of the flow of energy in humans. The ultimate source of the energy we need and use is from the sun and was transformed into chemical forms by photosynthetic organisms. The catabolic processes provide chemical forms of energy to be consumed in anabolic processes.



Chemical forms that have bond structures that are energy enriched and in which can be made available to do work.





This table shows the inextricable linkage between Keq and standard free energy of chemical reactions. Both indicate how far a reaction will proceed to reach equilibrium.





High energy content metabolites central to cellular function. All contain phosphorus, even CoA.





ATP and nucleoside triphosphates are versatile and offer three high energy linkages that can be exploited in chemical reactions.

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Presentation Notes

FIGURE 13-20 Nucleophilic displacement reactions of ATP. Any of the three P atoms (α, β, or γ) may serve as the electrophilic target for nucleophilic attack—in this case, by the labeled nucleophile R—18O:. The nucleophile may be an alcohol (ROH), a carboxyl group (RCOO–), or a phosphoanhydride (a nucleoside mono- or diphosphate, for example). (a) When the oxygen of the nucleophile attacks the γ position, the bridge oxygen of the product is labeled, indicating that the group transferred from ATP is a phosphoryl (—PO32–), not a phosphate (—OPO32–). (b) Attack on the β position displaces AMP and leads to the transfer of a pyrophosphoryl (not pyrophosphate) group to the nucleophile. (c) Attack on the α position displaces PPi and transfers the adenylyl group to the nucleophile.



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FIGURE 13-11 Chemical basis for the large free-energy change associated with ATP hydrolysis. 1 The charge separation that results from hydrolysis relieves electrostatic repulsion among the four negative charges on ATP. 2 The product inorganic phosphate (Pi) is stabilized by formation of a resonance hybrid, in which each of the four phosphorusﾐoxygen bonds has the same degree of double-bond character and the hydrogen ion is not permanently associated with any one of the oxygens. (Some degree of resonance stabilization also occurs in phosphates involved in ester or anhydride linkages, but fewer resonance forms are possible than for Pi.) 3 The product ADP2– immediately ionizes, releasing a proton into a medium of very low [H+] (pH 7). A fourth factor (not shown) that favors ATP hydrolysis is the greater degree of solvation (hydration) of the products Pi and ADP relative to ATP, which further stabilizes the products relative to the reactants.





Reactions with strongly negative ΔG´° values.

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FIGURE 13-13 Hydrolysis of phosphoenolpyruvate (PEP). Catalyzed by pyruvate kinase, this reaction is followed by spontaneous tautomerization of the product, pyruvate. Tautomerization is not possible in PEP, and thus the products of hydrolysis are stabilized relative to the reactants. Resonance stabilization of Pi also occurs, as shown in Figure 13-11. FIGURE 13-14 Hydrolysis of 1,3-bisphosphoglycerate. The direct product of hydrolysis is 3-phosphoglyceric acid, with an undissociated carboxylic acid group, but dissociation occurs immediately. This ionization and the resonance structures it makes possible stabilize the product relative to the reactants. Resonance stabilization of Pi further contributes to the negative free-energy change. FIGURE 13-15 Hydrolysis of phosphocreatine. Breakage of the P—N bond in phosphocreatine produces creatine, which is stabilized by formation of a resonance hybrid. The other product, Pi, is also resonance stabilized. FIGURE 13-16 Hydrolysis of acetyl-coenzyme A. Acetyl-CoA is a thioester with a large, negative, standard free energy of hydrolysis. Thioesters contain a sulfur atom in the position occupied by an oxygen atom in oxygen esters. The complete structure of coenzyme A (CoA, or CoASH) is shown in Figure 8-38.



What is free energy and why is this important?



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FIGURE 13-17 Free energy of hydrolysis for thioesters and oxygen esters. The products of both types of hydrolysis reaction have about the same free-energy content (G), but the thioester has a higher free-energy content than the oxygen ester. Orbital overlap between the O and C atoms allows resonance stabilization in oxygen esters; orbital overlap between S and C atoms is poorer and provides little resonance stabilization.



What is going on here?

ATP

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FIGURE 13-18 ATP hydrolysis in two steps. (a) The contribution of ATP to a reaction is often shown as a single step, but is almost always a twostep process. (b) Shown here is the reaction catalyzed by ATP-dependent glutamine synthetase. 1 A phosphoryl group is transferred from ATP to glutamate, then 2 the phosphoryl group is displaced by NH3 and released as Pi.




• Fifth Edition



CHAPTER 2

The Foundations of Biochemistry



WaterWater is pure, Water is natural, Water is healthy, Water can help all

Water is simple, Water is free, Water can help the lives, The lives of you and me

Olivia Taylor



Water is the solvent of life

It is required for the function of all active cells and biological entities associated with cells. The properties of water are so very extremely profound in their biological significance. The ultimate structure and shape of many of the compounds on which life is based, proteins, nucleic acids, lipids, and carbohydrates result from their direct interaction with water.

The physical properties of water are so unusual relative to other compounds of similar composition and size that it is reasonable to assume that where there is no water there can be no life.



Water is present in high concentration (~55.6 M), its shape is triangular, and it has a polar character. Having polarity means that one side of the molecule is charged differently than the opposite side. The asymmetrical charge distribution in water is caused by the propensity of the oxygen nucleus to draw the electrons shared with the two hydrogen atoms closer to it and away from the hydrogen.

A polar compound is said to be a dipole. Water is a dipole of 1.84 debye units or a dielectric constant of 80.2 at 20°C (1 debye = 0.21 Å in charge separation).

The highly polar nature of water makes it an extremely good solvent for compounds with polar character.



In fact, as the dielectric constant of a medium increases its ability to keep opposite charged molecules apart increases according to Coulomb's Law

1. F = kq1q2Dr2

where F is the force between electrical charges, q1 and q2 separated by a distance r. D is the dielectric constant of the medium (solvent) and k is the proportionality constant (8.99 X 109).

The high dielectric constant of water means that the force between two electrical charges will be reduced.

Thus NaCl readily dissolves in water.



Water molecules can interact with charged components of compounds forming hydration shells that neutralize charge and helps to keep oppositely charged atoms from coalescing into an insoluble salt.

Many of the compounds that make up cells are polar, but some are nonpolar. These compounds do not mix well or dissolve in water. An example of this is the mixing or nonmixing of oil with water (the salad dressing dilemma). In this case, the theme is "like dissolves like."

This leads to an important paradigm of hydrophilic and hydrophobic interactions in biochemical structure and function.

Nonpolar compounds in an aqueous environment tend to associate together excluding water from the interior of the conglomeration and this is called hydrophobic attraction.



In the case of a fatty acid suspended in water, micelle formation will occur. In reality, there is no actual attraction between nonpolar molecules that causes them to associate in this manner.

When a nonpolar compound is surrounded by water the normal positioning of the water molecules in the vicinity of that compound is disrupted. This disruption of the normal positioning and free movement of water molecules leads to the formation of an ordered arrangement of a large number of water molecules into a cage like structure. Such structures are called clathrates. Highly ordered clathrate structures are energetically not favored.

The coalescence of individual nonpolar molecules together is more energetically favorable. Consequently, the energy of the system is minimized as the nonpolar molecules coalesce out of contact with water.



Given the polarity of the water molecule, electrons are unequally distributed and this gives rise to an attraction between a hydrogen atom of one water molecule with the electrons of an oxygen atom of another water molecule.

Bond angles of water are similar to those expected of a tetrahedral and suggest that the polar nature of the orbitals is equivalent to sp3 hybrids.

Water is not linear. The two hydrogen atoms occupy the corners of the tetrahedron and the nonbonding pairs of oxygen’s electrons occupy the other corners.

The electron orbital structure of the oxygen in a water molecule is such that two hydrogen atoms from two different water molecules can interact with the oxygen of another water molecule.



This linear attraction between the oxygen and the hydrogen atoms of different water molecules is none other than a hydrogen bond (H-bond). Actually, the hydrogen atom can form a H-bond with an N or O atom that is weakly basic and can act as an acceptor. This is made possible by the hydrogen atom's small size, which permits it to approach a lone-pair electron cloud of acceptor atoms.

The nucleus of H is the only one small enough to approach within the van der Waals distance and get close enough for an electrostatic interaction to exist.

The H-bond distance of water is about 0.18 nm.

Please view this video: http://www.youtube.com/watch?v=LGwyBeuVjhU&feature=related

http://www.youtube.com/watch?v=LGwyBeuVjhU&feature=related�



This capacity to form H-bonds gives water its uncommon physical properties. Although weak compared to covalent bonds (20 kj/mol vs ca 150-500 kj/mol), the enormous number of H-bonds (4-5 kcal/mol; 1 cal = 4.18 joules) that exist at any given moment in water are central to its physical properties.

The turnover rate of a H-bond in water is extremely rapid at 10-11/sec.

Each water molecule can form as many as four H-bonds.

Liquid water possesses about 85% of the maximal H-bonds possible. In contrast, ice contains roughly 99% of the possible H-bonds.

Hydrogen bonding makes liquid water molecules extremely cohesive.



The constant translational and rotational movements of water molecules in solution are surprisingly not as well understood as one may assume, yet are believed to underlie water’s unique physical properties and its central role in biological systems.

Our understanding of bulk water’s structure and dynamics despite intensive study remains limited.

Recent experimental advances are helping to reveal the “dance” of water molecular dynamics when interacting with simple molecules and ions.

Some of the latest experiments use "ultrashort" infrared pulses with durations on the order of 50 x 10–15 s, or 50 fs. By using such short pulses it becomes possible to resolve some of the elementary “dance” steps.



An infrared frequency is chosen to excite stretching vibrations of the OH bonds.

The vibrational frequency of the OH group shifts when the H atom changes its hydrogen-bond environment. The ultrashort infrared pulses can be used to measure time-dependent fluctuations in the vibrational frequency in a process called "spectral diffusion," where therotational dynamics of the OH bond is measured by using polarized light.

In one example, it has been shown that a water molecule shifts its donated H-bonds betweenwater and perchlorate (anion) acceptors by a large, angular rotation of 49 ± 4°. http://www.youtube.com/watch?v=D3mbOzOYzQ0

As we better understand water’s unique molecular properties, we will better understand why it is the solvent of life.

http://www.youtube.com/watch?v=D3mbOzOYzQ0�



The unusual physical properties of water include:

a high surface tension relative to NH3, creating a high cohesiveness. This cohesiveness is so important to plants because it permits the capillary movement of water from the roots to the canopy. Without the high cohesiveness of water, trees and certainly trees like the Redwoods could not exist;

a high viscosity;

a high heat capacity 1 cal/g/°C between 14.5° and 15.5°C; a high boiling point; otherwise would boil at -100°C without H-bonds;

Physical Property H2O NH3 CH4 H2S Molecular Weight 18 17 16 34 Boiling Point (C) 100 -33 -161 -61

Freezing Point (C) 0 -78 -183 -86 Viscosity (centipoise) 1.01 0.25 0.10 0.15



a high heat of fusion 80 cal/g or 6 kj/mol which decreases the H-bonds from 99% down to 85%;

and a high heat of vaporization 540 cal/g or 41 kj/mol which decreases H-bonds from 85% down to near 0%. This is significant because vaporization consumes much heat and provides significant cooling capacity.

Water can act as a Bronsted-Lowry acid:

H2O + NH3 ↔ OH‾ + NH4+

acid base base acid

or as a base:

H2O + HCO3‾ ↔ H3O+ + CO32-

base acid acid base



However, absolutely pure water has a neutral pH! So, what is the pH of water that has reached equilibrium with the ambient atmosphere?

Why is this becoming an increasingly worrisome problem?




• Fifth Edition



CHAPTER 2

The Foundations of Biochemistry



Figure 2-1 Structure of the water molecule

Page

39

From Biochemistryby Voet & Voet, Third edition, 2004



Figure 2-2 Hydrogen bond between two water molecules.

Page

40




Figure 2-3 Structure of ice.

Page

40




Figure 2-5 Solvation of ions by oriented water molecules.

Page

42

From Biochemistry by Voet & Voet, Third edition, 2004



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Presentation Notes

FIGURE 2-7a Amphipathic compounds in aqueous solution. (a) Long-chain fatty acids have very hydrophobic alkyl chains, each of which is surrounded by a layer of highly ordered water molecules.



Water is not like any of the other compounds listed above. It has a high melting point, a high boiling point and a high heat of vaporization.



The water cages are know as clathrates, and because they are highly ordered, they represent one form of energy described in the Gibbs equation.



The following slides are included but were not shown in the lecture 5 video.



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Presentation Notes

FIGURE 2-1 Structure of the water molecule. (a) The dipolar nature of the H2O molecule is shown in a ball-and-stick model; the dashed lines represent the nonbonding orbitals. There is a nearly tetrahedral arrangement of the outer-shell electron pairs around the oxygen atom; the two hydrogen atoms have localized partial positive charges (δ+) and the oxygen atom has a partial negative charge (δ–). (b) Two H2O molecules joined by a hydrogen bond (designated here, and throughout this book, by three blue lines) between the oxygen atom of the upper molecule and a hydrogen atom of the lower one. Hydrogen bonds are longer and weaker than covalent O—H bonds.



Diagram of the structure of ice.

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FIGURE 2-2 Hydrogen bonding in ice. In ice, each water molecule forms four hydrogen bonds, the maximum possible for a water molecule, creating a regular crystal lattice. By contrast, in liquid water at room temperature and atmospheric pressure, each water molecule hydrogen-bonds with an average of 3.4 other water molecules. This crystal lattice structure makes ice less dense than liquid water, and thus ice floats on liquid water.



Hydrogen bonding

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FIGURE 2-3 Common hydrogen bonds in biological systems. The hydrogen acceptor is usually oxygen or nitrogen; the hydrogen donor is another electronegative atom. FIGURE 2-4 (part 4) Some biologically important hydrogen bonds.



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FIGURE 2-5 Directionality of the hydrogen bond. The attraction between the partial electric charges (see Figure 2-1) is greatest when the three atoms involved in the bond (in this case O, H, and O) lie in a straight line. When the hydrogen-bonded moieties are structurally constrained (when they are parts of a single protein molecule, for example), this ideal geometry may not be possible and the resulting hydrogen bond is weaker.



Note: What does amphipathic mean?



Polar molecules are hydrophilic, meaning they love water, and in the case shown here water can dissociate the Na and Cl from the crystal lattice.

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FIGURE 2-6 Water as solvent. Water dissolves many crystalline salts by hydrating their component ions. The NaCl crystal lattice is disrupted as water molecules cluster about the C– and Na+ ions. The ionic charges are partially neutralized, and the electrostatic attractions necessary for lattice formation are weakened.



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BCH 5045 Graduate Survey of Biochemistry - Environmental ...hort.ifas.ufl.edu/faculty/guy/bch5045/Lecture Files/Lecture 5.pdf · Bioenergetics and Biochemical Reaction Types