introduction, noncovalent bonds, and properties of...

BIOC 460 Summer 2010

Introduction, Noncovalent Bonds, and Properties of Water 1

Introduction, Noncovalent Bonds, and Properties of Water

Reading: Berg, Tymoczko & Stryer: Chapter 1problems in textbook: chapter 1, pp. 23-24, #1,2,3,6,7,8,9, 10,11•practice problems in Gen Chem Review.pdf (see Lecture Notes page)•an elementary explanation of logarithms (Lec. Notes page).

A very well done animation on the life of a cell everyone should watch:http://multimedia.mcb.harvard.edu/anim_innerlife.html

Nelson & Cox, LehningerPrinciples of Biochemistry,Fig. 2-7b

Key Concepts in Biochemistry• Cells -- important structural features; compartments

(plasma membrane, nucleus or nucleoid, cytoplasm, ribosomes, organelles like mitochondria, chloroplasts, endoplasmic reticulum and Golgi apparatus)

• Chemical unity of living systems • Transformation of energy and matter from surroundings -

-> complex, orderly structures• Biomolecules -- functional groups; condensation reactions• Proteins -- molecular workhorses of living systems• Enzymes increase rates of biological reactions to permit life y g p

on a biological timescale.• Rates of processes exquisitely regulated to maintain

dynamic steady state. • 3-D structures of biomolecules determine their functions --

role of noncovalent interactions in structure and function.



Key Concepts, continued• Noncovalent interactions: ionic interactions, hydrogen

bonds, van der Waals interactions, hydrophobic “interactions”

individually much weaker than covalent bonds– individually much weaker than covalent bonds– collectively very strong – crucial to structures and functions of biomolecules

• Properties of water -- “solvent”/milieu for living systems• Most biomolecules have functional groups that are weak

acids or bases– ionization properties of weak acids/bases crucial to

structures and functions of biomolecules– pH determines state of ionization of biomolecular weak

acids and bases.

• Buffers (intracellular and extracellular)

Bacterial CellsREVIEW OF CELL STRUCTURES

Nelson & Cox, Lehninger Principles of Biochemistry, 4th ed., Fig.1-6



Eukaryotic Cells

Nelson & Cox, Lehninger Principles of Biochemistry, 4th ed., Fig.1-7a

Eukaryotic Cells

Nelson & Cox, Lehninger Principles of Biochemistry, 4th ed., Fig.1-7b



Review of Functional Groups in Biomolecules

• Review of Important Functional Groups in Biomolecules



Review of Condensation and Hydrolysis Reactions1. Esters

2. Amides



3. Anhydrides

Biochemistry: the chemical nature of life = the relationship between structure and function!

“You Say You Want a Revolution?”:• Metabolic: 50’s – 80’s via radioisotopes

Structural: 60’s present via X ray crystallography• Structural: 60’s – present via X-ray crystallography and electron microsopy

• Molecular biology: 80’s – present allows for site specific mutagenesis of proteins

• Genomic: 2000 Human Genome solved (24,000 genes for 800,000 – 1,000,000 proteins

• Proteomic: 2000 – present via advances in mass spectrometry and computer technology

Dawning of the Age of Systems Biology (i.e., putting it all together)



As of May 2010:Genomes solved

Protein structures determined: 65,260

Things move fast in this business

The Central Dogma:

DNA: 4 bases (A,T,C,G)

transcription

mRNA translation

Proteins from 20 naturally occurring amino acids

F ti !Function !



Biochemistry is based on four fundamental concepts• Space (structure, proximity)• Time (dynamic interacting processes)• Bonds/Energy (the reversal of disorder): Thermodynamics• Organic chemistry(a bit of inorganic): C, S, O, N, P, Fe, Cu, Zn, Mo

SPACE (1 A {Angstrom} = 10-10 m = 0 1 nm)SPACE (1 A {Angstrom} = 10 10 m = 0.1 nm)• Bonds: 1.5 - 3.0 A• Molecules: 5 - 100 A• Aggregates: 100 - 1000 A• Cells: 104 - 105 A

TIME• Photosynthesis/vision: femto (10-15) to pico (10-12) seconds• Protein motion: nano (10-9) to seconds• Enzyme catalysis: micro (10-6) to seconds• Cell division: 103 to 105 seconds• Your life years

Chemical Bonds/Interactions in Biomolecules• Covalent bonds: single (356 kJ/mol) & double (730)

– 2 atoms share a pair of electrons to fill an orbital on each atom

– Equal or nearly equal sharing --> nonpolar group or molecule

• Examples: C–C and C–H bonds (not polar)– Unequal sharing --> a polar group or molecule

• 1 atom has partial positive charge (δ+)• other atom has partial negative charge (δ )• other atom has partial negative charge (δ–)• Example: an O–H bond (polar)

• Interaction energy (bond energy): the energy releasedduring formation of the bond/interaction (so that much energy would have to be put in to break the bond).



“Weak” individually; collectively quite considerable!

from Nelson & Cox, Lehninger Principles of Biochemistry, 4th ed.)

Ionic Interactions (charge-charge interactions =salt bridges): electrostatic attraction or repulsion between charged groups (-6 kJ/mol in H2O; -231 in hexane)

Coulomb’s Law:

E = Energy of interactionq’s = charges (+/-, +/+, or -/-)D = dielectric constant (1 for vacuum, ~2 for hexane,D dielectric constant (1 for vacuum, 2 for hexane,

~80 for H2O)r = distance between charged atomsk = proportionality constant; value depends on units

desired for expressing energy



• H-bond Donor H-bond Acceptor (4 – 20 kJ/mol)–N---H ----------> : N–O---H ----------> : O

• Electrostatic in nature.• Effect of polarity of solvent on dielectric constant (D).

B i ifi t h H b d i b i d i l t i

Hydrogen Bonds

Becomes significant when H-bond is buried in non-polar protein interior.

• Strict distance dependence: 1.5 – 2.6 A (H O,N)2.4 - 3.5 A (N O)

Beyond these distances, H-bonds do not form.

H-bonds cont’d

The energy of the H-bond is very dependent on orientation of the three atoms involved.



• van der Waals Interactions (2 – 4 kJ/mol)– Energy of a van der Waals interaction as 2 atoms approach one

another within about 4-5 Å– Individually very weak and nonspecific, but sum of many can be very

important in steric (shape) complementarity– Energy drops off by 1/r6 (very dramatic)

Berg et al., Fig. 1.10



Bonding CharacteristicsEnergy (kJ/mol) Distance (A)

C l t 356 (single) 1 54 C CCovalent 356 (single)730 (double)

1.54 C-C1.34 C=C

Ionic -6 (@3.4 A;H2O)-231 (3.4 A;hexane)

• E drops off 1/r2

Variable

H-bond* 4 – 20 1.5 – 2.6 H->O2 4 3 5 N >O2.4 – 3.5 N->O

van der Waals* 2 - 4 per atom pair• E drops off 1/r6

4 - 5

* Individually very weak contribution, collectively enormous!

ThermodynamicsBiology depends on the laws of thermodynamics: the creation of order (life) from disorder (the universe).

1. First Law: The total energy of the SYSTEM and its SURROUNDINGS are CONSTANT.

2. Second Law: The total entropy of the SYSTEM and its SURROUNDINGS always increases for a spontaneousprocess. Increased entropy = death in the bio-world. Thus, nature favors disorder, life is the result of order.

∆G ∆H T∆S RT l K∆G = ∆H – T∆S = -RT lnKeq

“At Equilibrium” is a very important concept!



Hydrophobic "Interactions"• hydrophobic effect, the "oil drop" effect• Nonpolar molecules (the system) cannot interact with H2O (the

surroundings) by H-bonds or ionic interactions.• Nonpolar molecule “forces” water to organize itself, decreasing entropy of

H2O.• In order to increase total entropy water “forces” nonpolar molecules toIn order to increase total entropy , water forces nonpolar molecules to

interact with each other (decreasing system’s entropy but increasing surrounding’s entropy).

• Results in "hydrophobic” or nonpolar interactions, which are entropicallydriven.

Berg et al.Fig. 1-12

Properties of Water• Polarity

– asymmetric charge distribution makes molecule dipolar (polar)O atom ™ H atom ™+– O atom ™–, H atom ™+

– strong ionic character to O-H bond• Hydrogen bonding

– Water molecule bent:– In how many hydrogen bonds can

1 H2O molecule participate?2 p p



H2O H-bonding in ice

Nelson & Cox, Lehninger Principles ofBiochemistry, 4th ed., Fig. 2-2

Solvent Properties of Water • Excellent solvent for:

– Ions/charged groups

– Neutral polar compounds, e.g. sugars:highly solvated (H-bonds to solvent HOH)



Solvent Properties of Water, continued• Poor solvent for hydrophobic groups - fatty acid alkyl “tail”

Nelson & Cox, LehningerPrinciples of Biochemistry, 4th ed., Fig. 2-7a

• Fatty acid: example of an amphipathic (amphiphilic) molecule

Gen Chem Review (Yes! You have to know how to use that stuff)• Review from general chemistry -- see Gen Chem review

material for details. – Review notes posted, will be covered in review session

tomorrowtomorrow.– General chemistry is a prerequisite for biochem, and

you need to understand it -- review it on your own and/or come to review session.

• General concepts of chemical equilibrium and equilibrium constants

• Importance of H+ (proton) concentration in cells and in p (p )extracellular media

• State of ionization of weakly acidic groups on biomolecules important to structure and function

• Titration curves• Buffers



Ionization Properties of H2O• H2O and acids in aqueous solution dissociate to yield protons

(H+) (hydrated, forming H3O+ etc.)• Proton concentrations often expressed on log10 scale as pH:

pH = –log[H+]• Tendency of Bronsted acid to donate proton to H2O

(dissociate the proton) is described by its equilibrium acid dissoc. constant Ka, i.e. by its pKa = –log Ka.

Note: p(anything) = –log (anything)• pKa values measured experimentally by titration curves as

the pH at half equivalence points• Relationship between pH, pKa, and ratio of conjugate

base/conjugate acid described by the Henderson-Hasselbalch Equation:

Buffers• Homeostasis: maintenance of constant conditions in internal environment• Fluids in living systems -- pH is regulated, almost constant• pH regulated by buffer systems

– Buffer: aqueous system that resists changes in pH when small amounts of acid or base are addedamounts of acid or base are added

– Buffer system: aqueous solution of a weak acid and its conjugate base– Buffer range of a weak acid: pH values near its pKa, about ±1 pH unit

from pKa (Maximum buffering capacity is at the pKa.)• Equilibrium acid dissocation reaction (remember Le Chatelier’s Principle, the

“law of mass action”):HA <==> H+ + A–

Th hi h th [H+] (th l th H) th ilib hift t l ft– The higher the [H+] (the lower the pH), the more equilib. shifts to left.– The lower the [H+] (the higher the pH), the more equilib. shifts to right.– Exact ratio of base/acid (A–/HA) depends on Henderson-Hasselbalch Eq:

– When pH = pKa [A–]/[HA] = __________?



2 Physiologically Important Buffer Systems• Intracellular: Phosphate species

– Inorganic phosphate (phosphoric acid)

– Organic phosphates, e.g., phosphomonoesters

• Extracellular (blood plasma of mammals):carbonic acid / bicarbonate buffer system

pKa~6.1



1. Physiologically, how would a mammal deal withacidosis (blood pH ↓, [H+] ↑) in the short term?

2. Physiologically, how would a mammal deal with alkalosis (blood pH ↑; [H+] ↓) in the short term?

Learning Objectives and Review Material• Review basic structures of cells and organelles -- important

structural features and compartments (nucleus or nucleoid, plasma membrane, cytoplasm, ribosomes, mitochondria, chloroplasts, and endoplasmic reticulum).

• Review (from posted lecture notes here) functional groupsimportant in biomolecules, and condensation reactions involvingimportant in biomolecules, and condensation reactions involving some of these functional groups.

• List and explain the characteristics of 4 types of noncovalent interactions important in structures and interactions of biomolecules. Answer the following questions:

1. What is an ionic interaction (charge-charge interaction), and what other terms are used to describe the same thing?• How does the distance between two charged groups affect the

energy of their interaction?energy of their interaction? • What are the relative values of the dielectric constants for a

nonpolar solvent and a polar solvent?• How does solvent polarity affect strength of ionic interactions?• What type of solvent is water? • Is an ionic interaction stronger in a polar solvent or in a

nonpolar solvent?



Learning Objectives and Review, continued(Noncovalent interactions, continued)

2. What is a hydrogen bond, what is a hydrogen bond donor, and what is a hydrogen bond acceptor?• How does the strength of a hydrogen bond relate to its

directionality?B bl t id tif h i l ( d th ifi t• Be able to identify chemical groups (and the specific atoms involved) that can serve as hydrogen bond donors and groups which can serve as hydrogen bond acceptors. [Do not confuse a hydrogen bond donor with a proton donor (Bronsted acid).]

3. What are van der Waals interactions?• How (qualitatively, not an equation) does their strength relate

to the distance between atoms?• Why are such weak, nonspecific interactions important in y p p

biochemistry?4. What is the “hydrophobic effect”?

• Explain the idea of ”hydrophobic interactions" and the roles they play in biological systems. (Roles will become more apparent as the semester progresses).

Learning Objectives and Review, continued• Explain the properties of H2O (its ionization properties, polarity,

hydrogen bonding ability, and solvent properties) that are so important to its role as the major constituent of living systems.

– Explain: titration curve, buffer, and pKa. Relate the strength of a weak acid to its pKa.Write out the 3 acid dissociation reactions of phosphoric acid and– Write out the 3 acid dissociation reactions of phosphoric acid, and write out condensation reactions showing formation of a phosphomonoester and of a phosphodiester.

– See practice problems at end of Gen Chem Review notes:

Explain relationships between (and be able to do calculations involving these relationships): 1. [H+] and pH2. Ka (acid dissociation constant) and pKa3. ratio of [conjugate base]/[conjugate acid] and pH and pKa4. ratio of [conjugate base/[conjugate acid] and fraction or

percent of a functional group that's in the form of the conjugate acid or the conjugate base.

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