Chapter 02_2015

transcript

BIOCHEMISTRY I

Midterm exam schedule (Chap 2-7)

4/24 (Fri) 8pm

No withdrawal allowed!

• Key term report submission:

3/4 Chap 2. aqueous system.

3/9. Chap 3. amino acids, peptides, and proteins

3/16. Chap 4. tertiary and quaternary structure of proteins

3/23. Chap 5. protein functions

4/1. Chap 6. enzymes

4/8. Chap 7. carbohydrates and glycobiology

# Quizes

3/23 (Mon): Chapter 2 and 3

4/6 (Mon): Chapter 4 and 5

2 | Water and Aqueous Solutions

CHAPTER 2 Water and Aqueous Solutions

• What kind of interactions occur between molecules

• Why water is a good medium for life

• Why nonpolar moieties aggregate in water

• How dissolved molecules alter properties of water

• How weak acids and bases behave in water

• How buffers work and why we need them

• How water participates in biochemical reactions

Learning goals:

Water is the most abundant substances in

living system, 70% or more of the weight of

most organisms

All aspects of cell structure and function are

adapted to the physical and chemical

properties of water

Hydrogen bonding gives water its unusual properties

Water has higher melting point, boiling point, and heat of

vaporization than most other common solvents

: Hydrogen bonding

FIGURE 2-1a Structure of the water molecule. (a) The dipolar

nature of the H2O molecule; the two hydrogen atoms have localized

partial positive charges (δ+) and the oxygen atom has a partial

negative charge (δ–).

Hydrogen Bonds

• Strong dipole-dipole or charge-dipole interaction that arises

between an acid (proton donor) and a base (proton acceptor)

• Typically 4–6 kJ/mol for bonds with neutral atoms,

and 6–10 kJ/mol for bonds with one charged atom

• Typically involves two electronegative atoms (frequently

nitrogen and oxygen)

• Hydrogen bonds are strongest when the bonded molecules are

oriented to maximize electrostatic interaction

• Ideally the three atoms involved are in a line

Hydrogen Bonding in Water

• Water can serve as both

– an H donor

– an H acceptor

• Up to four H-bonds per water molecule gives water its

– anomalously high boiling point

– anomalously high melting point

– unusually large surface tension

• Hydrogen bonding in water is cooperative

• Hydrogen bonds between neighboring molecules are weak

(20 kJ/mol) relative to the H–O covalent bonds (420 kJ/mol)

Ice: Water in a Solid State

• Water has many different crystal forms;

the hexagonal ice is the most common

•Hexagonal ice forms a regular lattice,

and thus has a low entropy

•Hexagonal ice contains more hydrogen bonds/water molecule

• Thus, ice has lower density than liquid water;

• and, ice floats

Water forms hydrogen bonds with polar solutes

- Water is a polar solvent

: charged or polar compounds can be readily dissolved.

- Chloroform and benzene are nonpolar solvents

- Hydrophilic compounds: dissoves easily in water.

- Hydrophobic compounds: nonpolar molecules

Importance of Hydrogen Bonds

• Source of unique properties of water

• Structure and function of proteins

• Structure and function of DNA

• Structure and function of polysaccharides

• Binding of substrates to enzymes

• Binding of hormones to receptors

• Matching of mRNA and tRNA

Hydrogen Bonds: Examples

Biological Relevance of Hydrogen Bonds

Water dissolves many salts

• High dielectric constant reduces attraction

between oppositely charged ions in salt crystal;

almost no attraction at large (> 40 nm) distances

• Strong electrostatic interactions between the

solvated ions and water molecules lower the

energy of the system

• Entropy increases as ordered crystal lattice is

dissolved

F=Q1Q2

The strength of ionic interaction (F):

Q: magnitude of charge

ε: dielectric constant

r: the distance between charged groups

For water, at 25 °C, ε = 78.5: weak ionic interaction

For benzene, ε = 4.6 : strong ionic interaction

Ionic attractions or repulsion operate only over short distances

The Hydrophobic Effect

• Refers to the association or folding of nonpolar

molecules in the aqueous solution

• Is one of the main factors behind:

– protein folding

– protein-protein association

– formation of lipid micelles

– binding of steroid hormones to their receptors

• Does not arise because of some attractive direct

force between two nonpolar molecules

Solubility of Polar and Nonpolar Solutes

Why are nonpolar molecules poorly soluble in water?

Low solubility of hydrophobic solutes can be explained by entropy

• Bulk water has little order:– high entropy

• Water near a hydrophobic solute is highly ordered:– low entropy

Low entropy is thermodynamically unfavorable, thus hydrophobic solutes have low solubility.

Water surrounding nonpolar solutes has lower entropy

Origin of the Hydrophobic Effect (1)

• Consider amphipathic lipids in water

• Lipid molecules disperse in the solution; nonpolar

tail of each lipid molecule is surrounded by ordered

water molecules

• Entropy of the system decreases

• System is now in an unfavorable state

Origin of the Hydrophobic Effect (2)

• Nonpolar portions of the amphipathic molecule aggregate so that fewer water molecules are ordered

• The released water molecules will be more random and the entropy increases

• All nonpolar groups are sequestered from water, and the released water molecules increase the entropy further

• Only polar “head groups” are exposed and make energetically favorable H-bonds

Hydrophobic effect favors ligand binding

• Binding sites in enzymes and receptors are often hydrophobic

• Such sites can bind hydrophobic substrates and ligands such as steroid hormones

• Many drugs are designed to take advantage of the hydrophobic effect

van der Waals Interactions

• van der Waals interactions have two components:

– Attractive force (London dispersion) depends on the polarizability

– Repulsive force (Steric repulsion) depends on the size of atoms

• Attraction dominates at longer distances (typically 0.4–0.7 nm)

• Repulsion dominates at very short distances

• There is a minimum energy distance (van der Waals contact distance)

Biochemical Significance of van der Waals Interactions

• Weak individually

– easily broken, reversible

• Universal

– occur between any two atoms that are near each other

• Importance

– determines steric complementarity

– stabilizes biological macromolecules (stacking in DNA)

– facilitates binding of polarizable ligands

Examples of Noncovalent Interactions

Name and briefly define four types of

noncovalent interactions that occur between

biological molecules.

(1) Hydrogen bonds: weak electrostatic attractions between one

electronegative atom (such as oxygen or nitrogen) and a

hydrogen atom covalently linked to a second electronegative

(2) electrostatic interactions: relatively weak charge-charge

interactions (attractions of opposite charges, repulsions of

like charges) between two ionized groups

(3) hydrophobic interactions: the forces that tend to bring two

hydrophobic groups together, reducing the total area of the

two groups that is exposed to surrounding molecules of the

polar solvent (water)

(4) van der Waals interactions: weak interactions between the

electric dipoles that two close-spaced atoms induce in each

other.

Weak interactions are crucial to macromolecular structure and function

Hydrogen bond

Ionic interaction

Hydrophobic interaction

Van Der Waals interaction

continually forming and breaking

=> individually insignificant but cumulatively very significant

Effects of Solutes on Properties of Water

• Colligative Properties

– Boiling point, melting point, and osmolarity

– Do not depend on the nature of the solute, just the concentration: number of solute particles

• Noncolligative Properties

– Viscosity, surface tension, taste, and color

– Depend on the chemical nature of the solute

• Cytoplasm of cells are highly concentrated solutions and have high osmotic pressure

Osmotic Pressure

Osmosis => water movement across a semipermeable membrane

driven by differences in osmotic pressure.

Isotonic solution

Hypertonic solution

Hopotonic solution

Effect of Extracellular Osmolarity

Ionization of Water

• O-H bonds are polar and can dissociate heterolytically

• Products are a proton (H+) and a hydroxide ion (OH–)

• Dissociation of water is a rapid reversible process

• Most water molecules remain un-ionized, thus pure water has very low electrical conductivity (resistance: 18 M•cm)

• The equilibrium is strongly to the left

• Extent of dissociation depends on the temperature

H2O H+ + OH-

Proton Hydration

• Protons do not exist free in solution.

• They are immediately hydrated to form hydronium ons.

• A hydronium ion is a water molecule with a proton associated with

one of the non-bonding electron pairs.

• Hydronium ions are solvated by nearby water molecules.

• The covalent and hydrogen bonds are interchangeable. This

allows for an extremely fast mobility of protons in water via

“proton hopping.”

Proton Hopping

Water chain in cytochrome f, which is part of the energy-trapping machinery of

photosynthesis in chloroplasts. Proton hopping is involved in movement of protons.

Ionization of Water: Quantitative Treatment

Concentrations of participating species in an equilibrium process are not independent but are related via the equilibrium constant:

H2O H+ + OH- Keq = ————[H+]•[OH-]

Keq can be determined experimentally, it is 1.8•10–16 M at 25C.[H2O] can be determined from water density, it is 55.5 M.

Ionization of water is expressed by an equilibrium constant

Keq =[H+][OH-]

[55.5 M]at 25°C

( 55.5 M )( Keq ) = [H+][OH-] = Kw

Kw = [H+][OH-] = 1.0 X 10-14 M2

Ion product of water, Kw

Neutral pH:

equal amount of [H+] and [OH-] as in pure water.

[H+] = [OH-] = 10-7 M

What is pH?

• pH is defined as the negative

logarithm of the hydrogen ion

concentration

• Simplifies equations

• The pH and pOH must always

add to 14

• In neutral solution, [H+] = [OH–]

and the pH is 7

• pH can be negative ([H+] = 6 M)

pH = -log[H+]

214- M101]OH][H[ wK

14]OHlog[]Hlog[ -

14pOHpH

pH scale is logarithmic: 1 unit = 10-fold

pH of Some Common Liquids

Dissociation of Weak Electrolytes: Principle

• Weak electrolytes dissociate

only partially in water.

• Extent of dissociation is

determined by the acid

dissociation constant Ka.

• We can calculate the pH if the

Ka is known. But some

algebra is needed!

+ H2O-+ H3O

]OH[ 2 eqa KK

M1074.1COOH]CH[

]COOCH][H[ 5

]COOCH[

]COOHCH[][

pKa measures acidity

• pKa = –log Ka (strong acid large Ka small pKa)

Weak acids have different pKas

Imidazole pKa = 7.0

Protonated? Deprotonated?

At pH = 8.0

At pH = 6.0

pKa=6.0

Buffers are mixtures of weak acids and their anions (conjugate base)

• Buffers resist change in pH

•At pH = pKa, there is a 50:50 mixture of acid and anion forms of the compound

•Buffering capacity of acid/anion system is greatest at pH = pKa

•Buffering capacity is lost when the pH differs from pKa by more than 1 pH unit

Acetic Acid-Acetate as a Buffer System

Henderson–Hasselbalch Equation:Derivation

HA H+ + A-

]A[logppH

]A][H[ -

HA][][H

HA][log-log]log[H- aK

Chapter 2: Summary

In this chapter, we learned about:

•The nature of intermolecular forces

•The properties and structure of liquid water

•The behavior of weak acids and bases in water

Chapter 02_2015

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