BIOCHEMISTRY I

Water

Lecture 03

9:30 AM-10:45 AM

Outline of chapter

Hydrogen bonding in water

Some biologically important hydrogen bonds

Gibbs free energy and competition between entropy

and enthalpy

Some examples of polar, non-polar and amphipathic

compounds

water as solvent

amphipaths suspended in water

role of water in formation of enzyme-substrate

complexes (desolvation)

proton hopping

bond strengths

interactions in the aqueous environment

Ionization of water, weak acids, and weak bases

Henderson-Hasselbalch equation (pH, Pka and Buffer

concentration)

Titration of biological weak acids

Enzymatic activity and pH

buffering in the blood

Chapter 2: Water 4

Hydrogen bonding in waterRoughly tetrahedral geometry

Dipole moment

Oxygen is more electronegative and has two partial negative charges (-)

Each hydrogen nucleus (proton) has partial positive charge (+)

(a) Ball and stick model and (b) space filling model

(c) Two water models joined by a hydrogen bond (three lines)

Hydrogen bonds longer (~2-3x) and weaker (~15-20x) than covalent bond

Chapter 2: Water 5

--Waters capacity to form multiple hydrogen bonds with other water molecules alter its chemical properties with respect to other similarly sized molecules

Water Methanol EthanolBoiling point 373K 338K 352K

Heat of vaporization

40.7 kJ/mol 35.3 kJ/mol 38.6 kJ/molViscocity 1mPS 0.5mPS 1.1mPS

Surface tension 72dyn/cm 23dyn/cm 22dyn/cmDielectric constant 80.1 32.7 24.5

Chapter 2: Water 6

Hydrogen bonding in water

--Water at 0C (ice) forms H-bonds with 4 other water molecules (donates two protons and oxygen accepts two protons)

--Water at room temperature has an average of 3.4 H-bonds per water molecule (donates two protons and oxygen accepts 1.4 protons)

HYDROGEN BONDING IN BIOLOGICAL SYSTEMS

Hydrogen bonding is not unique to water, viz.:

In fact: H-bonds can form between an electronegative atom and an H-donor moiety. The latter is an H atom bonded to another electronegative atom, typically N, O, or F .; i.e., the

ensemble consists of H sandwiched between two electronegative atoms

Examples of H-bond structures ubiquitous in biological systems:

Directionality of the Hydrogen Bond

electrostatic interactions maximized when acceptor atom and H atom are in a straight line

directionality contributes, i.e., to the ability of proteins and nucleic acids (with many such bonds) to form precise 3D structures

Intermolecular vs Intramolecular Hydrogen bonding

Chapter 2: Water 9

Rudimentary Thermodynamics

G = H - TS

Reactants Products

Keq = [Products] / [Reactants]

G = -RT ln (Keq)

G is the Gibbs free energy:

If G = (-) then Keq is large and Products are favored

If G = (-) then Keq is small and reactants are favored

H is change in enthalpy (internal energy of system). H = (-) means that the system is more stable

S is change in entropy (disorder)

Water as a solvent

Chapter 2: Water 9

Water is a polar solvent and dissolves most biological molecules that are hydrophilic (water-loving).

Hydrophobic (water-fearing) and non-polar molecules like lipids and waxes dissolve poorly in water.

Water can dissolve salts like NaCl (below) by hydrating and stabilizing the ions. This weakens the electrostatic interactions of the ions acting against the ions tendency to form a crystalline lattice.

Water as a solvent

random orientation

Chapter 2: Water 10

Amphipathic compounds contain polar (or charged) regions and nonpolar regions.

Suspension of amphipaths in water

The hydrophobic alkyl groups cluster together to present the smallest hydrophobic area to water, and the polar head groups arrange to maximize their interaction with water.

Chapter 2: Water 11

H-bonds between water molecules are broken (H = Hwater - Hmixture = +)

Water molecules are forced to form a cage around the hydrophobic tail (S = -)

G = H - TS (H = +; S = -; thus

G = +)

This is an unfavorable reaction. Lipids have low solubility in water (biochemists/biologists refer to lipids as being suspended in water rather than dissolved in water)

Why is S = - when it is + in the NaCl example?

Chapter 2: Water 12

--Nonpolar regions cluster together to minimize contact with water (hydrophobic interactions) (H = -, but small)

--Polar regions arrange to maximize interaction with water--Driving force is water, not hydrophobic clustering/packing!

--Thermodynamic stability gained by minimizing # of ordered H2O molecules S = + and large

Chapter 2: Water 13

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--release of water molecules is an increase in entropy (S = +)

packing of hydrophobic chains is decrease in enthalpy (H = -)

Critical micelle concentration (CMC): The concentration at which micelles begin to readily form. A property of each lipid -think of it as Keq.

Chapter 2: Water 14

Hydrophobic molecules (e.g. hexane, testosterone) dissolve in water via a similar concept.

Water forms an ordered cage (clathrate structure) around

hydrophobic molecules.

Chapter 2: Water 15

--Enzyme and substrate have a shell of water that hydrates/surrounds them

--Binding of enzyme to substrate releases some of this. This is ann increase in entropy of water (S = +)

--Decrease in enthalpy (H = -) as a result of more H-bonds, ionic interaction, and hydrophobic interaction.

----------------------------

Water in biochemistry

Chapter 2: Water 17

Proteins arrange water molecules next to one another like chain/wire

Chain of water molecules allows for protons to hop from one end to

the other end

Proton hopping is faster than diffusion

Used by some membrane proteins to regulate influx of protons into

organelles

Hydronium ion: (H3O+)

Proton hopping is much faster than diffusion explains the high mobility of H+ ions compared to other noncovalent cations such as Na+ or K+.

Proton Hopping

Bond strengthsCovalent bonds (sharing of electron pairs between atoms. ~300 - 400 kJ/mol)

19Chapter 2: Water

Non-covalent bonds (~3-40 kJ/mol)

Interactions in the aqueous environmentNon-covalent bond interactions in water (3-40 kJ/mol)

20

Ionic EquilibriaAcids are proton donors

Bases are proton acceptors

Many biological molecules are weak acids or weak bases

!The ionization of water and the ion product

[H2O] + [H2O] [H3O+]+[OH-]

[H2O] [H+]+[OH-]

[H+][OH-] = Kw = 1 x 10-14 M2

[H+]=[OH-] = 1 x 10-7 M

The dissociation of an acid can be written as

HA+ A + H+HA A- + H+HA- A2- + H+

The equilibrium constant (Ka) is defined as

Ka = [H+][A-]/[HA]

The strength of the acid is usually expressed in terms of pKa

pKa = -log(Ka)

Chapter 2: Water 25

pH = - log [H+]

pX = - log [X]

------------------------------------------------------

-----------------------------------------------------Neutral

acidic

basic

--logarithmic, not arithmetic

--1 pH unit = 10X change in [H+]

--pH affects activity of biological macromolecules (e.g., enzymes)

--pH of blood or urine commonly used in medical diagnoses (acidosis vs. alkalosis)

Chapter 2: Water 26

pH of Common Aqueous Liquids

Acids are proton donors

Bases are proton acceptors

A proton donor and its corresponding acceptor are called a conjugate acid-base pair

The Henderson-Hasselbalch EquationDescribes the chemical composition of a buffer as a function of pH

[HA] is the concentration of the weak acid

[A-] is the concentration of the conjugate base

[H+] is the concentration of proton

Ka is a known quantity particular for an acid/base

Chapter 2: Water 27

[HA] [H+] [A-]

deeStamp

Henderson-Hasselbalch Equation

pKa is a measurement of the tendency for a weak acid (of a

conjugate acid-base pair) to lose a proton in aqueous media

if the pH is less than the pKa (by at least one pH unit), then

the H-HEqn indicates that the weak acid form (i.e., the proton

donor) will be the dominant species ([HA]>>[A]-)

if the pH=pKa, then the weak acid (proton donor) and its

conjugate base (proton acceptor) will be present in equal

amounts ([HA]=[A]-)--a.k.a.,"half-equivalence point". Also,

effective pH buffer region (i.e., pH range resistant to change

in pH upon addition of acid or base) = pKa 1

if pH>pKa (by at least one pH unit), then the conjugate base

(proton acceptor will be the dominant species ([A]->> [HA])

this is reason why at physiological pH (~7), amino acids are

zwitterionic, viz.,

Chapter 2: Water 28

Titration of weak acidsHenderson-Hasselbalch Equation: relates pH, pK, and buffer concentration. If you know any two the third can be calculated

[A-]

pH = pKa + log [proton acceptor][proton donor]

acetate ion

Acetic acid[HA]

What is the pH of a buffer mixture containing 1M acetic acid and 0.5M sodium acetate?

The Ka of acetic acid is 1.74 x 10-5 M

The blue box is the pH at which acetic acid has its greatest buffering capacity

Histidine Titration Curve

Titration of Three Weak Acids

Weakest Acid

Ammonia

Ammonium ion

Acetic acidStrongest acid (loses proton most readily)

Di-H

Mono-H

The buffering capacity of any weak acid/base is its pKa 1

A weak acid is 1% ionized at 2 pH units below its pKa and 99% ionized at 2 pH units above its pKa

Chapter 2: Water 29

Enzymes typically have a pH optimum, as seen for pepsin (active in stomach), trypsin (acts in small intestine), and alkaline phosphatase (active in bone, role in bone mineralization).

Chapter 2: Water 30

Enzymatic activity and pH

Chapter 2: Water 31

--The bicarbonate buffer system is important in blood

--Three reversible reactions: CO2 in lungs affects equilibrium

--Hyperventilation: Excessive breathing reduces CO2 in blood and increases pH of blood. Hypoventilation is the opposite!

(bicarbonat

(carbonic

(dissolved

ga

pK=6.37

Buffering the blood