An isolated system cannot exchange matter or energy with its surroundings. A closed system can exchange only energy with its surroundings. A closed system inevitably reaches equilibrium. Open systems exchange both matter and energy with their surroundings and therefore cannot be at equilibrium. Living organisms must exchange both matter and energy with their surroundings and are thus open systems. Living organisms tend to maintain a constant flow of matter and energy, referred to as the steady state.

Lecture 2: Physical Properties of Water

• Water is a polar molecule

• Hydrogen bonding

• Noncovalent interactions

• Solubility –water of hydration

• Hydrophobic Effect

• Osmosis, Diffusion, Dialysis

Hydrogen Bond between water molecules O atom has partial – charge of – 0.66e H atom has partial + charge of +0.333e

Ice Structure Each water molecule interacts tetrahedrially with 4 other water molecules. H bonds are ~ 0.5Å shorter than van der Waals distance between H bonded atoms. Ice has open structure – water expands on chilling thus ice is less dense than water and floats

H bonding in liquid water Irregular structure with 3, 4 or 5 molecules arranged in ring structures. Very hard to get precise structure for liquid water since molecules are constantly in motion

Noncovalent interactions between neutral molecules

Collectively known as Van der Waals forces

Solvation of ions by water

Water dipoles orient according to ion charge

Uncharged polar molecules dissolve by H bonding of functional groups with solvent water molecules

Water molecules maximize H bonding by forming a “cage” around nonpolar solute

Aggregation of nonpolar molecules in water is entropy driven

(a) individual hydration of dispersed molecules decreases entropy of system (reduces H bonding possibilities of solvent water molecules

(b) aggregation increases entropy of system since fewer water molecules are needed to hydrate aggregate and more H bonding of water molecules is possible

Amphiphiles form micelles (a) and bilayers (b)

Osmosis is movement of solvent across membrane from high concentration (pure water) to lower concentration (water + solute molecules). Osmotic pressure is the pressure required to prevent inward flow of water

Diffusion is random movement of water & solute across membrane permeable to both from high concentrations to low concentrations – thus water moves in and solute moves out of the dialysis bag.

Lecture 3: Chemical Properties of Water

• Ionization of Water

• Acids and Bases

• pH

• Buffers

Water is an important reactant in many biochemical reactionsCondensation Reactions - loss of water Hydrolysis Reactions - addition of water

Condensation

Hydrolysis

Hydrolysis

R O P OHO

O-HO P O

O

O-R O P O P O-

O O

O- O-

-H2O ++Condensation

ACIDS AND BASES

I. Classical (Arrhenius) Definition

Bases are compounds that yield OH- ions whendissolved in H2O

Acids are compounds that yield H+ (H3O+) ionswhen dissolved in H2O

However, often talk about acid-base reactions that involveneither H3O+ or OH- ions.

2 other common systems for describing acid-base Reactions

II. Hydrogen-ion Transfer (Brönsted-Lowry System)

Brönsted acid can give up a proton to another speciesBrönsted base can accept a proton from another species

III. Electron-pair Transfer (more general system)

An acid-base Reaction consists of donation of a pairof electrons from the base to the acid

Brönsted-Lowry system most useful for typical biochemicalreactions such as buffering.

WEAK ACIDS & BASES BUFFERS

Of particular interest are reversible reactions involving hydrogenions. These include weak acids & weak bases which have beendefined by Brönsted as:

HA H+ + A-

Acid conjugate base

B + H+ BH+

Base conjugate acid

A Brönsted acid can ionize to form a proton and its conjugate base

A Brönsted base can react with a proton to form its conjugate acid

CH3 COOH H+ + CH3 COO-

HOAc H+ + OAc-

Incomplete dissociation of these acids and bases distinguish themfrom strong acids (HCl & HNO3) and from strong bases such asNaOH

Conjugate acid-base pair

Since most biochemical reactions take place in H2O, It isnecessary to consider the ionization of water. Water actsas a weak acid:

H2O H+ + OH-

At normal pressure & temperature the concentration ofundissociated H2O is ≅ constant:

1000gL-1 / 18g mol-1 ≈ 55.56 mol / L = 55.56 M

∴ define a new equilibrium constant Kw which includesconcentration of "solvent" H2O

Kw = [H+] [OH-] = 1 x 10-14 mol2/L2 at 25° Cnote the units

This is the basis of the Sorensen pH scale whichranges from 0 – 14 (acidic to basic)

Regardless of the concentration of H+ and OH- ions,which may be contributed by ionization of acids &bases in solution, the ion product of water Kw prevails:

[H+] x [OH-] = 1.0 x 10-14 under normal conditions

mole/L10x1.8O]2[H][OH][HKeq 16−=

−+=

≈ constant

Review: H2O H+ + OH-

116

2eq liter mole101.8

OH]][OH[HK −−

−+

×==

55.56 M pure water (1) Kw = [H+] [OH-] = 55.56 M × 1.8 X 10-16 M

=1.0 X 10-14 mole2 / liter2 @ 25oC

[H+] [OH-] = 1 X 10-14 mole2 / liter2

log [H+] + log [OH-] = -14

define pH = -log[H+] pOH = -log [OH-]

∴pH + pOH = 14

Dissociation of weak acid HA H+ + A-

weak acid conjugate base

[HA]]][A[HKeq

−+= dissociation constant Ka