Water Biochemistry by Laura Murray

8/9/2019 Water Biochemistry by Laura Murray

http://slidepdf.com/reader/full/water-biochemistry-by-laura-murray 1/59

Biochemistry

Dr. Amjad Mahasneh

Office: PH4 L0

Office Hours: Sun, Tue, Thu (10:15-11:15)Mon, Wed (2:30-3:30)

http://elearning.just.edu.jo



Water: The solvent for Biochemical reactions

Chapter 2



When two atoms of the same

electronegativity form a bond, the electrons

are shared equally non-polar bond However, if they differ in electronegativity,

electrons are not shared equally polar

bond



The bond in a molecule may be polar, but

the molecule itself can still be nonpolar because of the geometry e.g. CO2

- 2+ -

OCO

The two CO bonds are polar, but

because the CO2 molecule is linear, the

attraction of the oxygen for the electrons inone bond is cancelled out by the equal

and opposite attraction for the electrons by

oxygen on the other side of the molecule



Fig. 2-1, p.35

The structure of water (bent). Oxygen has a partial negative charge, and the hydrogens

have a partial positive charge. The uneven distribution of charge gives rise to the

large dipole moment of water. The dipole movement in this figure points in the direction

from negative to positive, the convention used by physicists and physical chemists;

organic chemists draw it pointing in the opposite direction.



Solvent Properties of Water

Why do some chemicals dissolve in water

while others do not?

The polar nature of water largely

determines its solvent properties.

Ionic compounds with full charges, such

as potassium chloride (KCL, K+ and Cl- in

solution), and polar compunds with partialcharges (i.e dipoles) such as ethylalcohol

or acetone, tend to dissolve in water



Hydration shells surrounding ions in solution. Unlike charges attract. The

partial negative charge of water is attracted to positively charged ions.

Likewise, the partial positive charge on the other end of the water molecule isattracted to negatively charged ions. Fig. 2-2, p.35



Ionic and polar substances are referred to as

hydrophilic

It is less favorable thermodynamically for water molecules to be associated with nonpolar

molecules than with other water molecules. As

a result, non-polar molecules like hydrocarbons

don¶t dissolve in water and are referred to ashydrophobic

A nonpolar solid leaves undissolved material in

water. A nonpolar liquid forms a two-layer system with water

The interaction between nonpolar molecules

are called hydrophobic interactionsTable 2-2, p.37



Why do oil and water mixed

together separate into layers? A single molecule may have both polar

(hydrophilic) and nonpolar portions

(hydrophobic) amphipathic e.g. fattyacid

Such a compound in the presence of

water tends to form structures calledmicelles



An amphiphilic molecule: sodium palmitate

An amphiphilic molecule: sodium palmitate

Amphiphilic molecules are frequently symbolized by a ball and zigzag line structure, where

the ball represents the hydrophilic polar head and the zigzag line represents the nonpolar

hydrophobic hydrocarbon tail. Fig. 2-4, p.37



Micelle formation by amphipathic molecules in aqueous solution.

When micelles form, the ionized polar groups are in contact with the water,

and the nonpolar parts of the molecule are protected from contact with the

water.

Fig. 2-5, p.37



Interactions between nonpolar molecules

themselves are very weak and depend on

the attraction between short-lived

temporary dipoles and the dipoles they

induce.

It is called van de Waals interactions



Hydrogen Bonds

H bond is of electrostatic origin and can be

considered a special case of dipole-dipole

interaction.

When H is covalently bonded to a veryelectronegative atom such as O or N, it

has a partial positive charge due to the

polar bond. This partial positive charge onH can interact with an unshared pair of

electrons on another electronegative atom.



A comparison of linear and nonlinear hydrogen bonds.

Nonlinear bonds are weaker than bonds in which all three atoms lie in a straight

line.

Fig. 2-6, p.38



A comparison of the numbers of hydrogen bonding sites inHF, H2O, and

NH3. (Actual geometries are not shown.) Each HF molecule has one

hydrogen-bond donor and three hydrogen-bond acceptors. Each H2O

molecule has two donors and two acceptors. Each N

H3 molecule has threedonors and one acceptor. Fig. 2-7, p.39



The geometric arrangement of H-bonded water

molecules has important implications for the

properties of water as a solvent.

The bond angel is 104.3° and the angel

between the unshared pairs of electrons is

similar. The result is a tetrahedralarrangement of water molecules.

The hydrogen bonding between water

molecules can be seen more clearly in the

regular lattice structure of the ice crystal

In liquid water, H-bonds are constantly

breaking and new ones are constantly forming



Tetrahedral hydrogen bonding inH2O: an array of H2O

molecules in an ice crystal. EachH2O molecule is hydrogen-

bonded to four others.Fig. 2-8, p.39



H-bonds are much weaker than normal covalent bonds

Table 2-3, p.40



H-bonding also plays a role in the behavior of water as a

solvent.

Fig. 2-9, p.41

Examples of hydrogen bonding between polar groups and water.



Other Biologically important H-bonds

H-bonds have a vital role in stabilizing the

3-dimensional structures of DNA, RNA

and proteins

Proteins helix and -pleated sheet



Table 2-5, p.41



Acids, Bases and pH

acid: a proton (hydrogen ion) donor

Base: a proton acceptor

Depends on the chemical nature of thecompounds.

Complete dissociation for strong acids to

no dissociation for very weak acids

Acid strength: amount of H ion released

when a given amount of acid is dissolved

in water acid dissociation constant or ka



What are acids and bases?

HA Aí + H+

HA + H2O() Aí + H3O+

acid

acid

base

Conjugate

base

Conjugate

Acid to H2O

Conjugate

Base to HA



What is pH?



The ionization of water

Fig. 2-10, p.43

H2O OHí+ H+



Fig. 2-11, p.43

The hydration of hydrogen ion in water



pH= - log10 [H+]

When a solution has a pH of 7 neutral

Acidic solutions have pH values lower than7 and basic solutions have pH values

higher than 7



In biochemistry, most of the acids

encountered are weak acids. These have

a Ka well below 1.

To avoid having to use numbers with large

negative exponents, a similar quantity pKahas been defined

pKa= - log10 Ka

The higher the Ka (smaller the pKa ) thestronger the acid



Table 2-6, p.45Acid dissociation constant ka and the pKa for a number of acids



Why do we want to know the pH?

An equation connects the Ka of any weak acid with thepH of a solution containing both that acid and its

conjugate base.

Necessary to control pH for optimum reaction conditions

Important biological macromolecules lose activity atextremes of pH

pH = pKa + log [A-]/[HA]

Henderson-Hasselbalch equation

Useful in predicting the properties of buffer used to control

the pH of reaction mixtures



Fig. 2-12, p.45

pH versus enzymatic activity.



Titration Cur ves

A titration is an experiment in whichmeasured amounts of base are added to a

measured amount of acid

Reaction is followed using a pH meter The point the titration at which the acid is

exactly neutralized is called the

equivalence point

The inflection point in the titration curve is

reached when the pH equals the pKa of

the acid



Fig. 2-13, p.46Titration curve for acetic acid. Note that there is a region near the p K a at which the titration curve is relatively flat.In other words, the pH changes very little as base is added in this region of the titration curve.



When the pH of a solution is less than the

pKa of an acid, the protonated form

predominates

When the pH of a solution is greater than

the pKa of an acid, the deprotonated

(conjugate base) form predominates



Buffers

A buffer is something that resists change.

A buffer solution tends to resist change in

pH when a small to moderate amounts of

a strong acid or base are added A buffer consists of a mixture of a weak

acid and its conjugate base



How do buffers work?



Fig. 2-14, p.49



How do we choose a buffer?

The pH of a sample being titrated changesvery little in the vicinity of the inflection

point of a titration curve (fig. 2.15a)

A buffer solution can maintain the pH at arelatively constant value because of the

presence of appreciable amounts of both

the acid and its conjugate base

This condition is met at pH values at or

near the pKa of the acid (Table 2.7)



The relationship between the titration curve and buffering action in H2PO4 -.

(a) The titration curve of H2PO4 -, showing the buffer region for the H2PO4-/HPO4

2-

pair. Fig. 2-15a, p.50



The H2PO4/HPO4

-

pair is suitable as abuffer near pH 7.2

At pH values below the pKa, the acid form

predominates, and at pH values above the

pKa, the basic form predominates. The

plateau region, where the pH does not

change rapidly covers a pH range

extending approximately one pH unit oneach side of the pKa (fig. 2.15b)



The relationship between the titration curve and buffering action in H2PO4 -.

(a) The titration curve of H2PO4 -, showing the buffer region for the H2PO4-/HPO4

2-

pair. Fig. 2-15a, p.50



Fig. 2-15b, p.50

The relationship between the titration curve and buffering action in H2PO4 -

(b) Relative abundance of H2PO4 ± and HPO4

2- .



Table 2-7, p.51



Buffering capacity

A buffer solution with low concentrations of

both the acid and base forms is said to

have a low buffering capacity.

A buffer that contains greater amounts of both acid and base has a higher buffering

capacity



How do we make buffers in the laboratory?

To have a buffer, all that is necessary arethe two forms of the buffer present in the

solution at reasonable quantities.

By adding predetermined amounts of the

conjugate base form A- to the acid form

HA

To make buffer we could start with HA

form and add NaOH until the pH is correctas determined by pH meter

Or we could start with A- and add HCl

until the pH is correct



Fig. 2-16a, p.53

Two ways of looking at buffers. In the titration curve, we see that the pH varies only

slightly near the region in which [HA] = [A-].

In the circle of buffers, we see that adding OH- to the buffer converts HA to A-.

Adding H+ converts A- to HA.



Are naturally occurring pH buffers present in

living organisms?

The H2PO4 ± /HPO4

2- pair is the principal

buffer in cells.

In blood, phosphate ion levels are

inadequate for buffering, and a differentsystem operates

Based on carbonic acid (H2CO3):

H2CO3 H+ + HCO3-

pKa of H2CO3 is 6.37, The pH of the blood

is 7.4, is near the end of the buffering

range of this system



CO2 can dissolve in water and in water-

based fluids such as blood

CO2 (g) CO2 (aq)

CO2 (aq) + H2O () H2CO3 (aq)

H2CO3 (aq) H+(aq) + HCO3- (aq)Net: CO2 (g) + H2O () H+ + HCO3

- (aq)

At pH of blood, most of the dissolved CO2 is

present at HCO3-

The CO2 being transported to the lungs to

be expired takes the form of bicarbonate

ion



Zwitterions: compounds that have both a

positive and negative charge (Table 2.8)



Table 2-8, p.54



The End



Fig. 2-CO, p.34



p.42



How do buffers work?

Lets compare the changes in pH that

occur on the addition of equal amounts of

strong acid or strong base to pure water at

pH 7 and to a buffer solution at pH 7

If 1 ml of 0.1 M HCl is added to 99 ml of

pure water pH drops drastically

If 0.1 M NaOH added instead of HCl pHrises drastically

Water Biochemistry by Laura Murray

Documents

Transcript of Water Biochemistry by Laura Murray