Water: The molecule of lifePhysical & Chemical Properties

pH concepts & buffers

COURSE TITLE: BIOCHEMISTRY 1COURSE CODE: BCHT 201

PLACEMENT/YEAR/LEVEL: 2nd Year/Level 4, 1st Semester

M.F.Ullah, Ph.DShowket H.Bhat, PhD

Water: The molecule of lifeWater is a chemical substance with thechemical formula H2O. A water moleculecontains one oxygen and two hydrogenatoms connected by covalent bonds.

Water appears in nature in all three common states of matter and may take many different forms on Earth: water vapor and clouds in the sky; seawater in the oceans; icebergs in the polar oceans.

Chemical and physical properties

The major chemical and physical properties of water are:1. Water is a liquid at standard temperature and pressure. It is tasteless and odorless. The

intrinsic colour of water and ice is a very slight blue although both appear colorless in small quantities. Water vapour is essentially invisible as a gas.

2. Water is transparent in the visible electromagnetic spectrum. Thus aquatic plants can live in water because sunlight can reach them. Infrared light is strongly absorbed by the hydrogen- oxygen or OH bonds.

3. Since the water molecule is not linear and the oxygen atom has a higher electronegativity than hydrogen atoms, it carries a slight negative charge, whereas the hydrogen atoms are slightly positive. As a result, water is a polar molecule with an electrical dipole moment. Water also can form an unusually large number of intermolecular hydrogen bonds (four) for a molecule of its size.

8. The boiling point of water (and all other liquids) is dependent on the barometric pressure. B.P is inversely related to pressure. For example, on the top of Mt. Everest where pressure is greater, water boils at 68 °C , compared to 100 °C at sea level where standard pressure is lesser than at higher altitudes (on the top of Mt. Everest). Conversely, water deep in the ocean near geothermal vents can reach temperatures of hundreds of degrees and remain liquid.

4. Water is a good solvent and is often referred to as the universal solvent. Substances that dissolve in water, e.g., salts, sugars, acids, alkalis, and some gases – especially oxygen, carbon dioxide (carbonation) are known as hydrophilic (water-loving) substances, while those that are immiscible with water (e.g., fats and oils), are as hydrophobic (water-fearing) substances.

6. Most of the major components in cells (proteins, DNA and polysaccharides) are also soluble in water.

7. Pure water has a low electrical conductivity, but this increases significantly with the dissolution of a small amount of ionic material such as sodium chloride.

5. Water is miscible with many liquids, such as ethanol, in all proportions, forming a single homogeneous liquid. On the other hand, water and most oils are immiscible, usually forming layers according to increasing density from the top. As a gas, water vapor is completely miscible with air.

9. Water has a high specific heat capacity, as well as a high heat of vaporization , both of which are a result of the extensive hydrogen bonding between its molecules.

10. Normally when liquids attain their solid state they tend to become dense. The maximum density of water occurs at 3.98 °C . It has the anomalous property (unusual behaviour) of becoming less dense, not more, when it is cooled down to its solid

form, ice. It expands to occupy 9% greater volume in this solid state, which accountsfor the fact of ice floating on liquid water, as in icebergs. Its density is 1,000 kg/m3 (62.428 lb/cu ft or 8.3454 lb/US gal) liquid (at 4 °C; ice has a density of 917 kg/m3). This property of water is very important for survival of aquatic life in countries where temperature drops to below freezing point. It prevents the whole water of lakes and ponds and other water bodies to freeze. As a layer of water freezes, it starts floating on the surface and traps the heat inside the water bodies to allow quatic plant and animal to survive the freezing temperature.

pH is defined as the negative logarithm of hydrogen ion concentration of a given solution. pH is related inversely to the hydrogen ion concentration in a solution.[H+] pH[H+] pH

pH = -log [H+]pH only refers to hydrogen ion concentration and is meaningful when applied to aqueous (water-based) solutions.

pH (potential of Hydrogen) : Basic Concepts

When water dissociates, it exist in equilibrium with hydrogen ions and hydroxide ions: H2O <--> H+ + OH-

The dissociation constant of pure water (Kw) is the product of its hydrogen ion and hydroxyl ion concentrations and is given as:Kw = [H+] [OH-]

Experimentally, it has been found that in pure water the concentration of:[H+] = [OH-] = 10-7

Therefore: Kw = [10-7][ 10-7] = [10-14](To multiply exponential numbers - simply add the exponents)

Water Equilibrium Principle: The multiplication product of H+ and OH- ion concentration must always be equal to 10-14.

The pH of pure water:The hydrogen ion concentration of pure water [H+] = 1x10-7 pH= -log [10-7]= 7

However as acids (which give hydrogen ion) or bases (which give hydroxyl ion) are added to the pure water, the concentration of hydrogen ion and hydroxyl ion varies (changes). The change in the effective concentration of hydrogen ion in the solution by acid (which directly increases hydrogen ions) or by base (which decreases hydrogen ions by neutralization with hydroxyl ion) changes the pH of the pure water.

pH of pure water is taken as a point of neutrality i.e pH 7 is the neutral pH value (neither acidic nor basic). Pure water is neutral. However when chemicals are mixed with water, the mixture can become either acidic or basic.

For an Acidic Solution: [H+] > 1x10-7 pH< 7 (acids have pH below 7) Basic Solution: [H+] < 1x10-7 pH>7 (bases have pH above 7)When calculating pH, remember that concentration [] refers to Molarity, M.

Example: 1. If an acidic solution has an H+ concentration of 0.0001 M, find the pH.

Solution:First convert the number to exponential notation, find the log, then solve the pH equation.

[H+] = 0.0001M = 10-4;

pH = - log [ H+]

= - log (10-4) = - (-4) = +4 = pH

Or simply pH = -log (0.0001)= 4

The purpose of the negative sign in the log definition is to give apositive pH value.

2. If a basic solution has an OH- concentration of 0.001M, find the pH.

Solution: In order to calculate the pH, first find the hydrogen ion concentration using the water dissociation constant equation (Kw).Kw= [H+] [OH-]From our knowledge of the Water Equilibrium Principle, we know that Kw = 10-14 and the product of [H+] and [OH-] must always be equal to Kw (10-14)Thus, [H+] = Kw/[OH-][H+]= 10-14/0.001= 10-11

pH= -log [10-11] = 11

Calculations:

pH Scale:

The pH scale, (0 - 14), is the full set of pH numbers which indicate the concentration of H+ and OH- ions in water.

Solutions with pH below 7 down to 0 are acidicSolutions with pH above 7 up to 14 are basic

A buffer solution is one that resists pH change on the addition of small amount of acid or alkali (base).

Buffer solutions are used in many biochemical experiments where the pH needs to be accurately controlled. Many life forms thrive only in a relatively small pH range so they utilize a buffer solution to maintain a constant pH.

One example of a buffer solution found in nature is blood. In living beings physiological pH is maintained in blood by means of physiological buffers such as bicarbonate, phosphate and proteins.

A buffer is an aqueous solution consisting of a mixture of a weak acid and its conjugate base or a weak base and its conjugate acid.

Buffer Solutions

Strong Acids: these are acids which dissociate completely in aqueous solution into hydrogen ion and their conjugate base so that the hydrogen ion concentration is same as that of the acid.

HA H+ + A- (completely ionized/dissociated)(acid) (hydrogen ion) (conjugate base)

e.g. HCl H+ + Cl- (Hydrochloric acid)

Weak Acids: Weak acids are only slightly ionized in solution and equilibrium is established between the acid and the conjugate base.

HA H+ + A- (Partially ionized/dissociated)(acid) (hydrogen ion) (conjugate base)

e.g. CH3COOH H+ + CH3COO- (Acetic acid)

Buffer is made from weak acid and its conjugate base or a weak base and its conjugate acid. The buffer works on the following principle to resist large changes in pH on addition of small amount of strong acids or bases.

1. When challenged by a strong acid, the hydrogen ion from the acid associates with the conjugate base of the weak acid of the buffer to form the weal acid and since this weak acid is weakly ionized (partially ionized) the hydrogen ion concentration of the solution does not exceeds much and pH variation is less.

e.g: In case of acetic acid buffer

CH3COOH H+ + CH3COO- (Acetic acid)

HCl H+ + Cl-

Addition of HCl (a strong acid) will add H+ ions to the buffer solution but these H+ ions will associate with acetate ions (conjugate base in the buffer) to form acetic acid which is a week acid and does not ionized to release large amount of H+ ions and therefore H+ concentration is not affected much maintaining the pH.

Principle of buffer action

2. When challenged by a base, the hydroxyl ion from the base extracts the hydrogen from the weak acid of the buffer to form water and conjugate base. Thus the addition of hydroxyl ions does not change the concentration of hydrogen ions in the solution (which it would have done through neutralization reaction in case of strong acids in buffer). Thus the pH is not much affected.

e.g: In case of acetic acid buffer

CH3COOH H+ + CH3COO- (Acetic acid)

NaOH Na+ + OH-

When NaOH is added to the buffer solution, Hydroxyl ions from the base extracts the hydrogen ion from the acetic acid forming water and acetate (Conjugate base). Therefore H+ ion concentration in the solution remains least affected. Thus the pH variation is least.

Selection of weak acids for buffer solution:

The pKa of an acid is the pH at which the acid is 50% dissociated so that the concentration of conjugate base [A-] is equal to the concentration of undissociated acid [HA].

The relationship between pH and pKa is given by the equation:pH= pKa + log10 [Conjugate base (A-)] / [Acid (HA)]

To select the weak acid for a buffer solution with desired pH , e.g. buffer with pH 4.0, a weak acid with pKa close to this desired pH is selected to prepare the buffer as buffering range lies at pKa ± 1.

E.g. For a pH 4.0 buffer, acetic acid with a pKa 4.8 is suitable as it gives buffering range between pH 4.8 ± 1, i.e 3.8 -5.8.

Buffers in blood: Buffer solutions are necessary to keep the correct pH for enzymes in many organisms to work. Many enzymes work only under very precise conditions; if the pH moves outside of a narrow range, the enzymes slow or stop working and can denature. In many cases denaturation can permanently disable their catalytic activity.

A buffer of carbonic acid (H2CO3) and bicarbonate (HCO3−) is

present in blood plasma, to maintain a pH between 7.35 and 7.45.

The majority of biological samples that are used in research are made in buffers, especially phosphate buffered saline (PBS) at pH 7.4.