ACID - BASE EQUILIBRIA

Mgr. Monika Šrámková

Department of medical chemistry and clinical

biochemistry,

2th Medical faculty of Charles Univerzity of

Prague and Motol Univerzity Hospital

CHEMICAL EQUILIBRIA- INTRODUCTION

Chemical equilibria:

a) Acid - base

b) Redox

c) Solubility

d) Complexation

Law of conservation of mass

The mass of reactants is exactly equal to the mass of products

AB + CD → AD + CB

Law of charge conservation

The total amount of electric charge entering the reaction is exactly

equal to the amount of charge going out of the reaction

0 ii zc

DYNAMIC EQUILIBRIA

At dynamic equilibrium, reactants are converted

to products and products are converted to

reactants at an equal and constant rate

The velocities of reactions are specified by Law

of Guldberg – Waag

v1 = v2 resp. k1. cA.cB = k2. cC.cD

Cato

Maximilia

n

Guldberg

Peter Waage

EQUILIBRIUM CONSTANT

We derive the equilibrium constant from the

dynamic equilibrium relation.

K is concentric equilibrium constant

If the stoichiometric coefficients in the chemical

reaction are not equal to one, chemically reacting:

k.K + l.L = m.M + n.N

BA

DC

2

1

c*c

c*c

k

kK

EQUILIBRIUM CONSTANT

(cM)m.(cN)n

K = ----------------

(cK)k.(cL)l

Equilibrium Concentration Constant is defined as the

fraction of the concentrations of reactants promoted to

stoichiometric coefficients and the product elevated to

stoichiometric coefficients

K 1 : in the reaction predominate reactants

K 1 : in the reaction predominate products

K is close to one: approximately the same amount of

reactants and products

THE PRINCIPLE OF ACTION AND REACTION-

PRINCIPLE OF LE CHATELIERŮV

Application of the general principle of action and

reaction

External action on the chemical system in

equilibrium causes the processes that the system

tries to eliminate external action

External action causes:

1) change of concentration of one component in

the reaction

2) change of temperature

3) change of pressure

Henry Louis

Le Chatelier

THEORY OF ACID-BASES (PROTOLYTIC)

REACTION

The reaction of acids and basis are protolytic

reaction

Arhenius theory (1884) -

acid is a hydrogen-containing compound that , in

water produces hydrogen ions (H+)

HCl → H+ + Cl-

H2SO4 → 2 H+ + SO42-

Base is a hydroxide-containing compund that , in

water, produces hydroxide ions (OH-)

NaOH → Na+ + OH-

Ca(OH)2 → Ca2+ + 2 OH-

THEORY OF ACID-BASES REACTIONS

Theory of Brönsted-Lowry (1923)

Acid is substance that can donate a protone

Base is substance that can accept protone

HNO3 → H+ + NO3-

NH3 + H+ → NH4+

HNO3 + H2O → H3O+ + NO3

-

Conjugate Acid-Bases pair is two species, one

an acid and one a base, that differ from each

other through the loss or gain of a proton

THE ION PRUDUCT OF WATER

The water is an ampholyte substance that can

either lose or accept a proton thus can function as

acid or base

Autodissosiation (self-ionization) of water:

The dissosiation ow water does not influence high concentration of molecular water

(55.56mol/l), therefore it can be considered as constant

The equation of ion product constant for water

is:

Kv = [H3O+] [OH-]=[10-7][10-7]=10-14

The concentrations of both ions are identical at 25 ° C and 100 kPa

ACIDITY AND BASICITY OF

SUBSTANCES,VALUE PH

According to the equation, H + and OH- are

inversely proportional

If c (H3O+) > c (OH-), solution is acidic

If c (H3O+) < c (OH-), solution is basic

If c (H3O+) = c (OH-), solution is neutral

Most of physiological solutions have a

concentration of hydrogen ions near a neutral

point

Søren Sørensen (1909) suggested a more practical

quantity known as pH

pH = – log [H3O+] pOH= 14 – pH

The letter p means „negative logarithm of“

DISSOCIATION CONSTANT

Describe reactions of acid and bases

The constant give an indication of the strenght of

acids and bases

Its value is given by the tendency of compound to

give or to accept a hydrogen ion

pK = – logK

If the value of K is high, it means, that practically

100% of the acid is dissociated.

The numeric value of weak acids/bases is low

The extent of proton transfer for weak acids is

ussualy less than 5%

STRENGHT OF ACIDS AND BASES

We differ three types of acid and bese acoording to

value of dissociation constANT:

a) Strong - K > 10-2, pK < 2 (HCl, HNO3,(COOH)2, NaOH)

b) Middle strong– K10-4 - K10-2 , pK 2 – 4 (H3PO4 )

c) Weak – K < 10-4, pK > 4 (H2CO3, NH3, urea )

HA

A OH3HA

K

B

OH BHB

K

STRONG AND WEAK ACIDS AND BASES

Strong acids are capable of complete dissociation

(release of proton), even in a strongly acidic

solution

Strong acids= high KA and low pKA

The weak acids in the acid solution release the

proton only partially

Bases behave similarly

Most of biochemicals compounds behave like

weak acids and bases

The pH of weak acids and bases must be

calculated using a dissociation constant

CALCULATION OF PH OF SOLUTIONS OF

STRONG MONOTROPIC ACIDS

Is acid that supplies one proton per molecule

during an acid-base reaction

Because the concentration of acid is equal to

concentration of H3O+, the pH can be obtained

from the realtionship: pH= – log(H3O

+)

According to number of hydrogen ion we can

classified also diprotic or triprotic acids, than we

multiply the concentration by number of

hydrogen.

OH3...3,2logpH

CALCULATION PH OF STRONG BASES

Process of calculation of pH of strong bases is

similar to strong acids

We obtained value of pH from relationship:

pH= 14 – pOH

OHlogpOH

CALCULATION OF PH OF WEAK ACIDS

There are oxonium cations and anions of the

acids together with acid and water molecules in

aqueous solutions of weak acids and therefore the

concentration of oxonium cations is not equal to

the concentration of acid

If the non-dissociated and dissociated form is of the same concentration,

the numerical value of pK is equal to pH

The pKA value is tabulated, if adds one equivalent of base to one

equivalent of acid while experomantal determination of PK, the measured

pH is then equal to pK

HXlogp2

1pH AK

AA logp KK

01

1log pKpKpH

DISTRIBUTIVE DIAGRAM

It is a graphical illustration of the equilibrium

composition of the protolytic system according to

pH

The diagram is formed by distribution curves

Ф = f(pH)

Ф is distribution coefficient, unnamed number

Ф HA + Ф A= 1

DISTRIBUTION DIAGRAM OF ACID- ACETIC

ACID, PK=4,755

pKCH3COOH

Ф

pH

DISTRIBUTION DIAGRAM OF POLYPROTIC

ACID-OXALIC ACID

pK1 = 1,25

pK2 = 4,285

Ф

pH

CALCULATION OF PH OF WEAK BASES

There are hydroxide anions of OH- and cations of

the metals together with the base and water

molecules In aqueous solutions of weak bases

Therefore, the concentration of hydroxide anions

is not equal to the concentration of the base

MOHKB logp2

114pH

BKlogpKB

HENDERSON-HASSELBACH EQUATION

It is derived from the classical dissociation

reaction

Indicates,if there is more conjugate base than

acid in a solution, the pH is higher than pKA

If there is more acid than conjugate base, the pH

is lower than pak

By equation is given relationship between pKA

and pH for buffer solution in which is acid-base

pair concentration ratio is something other than

1:1

From a practical point of view, the relationship is

formulated into pH dependence on the ratio of dissociated

and non-dissociated forms (conjugated bases and its acid)

HENDERSON-HASSELBACH EQUATION

HA

A H

K HA A H Kequilibrium constant

after cross-

multiplication

After a logaritmic

calculation and

multiplying by -1 of

both sides

A

HAloglogHlog K

Replacing -log [H +]

with pH and -log K for

pK is obtained

HA

AlogppH

-

K

BUFFERS

Keep pH of a solution from changing very much

when either strong acids or strong bases are

added to it

To handle the addition of both acids and bases, a

buffer contains acid (to react with added base)

and a base (to react with added acid)

Therefore, a combination of an acid-base

conjugate pair is used to prepare a buffer

Most buffers solution consist of nearly equal

concentrations of a weak acid and a salt

containing its conjugate base

Buffers may also contain a weak base and salt

containing its conjugate acid (less common)

BUFFERS

They are of great importance in maintaining the

physiological environment of cells (even in the

extracellular space)

One of the most effective mechanisms to ensure

the appropriate course of biochemical reactions

The important systems of buffers are:

a) Hydrogencarbonate buffer (HCO3- / H2CO3

-)

b) Phosphate buffer(H2PO4- / HPO4

2- )

c) System of intracellular proteins The buffering effectis more egffective, if the value of pH of the solution

approaches the value of pK. In a simplified way: so it is necessary to add

more acid or conjugate base to change value of pH of solution(add or

decrease the concentration of proton)

THE PH OF AQUEOUS SALT SOLUTIONS

Hydrolysis of salts is another example of acid-

base equilibria in water solution

When acids and bases are mixed, they react with

one another and their acidic and basic properties

disappear, we say, that they neutralized each

other

Salts are the products of acid-base neutralization

Most of salts disociate to their cation and anion

Anion A- derived from weak acid will have

reaction like strong base for example We have to decide, if solution of salt is derived from strong or weak acid or

base

HYDROLYSIS OF SALT

We can devided salts into four catgories

according to character of their cation and anion

a) Type NaCl - salt of strong acid and strong

base-does not hydrolyze, the solution is neutral

b) Type NH4Cl – salt of strong acid and weak

base hydrolyze to produces an acidic solution

c) Type CH3COONa – salt of weak acid and

strong base hydrolyze to produces a basic solution

d) Type CH3COONH4 – salt of weak acid and

weak base hydrolyze to produce a slightly acidic,

neutral, or slightly basic solution, depending on the

relative weaknesses of the acid and base

SOLUBILITY/PRECIPITATION EQUILIBRIA

Solubility equilibria exist in systems, where a

compound with low solubility is formed

A solubility product is a quantitative measure of

solubility of a precipitate

Ks value defines a limit product of concentrations

for which the given precipitate is not forming

It specifies an equilibrium of ion concentrations in

a saturated solution above the precipitate

pKs = - log Ks

The solubility of precipitation is usually very low (≤ 10-4 mol

l-1)

COMPLEXATION EQUILIBRIA

They are formed by mixing the metal ion and

ligand solutions to form sequentially the

coordination compounds

The complexes do not arise at one and the same

time, but by the sequential occupation of the

coordinating spheres of the ion

The equilibrium constant for the formation of the

complex ion is the formation constant – β

β = K1K2K3….Kn

The higher the equilibrium constant , the more

stable is the coordination complex

THANK YOU FOR YOUR ATTENTION