Solutions. Acid–base equilibrium in biological systems презентация

Solutions. Acid–base equilibrium in biological systems

Plan 0. Solutions and their colligative properties 1. The theory of electrolytic dissociation. Dissociation of bases, acides and salts in water solutions.Strong and weak electrolytes 2. Protolytic theory. 3. Dissociation of water. Hydrogen ion exponent. The homeostasis. 4. The importancy of pH maintenance in human body. 5. The concept of buffer solutions. 6. Hydrocarbonate buffer system 7. Phosphate buffer system 8. Protein buffer systems 9. Hemoglobin buffer system 10. Acidosis and alkalosis. Treatment of acidosis and alkalosis.

The theory of electrolytic dissociation

Electrolytic dissociation – process of decomposition of solutes in the solvent into ions.

1)                Substances dissociating in solutions or melts into positively charged Cat+(cations) and negatively charged An- (anions). The latter include acids, bases and salts. 2)                In electric field Cat+ move to cathode, An- move to anode. 3)                Electrolytes decompose into ions in different degree. 4)                Dissociation depend of: a)     nature of electrolyte; b)    nature of solvent; c)    concentration; d)    temperature.

Dissociation of bases, acides and salts in water solutions

Acides are compounds dissociating in aqueous solutions with the formation of positive ions of one species – hydrogen ions. HCl→H+ + Cl- Bases are compounds dissociating in aqueous solutions with the formation of negative ions of one species – hydroxide ions OH-. Ca(OH)2→Ca2++ 2OH- Medium salts dissociate to form metal cations and anion of acid radical.

Degree of dissociation α Ni - the number of molecules, dissociating into ions; Ntot – the total number of dissolved molecules.

Strong electrolytes Majority of salts. Some acids (HCl, HBr, HI, HNO3, HClO4, H2SO4). Alkalis (LiOH, NaOH, KOH, RbOH, CsOH, Ca(OH)2 , Sr(OH)2, Ba(OH)2)

Weak electrolytes Majority of acids and bases (H2S, H2CO3, Al(OH)3, NH4OH).

The dissociation of weak electrolytes is a reversible process CatAn Cat+ + An-

The equilibrium constant K is called the dissociation (ionization) constant

Ostwald dilution law

Acidity and basicity constants The dissociation constants of acids and bases, respectively called acidity constants (KA) and major (KB). Product constant acidity and basicity constants, with the acid conjugate base is the ion product of water:

Dissociation of water H2O H+ + OH-

Kw is constant, ion product of water.

Hydrogen ion exponent pH= -lg [H+]

Protolytic theory Danish physicist and chemist Johannes Brønsted and the English chemist Thomas Lowry in 1928-1929 was offered Protolytic (protonic) theory of acids and bases, according to which:

Base - a substance (particle) that can attach proton (i.e. base - proton acceptor). Acid- a substance (particle) that can donate proton (i.e. acid – proton donor) In the general form:

Salt - the reaction product of acid and base Example:

The homeostasis. The importancy of pH maintenance in human body The human body has mechanisms of coordination of physiological and biochemical processes proceeding inside it and maintenance constancy of internal medium (optimal value of pH, levels of different substances, temperature, blood preassure). This coordination and mantanance are called homeostasis.

The constancy of hydrogen ions concentration is one of important constant of internal medium of organism, because: 1) Hydrogen ions have catalytic effect on many biochemical processes; 2)Enzymes and hormones exhibit biological activity only at a specific range of pH values; 3)Small changes of pH in blood and interstitial fluids affect the value of the osmotic pressure in this fluids.

pH values of different biological fluids and tissues of the human body

The concept of buffer solutions Buffer solutions are solutions that resist change in hydrogen ion and the hydroxide ion concentration (and consequently pH) upon addition of small amounts of acid or base, or upon dilution.

The resistive action is the result of the equilibrium between the weak acid (HA) and its conjugate base (A−): H+(aq) + A−(aq) → HA(aq) OH-(aq) + HA(aq) → A−(aq) +H2O(l)

Henderson-Hasselbah equation

Buffer capacity Buffer capacity (B) - the number of moles of equivalents of strong acid or alkali to be added to 1 liter of buffer solution to shift the pH unit Вac.= Вbas.=

Buffer capacity Buffer capacity is maximal at a ratio of acid salt 1:1 => pH = pK. Good – at [pK+0.5, pK-0.5] Sufficient – at [pK+1, pK-1] The higher the concentration of the solution, the greater the buffer capacity. The concentration of acid and salt in the buffer solutions usually about 0.05-0.20 M.

The relative contribution%   buffer systems in the blood to maintain homeostasis it protolytic Buffer systems plasma Hydrogen carbonate 35% Protein 7% Hydrogen phosphate 1% TOTAL 43% Buffer systems erythrocytes Hemoglobin 35% Hydrogen carbonate 18% Hydrogen phosphate 4%

Hydrocarbonate buffer system HCO3- +H+ H2CO3 H2CO3+OH- HCO3-+ H2O CO2+ H2O H2CO3

pKa1(H2CO3)=6.1 pKa1(H2CO3)=6.1 pH of a blood plasma = 7.4

Alkaline reserve HCO3-+ H+ H2CO3 CO2+ H2O

Phosphate buffer system HPO42-+H+ H2PO4- H2PO4-+OH- HPO42-+H2O

The mechanism of action of phosphate buffer: 1. acid addition 2 Na++HPO42–+H++Cl - NaH2PO4+Na++Cl - 2. adding alkali : NaH2PO4 + NaOH  Na2HPO4 + H2O Excess hydrogen phosphate monobasic and removed through the kidneys. Full recovery of relations in the buffer occurs only 2-3 days.

pKa(H2PO4-)=6.8 pH of a blood plasma = 7.4

Protein buffer systems The plasma proteins (albumins, globulins) are less important than the hemoglobin for maintenance of pH.

PROTEIN acid-base buffer system

Hemoglobin buffer system

Hemoglobin acid-base buffer system BLOOD

Binding of hydrogen cations imidazole groups of hemoglobin.

Hemoglobin buffer system HHb + O2 HHbO2 Hemoglobin is a weaker acid (pKa HHb = 8.2) than oxyhemoglobin (pKa HHbO2 = 6.95). Therefore Hb- ions being anions of weaker acid are capable stronger to bind H+ ions than HbO2- ions. Undissociated molecules HHbO2 lose O2 easier than the ions HbO2-

a) the hemoglobin buffer system: HHb H+ + Hb-; b) the buffer system formed by oxyhemoglobin: HHbO2 H+ + HbO2-.

In erythrocytes: HHbO2 HHb + O2 (1) HHbO2 H+ + HbO2- (2) HbO2- Hb- + O2 (3)

In vessels of tissues

In vessels of tissues CO2+ H2O H2CO3 HbO2-+ H2CO3 HHbO2 + HCO3- HHbO2 HHb + O2

In lungs

In lungs HHb + O2 HHbO2 HHbO2+ HCO3- HbO2-+ H2CO3 H2CO3 CO2+ H2O

Acidosis and alkalosis

Literature 1. Medical Chemistry : textbook / V. A. Kalibabchuk [and al.] ; ed. by V. A. Kalibabchuk. - K. : Medicine, 2010. 2.http://www.chemeurope.com/en/encyclopedia/Buffer_solution.html