Medical chemistry Autumn 2010 Recommended textbooks: Táborská, Sláma: Medical Chemistry I (General and Inorganic Chemistry) Brno 2006 Dostál: Medical Chemistry II (Organic Chemistry) Brno 2006 Lecture files (ppt,pdf). Acid-base reactions •acid-base theory (recapitulation) •pH of acids and bases (recapitulation) •hydrolysis of salts (recapitulation) •buffers Solutions of electrolytes © Department of Biochemistry LF MU (E.T.) 2009 Terms from the secondary school chemistry that is necessary to repeat to understand the lecture topics: • electrolytes, their types and properties • chemical equilibrium, equilibrium constant • acids, bases, conjugate pairs • autoprotolysis of water, pH • strong and weak acids and bases • calculation of pH of strong and weak acids These terms are included in the textbook Medical chemistry I Electrolytes and non-electrolytes when a compound dissolves in water Molecules are not changed Nonelectrolyte Dissociation to ions Electrolyte Electrolytes + - + - Compounds that dissolve in water with the formation of ions. Electrolytes are ionic compounds or polar molecular compounds ions are surrounded by a certain number of water molecules (hydrated). Classification of electrolytes Strong electrolytes Weak electrolytes They are 100% ionized (fully dissociated) There is an equilibrium between ions and molecules Strong electrolytes: AB (s) " A+ (aq) + B-(aq) NaCl(s) " Na+ (aq) + Cl-(aq) Dissociation is complete Weak electrolytes: AB (s) " AB(aq) D A+ (aq) + B-(aq) HCN(g) " HCN(aq) DH+(aq) + CN-(aq) Dissociation is partial H2O H2O Characterization of electrolytes Compare arrows in dissociation equations: Strong electrolyte " Weak electrolyte D Electrolytes Strong Strong acids (HCl, HBr,HI, H2SO4, HNO3, HClO4, …………) Strong hydroxydes (NaOH, KOH, Ca(OH)2, Mg(OH)2… All salts Na2SO4, NaCl, K2CO3, FeCl3………. Weak Weak acids CH3COOH, H2CO3, HNO2… Weak bases Fe(OH)3, NH3, C6H5NH2….. Strong electrolytes NaCl Cl- Na+ Cl- Cl- Cl- Cl- Na+ Na+ Na+ Cl- Na+ NaCl NaCl NaCl NaCl j0078742[1] Weak electrolytes HCN HCN H+ CN- HCN HCN HCN HCN HCN HCN HCN H+ CN- Examples Nonelectrolyte Strong electrolyte Weak electrolyte Ethanol Urea Glucose Aceton Sucrose Glycerol HCl H2SO4 NaOH Ca(OH)2 NaCl Na2SO4 H2CO3 CH3COOH NH3 Mg(OH)2 HCN pyridine Acids and bases according to Brønsted concept Acid: molecule or ion that can lose a proton H+ HA " H+ + A- acid conjugate base Base: molecule or ion that can bind a proton H+ B + H+ " BH+ base conjugate acid Conjugate pair acid/base HA + B " A- + BH Conjugate pair Conjugate pair Acids and bases in water HA + H2O H3O+ + A- acid B + H2O BH+ + OH- base Water behaves as a base Water behaves as an acid pH Acidity of a medium is assessed according to the concentration of hydrogen cations similarly pH values of some body fluids Body fluid pH Blood plasma Urine Gastric juice Pancreatic juice Intracelular fluid (liver cells) Saliva 7,34 - 7,43 4,8 - 7,5 ≈ 2 7,5 - 8,8 6,4 - 6,5 7 - 8 Dissociation of a strong acid: HA + H2O ® H3O+ + A– The concentration [H+] is equal to the total strong acid concentration cHA : [H+] = cHA and pH = – log [H+] = – log cHA Dissociation of a strong hydroxide: MeOH(aq) ® Me+ + OH– In solutions of monobasic strong hydroxides, [OH–] = cMeOH , pOH = – log [OH–] = – log cMeOH, and pH = 14 – pOH pH of strong acids and strong hydroxides solutions Strong acid and strong hydroxides are strong electrolytes that are fully dissociated in aqueous solutions. What pH has vinegar ? What pH has a solution of vitamin C? Why can I drink solutions of carbonic and phosphoric acid, but not hydrochloric acid? pH of weak acids and bases HA + H2O H3O+ + A- B + H2O BH+ + OH- Disssociation of weak acids and bases Weak monoprotic acid Weak monobasic base equilibrium constant of ionization acid and base ionization constant What does express the pKA (pKB) values ? 1- 3…………..moderately strong 4-8…………….weak >8…………….very weak pK = - log K ~ pK KA – acid ionization constant KB – base ionization constant The lower the value of pKA, the stronger is the weak acid. pKA values of weak acids pKB values of weak bases Relation between KA a KB (pKA a pKB) pKA+ pKB = 14 KA . KB = Kv = 1. 10-14 Calculation of pH of weak acids and bases Weak monoprotic acid HA + H2O H3O+ + A- Weak monobasic base B + H2O BH+ + OH- Strong and weak acid with the same concentration– comparison of pH CH3COOH c=0,1 mol/l pKA = 4,7 pH = 2,35 + 0,5 = 2,85 HCl c=0,1 mol/l pH = -log cHA = - log 0,1 = 1 Why soap has alkaline reaction? Why solution of soda (Na2CO3) is alkaline? Why solution of Na3PO4 is alkaline? MCj00787110000[1] Salts are formed in reaction between an acid and base – neutralization. The final pH of a solution after the neutralization reaction is not always neutral. This is because some salts undergo to hydrolysis Hydrolysis of salts Cl- CH3COO- H2O CH3COOH + OH- Differences in properties of ions Cl- is a spectator anion, it does not react with water CH3COO- is an anion of weak acid, it tends to react with water and form acetic acid. Such amount of CH3COOH will be formed, to be in equilibrium with CH3COOH given by ionization constant Reactions of salts in water examples Example 1: CH3COONa – sodium acetate (salt of weak acid CH3COOH and strong hydroxide NaOH) 1. Dissociation CH3COONa " CH3COO- + Na+ 2. Hydrolysis of the anion of weak acid CH3COO- + H2O D CH3COOH + OH- pH is slightly alkaline Example 2: NH4Cl – ammonium chloride (salt of weak base NH3 and strong acid HCl) 1. Dissociation NH4Cl " NH4+ + Cl- 2. Hydrolysis of cation of weak base NH4+ + H2O D H3O+ + NH3 pH is slightly acidic Example 3: CuCl2 – copper (II) chloride (salt of Cu2+ cation derived from weak Cu(OH)2 and strong HCl) 1. Dissociation CuCl2 " Cu2+ + 2Cl- pH is slightly acidic 3. Hydrolysis of hydrated cation [Cu(H2O)6 ]2+ + H2O D [Cu(H2O)5OH]+ + H3O+ 2. Hydration of a metal cation Cu2+ + 6 H2O " [Cu(H2O)6]2+ Example 4: NH4NO2 (salt of weak base NH3 and weak acid HNO2) 1. Dissociation NH4NO2 " NH4+ + NO2- 2. hydrolysis of cation of a weak base and anion of a weak acid NH4+ + H2O D NH3 + H3O+ NO2- + H2O D HNO2 + OH- pH≈7 Example 5: NaNO3 (salt of strong hydroxide NaOH and strong acid HNO3) 1. Dissociation NaNO2 " Na+ + NO2- 2. No hydrolysis of cation nor anion Na+ + H2O D NO3- + H2O D pH = 7 Hydrolysis of salts - summary cation anion pH Strong base Strong acid Weak base Strong acid Strong base Weak acid Weak base Weak acid Composition of salt - origin of ions Complete How to maintain the constant pH of our blood? Buffer solutions • A buffer solution is a solution able to absorb a certain quantity of acid or base without undergoing a strong variation in pH • It serves to maintain a fairly constant pH value. Simple buffer solutions are mixtures of a weak acid and the conjugate base of that (e.g. acetic acid and sodium acetate) or a weak base and its conjugate acid (e.g. ammonia and ammonium chloride) Function of a buffer Example: solution of (CH3COOH + CH3COONa) Particles present in the solution: CH3COO- CH3COOH Na+ H+ mainly from the salt from the acid from the salt from the acid The equilibrium in the solution: CH3COOH D CH3COO- + H+ The presence of CH3COO- from the salt supresses dissociation of CH3COOH The concentration of H+ ions (and also pH) depends on the ratio of the acidic and basic component concentrations and KA. MCj00787110000[1] The logarithmic form of that relation is known as Henderson-Hasselbalch equation: • H+ ions of the strong acid are added to the solution: - concentration of H+ increases that upsets the equilibrium. the buffer base binds most of the added H+ ions which results in increase of the acidic buffer component New equilibrium will settle : CH3COO- + H+ CH3COOH [H+] increases proportionally to the increase of cacid / cbase, pH decreases proportionally to the decrease of the log cbase / cacid . How does the buffer work ? • OH- ions of the strong hydroxide are added to the solution CH3COOH + OH- CH3COO- + H2O Increase in OH– concentration withdraws H+ from the buffer acid that transforms into its conjugate base. CH3COO- + H+ CH3COOH [H+] decreases proportionally to the decrease of cacid / cbase, pH increases proportionally to the increase of the log cbase / cacid . pH of a buffer Henderson-Hasselbalch equation For the weak acid and its salt with a strong base from which Salt can be considered as conjugate base to the acid For weak base and its salt with the strong acid Henderson-Hasselbalch equation in general form: where cbase is the concentration of basic component and cAcid the concentration of acidic component of conjugate pairs that form the buffer. pH of a buffer depends on pKA value ratio of buffer components if cB/cA = 1 Capacity of a buffer – expresses the effectivity of a buffer Capacity is highest if cB/cA = 1 In sufficient buffer solutions, the ratio cbase/cacid should take values from 1:10 to 10:1, i.e. in the range pKA ± 1. Capacity depends also on the total buffer concentration Buffer systems in human body The pH value of blood is 7.40 ± 0.04 . Most biological happenings occur in the pH range 6 to 8. Blood buffer bases: Buffer: Hydrogen carbonate HCO3– / (H2CO3+CO2) Plasma proteins protein / protein-H+ Haemoglobin of red blood cells haemoglobin / haemoglobin-H+ Hydrogen phosphate HPO42– / H2PO4– All those buffer systems cooperate – a surplus of H+ is accepted by all buffer bases but distributed proportinally to their concentration in blood. Each of those four buffer systems has its own pKA . Hydrogencarbonate buffer CO2 + H2O D H2CO3 D H+ + HCO3- [CO2 + H2CO3] = [H2CO3]ef effective concentration It is proportional to the partial pressure of CO2 in blood •CO2 originates from metabolism •CO2 dissolves in water and its small part forms H2CO3 • The concentration of H2CO3 depends on concentration of CO2 • Instead the concentration [H2CO3] is used effective concntration [H2CO3]eff (water, 25 °C) pKAef = 6,37 In blood (t =37 oC, higher ionic stregth) pKAeff = 6,10 KA of carbonic acid is replaced by KAef Henderson-Hasselbalch equation for hydrogencarbonate buffer in blood parcial pressure of CO2 in kPa For coefficient 0,23 and pressure in kPa is expressed in mmol/l !!!!!! coeficient of solubility How does hydrogencarbonate buffer function? H+ H+ + HCO3- H2CO3 CO2 + H2O CO2 + H2O H2CO3 H+ + HCO3- lung OH- OH- + H2CO3 HCO3- + H2O CO2 + H2O Open system – the amount of CO2 may be regulated by ventilation kidney Hydrogencarbonate buffer functions as open buffering system Concentrations of both components can be regulated : CO2 by respiration HCO3- by function of liver and kidney Amino acid Ionizable group in the side chain pKA Aspartate Glutamate Histidine Cysteine Tyrosine Lysine Arginine b-carboxyl (-COOH) g-carboxyl (-COOH) imidazolium sulfanyl (-SH) phenolic hydroxyl e-ammonium (-NH3+) guanidinium –NH-C(NH2)=NH2+ 3,9 4,3 6,0 8,3 10,1 10,5 12,5 Plasma proteins and haemoglobin as buffers In all proteins, only ionizable groups can take part in acid-base reactions. At physiological pH values, imidazol groups of histidine residues alone act as effective buffer bases. The most important blood protein buffer is hemoglobin N N C H 2 - C H - C O O H N H 3 + H+ OH- Dissociation (protonization) of histidine base conjugate acid H2PO4- - acidic component HPO42- - bazic component pKA2 = 6,8 Buffer is of second-rate significance in the blood due to relatively low concentration. However, within the cells, phosphates with proteins are the major buffer bases. Hydrogen/dihydrogenphosphate buffer