1. LECTURE 2 THEME: Acid-base equilibrium in biological systems. Buffer solutions. ass. prof. Ye. B. Dmukhalska

2. Electrolytes in Aqueous Solution Electrolytes are substances such as NaCl or KBr, which dissolve in water to produce conducting solutions of ions. Nonelectrolytes are substances such as sucrose or ethyl alcohol, which do not produce ions in aqueous solution.

3. Anode Cathode (negative electrode) Solution of NaCl

4. Anode Solution of NaCl Cathode

6. According to the Degree of dissociation (α) electrolytes can be classified into the following: strong electrolytes are compounds that dissociate to a large extent (α> 30%) into ions when dissolved in water. For example, HCl, H 2 SO 4, HNO 3, HJ, NaOH, KOH, KCl. Medium electrolytes α = 2 - 30%. For istance, H 3 PO 4, H 3 PO 3. weak electrolytes are compounds that dissociate to only a small extent α<2%. For example, NH 4 OH, H 2 S, HCN, H 2 CO 3. nonelectrolytes α = 0 are compounds that don’t dissociate when dissolved in water.

7. Acid–Base Concept 1)The Arrhenius theory ACID A substance that provides H+ ions in water BASE A substance that provides OH- ions in water

8. 2) The Brønsted-Lowry Theory All Brønsted–Lowry bases have one or more lone pairs of electrons:

10. 3) The Lewis Acids and Base theory LEWIS ACID An electron-pair acceptor LEWIS BASE An electron-pair donor

11. Dissociation of Water Water acts as both an acid and a base at the same time: This reaction is characterized by the equilibrium equation Kw = [H 3 O+] [OH - ] The equilibrium constant Kw is called the ion-product constant for water.

12. The pH of a solution is defined as the negative base-10 logarithm (log) of the molar hydronium H3O+ ion concentration: 5. Equilibria in Solutions of Weak Acids The dissociation of a weak acid in water is characterized by an equilibrium equation. The equilibrium constant for the dissociation reaction, denoted Ka is called the acid-dissociation constant: Values of Ka and pKa = - log Ka for some typical weak acids are listed in Table The larger the value of Ka, the stronger the acid

13. Equilibria in Solutions of Weak Bases Weak bases, such as ammonia, accept a proton from water to give the conjugate acid of the base and OH- ions: The equilibrium constant Kb is called the base-dissociation constant: Relation Between Ka and Kb

14. BUFFERS Buffers are solutions which can resist changes in pH by addition of acid or alkali.

15. the Henderson-Hasselbalch equation. for acidic buffer solution [А - ] рН = pК a + log ---------- [АH] -for basic buffer solution [В] рН = 14 - pК b + log ----------- [HВ + ]

16. Buffers are mainly of two types: (а) mixtures of weak acids with their salt with а strong base (b) mixtures of weak bases with their salt with а strong acid. А few examples are given below: Н 2 СО 3 / NаНСО 3 (Bicarbonate buffer; carbonic acid and sodium bicarbonate) СН 3 СООН / СН 3 СОО Na (Acetate buffer; acetic acid and sodium acetate) Na 2 HPO 4 / NaH 2 PO 4 (Phosphate buffer)

17. Composition of buffer solutions by Brønsted– Lowry theory : (a) weak acid with their conjugate base Н 2 СО 3 / НСО 3 - Bicarbonate buffer solution СН 3 СООН / СН 3 СОО - -Acetate buffer solution H 2 PO 4 - / HPO 4 2- - Phosphate buffer solution (b) weak base with their conjugate acid NН 3 / NH 4 + - ammonia

18. Buffer solution of blood HHb/Hb - - Hemoglobin buffer solution HHbO 2 /HbO 2 - - Oxihemoglobin buffer solution H 2 PO 4 - /HPO 4 2- - Phosphate buffer solution H 2 CO 3 /HCO 3 - ; - Bicarbonate buffer solution PtCOOH / PtCOO - or NH 2 -R-COOH / NH 2 -R-COO - - Protein buffer solution

19. Factors Affecting pH of а Buffer The pH of а buffer solution is determined by two factors: 1. The value of pK: The lower the value of pK, the lower is the pH of the solution. 2. The ratio of salt to acid concentrations: Actual concentrations of salt and acid in а buffer solution may be varied widely, with по change in рН, so long as the ratio of the concentrations remains the same.

20. Buffer Capacity On the other hand, the buffer capacity is determined by the actual concentrations of salt and acid present, as well as by their ratio. Buffering capacity is the number of grams of strong acid or alkali which is necessary for а change in pH of one unit of one litre of buffer solution. The buffering capacity of а buffer is, definеd аs the ability of the buffer to resist changes in pH when an acid or base is added.

21. Buffers Act When hydrochloric acid is added to the acetate buffer, the salt reacts with the acid forming the weak acid, acetic acid and its salt. Similarly when а base is added, the acid reacts with it forming salt and water. Thus, changes in the pH are minimised. СН 3 СООН + NaOH = СН 3 COONa + Н 2 О СН 3 СООNа + HCI = СН 3 СООН + NaCI The buffer capacity is determined by the absolute concentration of the salt and acid. But the рН of the buffer is dependent on the relative proportion of the salt and acid (see the Henderson - Hasselbalch's equation). When the ratio between salt and acid is 10:1, the pH will be one unit higher than the pKa. When the ratio between salt and acid is 1:10, the pH will be one unit lower than the pKa.

22. Mechanisms for Regulation of pH (1) Buffers of body fluids, (2) Respiratory system, (3) Renal excretion. These mechanisms are interrelated.