1. BUFFERS Maintaining a stable pH when studying enzymes is a necessity! Buffer is a mixture of a weak acid (base) and a salt of this acid (base) Acid buffer HAn + KatAn Base buffer KatOH + KatAn

2. Requirements for buffers in biochemistry - Must be non biologically active - Must not interfere with the function of the enzyme being studied Good et al. (1966) developed a series of N-substituted taurine and glycine buffers. Those are non-toxic to cell at concentration 50 mM

3. Trivial nameAcid/Base namepKapKa Phosphate (pK 1 )―2.15 Acetate―4.76 MES2-(N-morpholino)ethanesulfonic acid6.10 Bis-Tris[Bis(2-hydroxyethyl)imino] tris(hydroxymethyl)methane 6.46 TES2-[Tris(hydroxymethyl)methylamino] ethanesulfonic acid 7.40 TrisTris(hydroxymethyl)aminomethane8.06 Piperidine―11.12 Dissociation constant of main buffer acids and bases

4. Buffer capacity Value that shows a quantity of a strong base or a strong acid, which must be added to 1 L of a buffer solution to change its pH by 1. HA ↔H + + A -, K a acidic buffer system BOH ↔ B + + OH - H + + OH - ↔ H 2 O, K w base contaminant

5. Assumption: [A - ] + [HA] = constant c – buffer concentration

6. Dependence of buffer capacity on pH β; c = 1M β; c = 0.1M pK a = 4 pK a = 8

7. Measurement of pH 1.Colorimetric method 2.Potentiometric method

8. Colorimetry The method is based on ability of some organic bases or acids to change color depending on pH as a result of the rearrangement of their structure Paper indicator liquid indicator pH: 1-3 5-6 7 8-9 11-13

9. IndicatorLow pH color Transition pH range High pH color Gentian violetyellow0.0–2.0blue-violet Methyl yellowred2.9–4.0yellow Bromophenol blueyellow3.0–4.6purple Congo redblue-violet3.0–5.0red Methyl orangered3.1–4.4orange Bromocresol greenyellow3.8–5.4blue Bromocresol purple yellow5.2–6.8purple Neutral redred6.8–8.0yellow Naphtholphthaleincolorless to reddish7.3–8.7greenish to blue Phenolphthaleincolorless8.3–10.0fuchsia Alizarine Yellow Ryellow10.2–12.0red Color indicators

10. Potentiometry Glass pH-sensitive electrode Measuring cell

11. Combined pH-selective electrode

12. Nernst Equation E – e.m.f. E 0 – standard e.m.f. z - charge of the ion, z H = 1 F – Faraday constant, 96500 Q

13. Experimental work 1.Prepare 0.1 solution of acetic buffer 2.Prepare 0.1 M solution of acetic acid 3.Prepare 1 M solution of a strong acid 4.Measure pH of the solutions