Acids and Bases

1. Acids and Bases Biotechnology I

2. Life Chemistry • Based on water • Cells contain 80-90% water • Proper pH essential to ALL living systems • Plants cannot live in poor pH soil • Animals die if blood pH is abnormal • Microorganisms need specific pH to grow & multiply • Maintaining proper pH is CRITICAL to survival of Cells and Biological systems

3. pH Environments • Biological and Industrial processes require specific pH environments • Food processing • Water purification • Rx production • Sewage treatment • Requires pH monitoring

4. Water • Water = H2O  H+ + OH- • Pure water at 25 C • Concentration of H+ = concentration of OH- [1 x 10-7 mole/L] • Aqueous = water based • H+ is the symbol for hydrogen ion • OH- is the symbol for the hydroxide ion

5. pH is • A way to express hydrogen ion concentration in a solution • Measurement of the acidity/alkalinity of an aqueous solution • pH is the –log of the H+ concentration • pH is measured on a scale • Ranges from 0 to 14 • Pure water • H+ concentration is 1x10-7 mole/L • The log of 1x10-7 = -7 • The – log of –7 = 7 • The pH of pure water = 7

6. Acids • Definition: electrolyte that donates hydrogen ions • Properties: • Acids in water conduct electricity • The stronger the acid the stronger the conductivity • Acids react w/metals to produce H2 gas • Acids are indicators; they cause reversible color changes • Phenolphthalein and litmus are two examples of acid-base indicators • Acids react w/hydroxide compounds to form water and salt; this type of reaction is called “neutralization” • Strong acids completely dissociate in water to release hydrogen ions = H+ • i.e. hydrochloric acid (HCl): HCL in water H+ + Cl-

7. Bases • Definition: electrolyte that yields hydroxide ions or accepts hydrogen ions • Properties: • Bases in water conduct electricity • The stronger the base the stronger the conductivity • Bases react with acids in neutralization reactions to form water and a salt • Bases cause reversible color changes in acid-base indicators (color is pH dependent) • Bases in water solution are slippery to the touch • Caution: even dilute bases can be caustic! • Strong bases completely dissociate in water to release hydroxide ions = OH- • NaOH in water  Na+ + OH- • TheOH- ions react with H + to form water, thereby  the concentration of hydrogen ions

8. Buffer  • Substance(s) that when in aqueous solution resists a change in H+ concentration even if acids or bases are added • Some buffers change pH as their temperature and/or concentration changes • Tris buffer is widely used in molecular biology; it is very sensitive to temperature and the pH will vary greatly at various temperatures.

9. Neutralization Reaction • One mole of H+ from an acid combines with one mole of OH- from a base to form H2O. • In addition, one mole of negative ions from the acid combine with one mole of positive ions from the base to form a salt. H+Cl- + Na+OH - H20 + NaCl

10. Logarithmic Scale • pH scale is logarithmic • Means each whole number increases by the factor of 10. • A solution with pH=6 is 10x more acidic than pure water with pH=7. • pH 5.0 has 10 x more H+ then pH of 6.0 • pH of 7.0 is 100 x less acidic than pH of 5.0 •  pH of 7.0 has 100 x less what then a solution with a pH of 5.0?

11. Quiz • What is OH- ? • What is the pH of a solution w/ an H+ ion concentration of 10-4 mole/L? • What is the concentration of H+ ions in a solution w/ a pH of 9.0?

12. Answers • Hydroxide ion • pH = -log [H+] = -log 10-4 = -(-4) = 4 • pH = -log[H+]; 9.0 = -log [H+] -9.0 = log [H+] antilog (-9.0) = 1 x 10-9 mole/L

13. Review Questions • Which pH value describes the most acidic solution? • 4 2 14 10 • What is one of the most common bases used in the lab? • Sodium Hydroxide • Describe it when it is in solution • Given what you know, what would you say about “Clorox” bleach? • It is slippery to the touch

14. Measuring pH • Indicators • Phenophthalein, phenol red, bromothymol blue, universal indicator to name a few • pH Paper • pH Meters

15. pH Meter • Meter / electrode system for measuring pH in laboratory • Provides greater accuracy, sensitivity than chemical indicators • Can measure pH of a solution to the nearest 0.1 unit • Can be used with variety of aqueous solutions • Consists of: • Voltmeter – measures voltage • Two electrodes connected to one another (sensor probe) • When immersed in the sample they develop an electrical voltage that is measured by the voltmeter • Calibration recommended with each use, when battery replaced and when fluid in sensor is changed

16. Calibration • Important in operating the pH meter • It tells the meter how to translate the voltage difference between the measuring and reference electrodes into units of pH • Temperature sensitive • Two buffers of known pH are used to calibrate a pH meter • Refer to pH meter manual

17. Adjusting the pH of a buffer • Most often you will adjust the pH using NaOH or HCL • Adjust the pH at the temperature it will be used at • For example, if you are running an enzyme assay at 37C then adjust the pH at 37C • When making a buffer, do not bring it to final volume until you have adjusted the pH. Why?

18. Adjusting the pH of a buffer • Place pH probe in solution • Check the pH and temperature • Add base or acid SLOWLY as required, soln. should be stirring • Re check pH to see if it is at specified pH.

19. Critical Tips for Using pH Meter • Depth of immersion – do not immerse to the bottom of a solution if there are particulates settled there • Make sure air bubbles are not trapped in the probe • Rinse probes w/ distilled water after each series of measurements • Be sure stir bars are not hitting the probe