Acids and Bases Biotechnology I

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

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

Water  Water = H 2 O  H + + OH -  Pure water at 25 C Concentration of H + = concentration of OH - [1 x mole/L]  Aqueous = water based  H + is the symbol for hydrogen ion  OH - is the symbol for the hydroxide ion

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

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 H 2 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 -

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 -  The OH - ions react with H + to form water, thereby  the concentration of hydrogen ions

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.

Neutralization Reaction  One mole of H + from an acid combines with one mole of OH - from a base to form H 2 O.  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 - H NaCl

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?

Quiz  What is OH - ?  What is the pH of a solution w/ an H + ion concentration of mole/L?  What is the concentration of H + ions in a solution w/ a pH of 9.0?

Answers  Hydroxide ion  pH = -log [H + ] = -log = -(-4) = 4  pH = -log[H + ]; 9.0 = -log [H + ] -9.0 = log [H + ] antilog (-9.0) = 1 x mole/L

Review Questions  Which pH value describes the most acidic solution?  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

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

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

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

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?

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.

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