Transcript Acids and Bases

Acids and Bases
Biotechnology I
Life Chemistry
Based on water
 Cells contain 80-90% water
 Proper pH essential to ALL living
systems
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 Plants cannot live in poor pH soil
 Animals die if blood pH is abnormal
 Microorganisms need specific pH to grow &
multiply
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Maintaining proper pH is CRITICAL to
survival of Cells and Biological systems
pH Environments
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Biological and Industrial processes require
specific pH environments
 Food processing
 Water purification
 Rx production
 Sewage treatment
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Requires pH monitoring
Water
Water = H2O  H+ + OH Pure water at 25 C
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 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
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pH is
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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
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Pure water
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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
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Definition: electrolyte that donates hydrogen ions
Properties:
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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-
Bases
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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.
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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.
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H+Cl- + Na+OH
-
H20 + NaCl
Logarithmic Scale
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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
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What is OH- ?
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What is the pH of a solution w/ an H+ ion
concentration of 10-4 mole/L?
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What is the concentration of H+ ions in a
solution w/ a pH of 9.0?
Answers
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Hydroxide ion
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pH = -log [H+] = -log 10-4 = -(-4) = 4
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pH = -log[H+]; 9.0 = -log [H+] -9.0 = log
[H+]
antilog (-9.0) = 1 x 10-9 mole/L
Review Questions
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Which pH value describes the most acidic
solution?
4
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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
Measuring pH
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Indicators
 Phenophthalein, phenol red, bromothymol
blue, universal indicator to name a few
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pH Paper
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pH Meters
pH Meter
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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
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Calibration recommended with each use, when
battery replaced and when fluid in sensor is
changed
Calibration
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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
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 For example, if you are running an enzyme
assay at 37C then adjust the pH at 37C
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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.
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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
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