Acid/Base Theory and Amino Acids 

 

 

I. Classification of substances in aqueous solutions

A. Acids

B. Bases

 C. Salt

1. increase H3O+ concentrations

2. taste sour

3. pH < 7

4. turn blue litmus red

5. react with active metals producing H2

6. react with carbonates/bicarbonates to produce CO2

1. Increase OH- concentrations

2. Taste bitter

3. pH > 7

4. Turn red litmus blue

5. Have slippery feel

 

 

1. ionic substance formed from the neutralization reaction of an acid and a base

2. pH ~ 7 (dependent upon strength of acid and base). See salt hydrolysis

 

 

  A. Definitions of Acids & Bases

    1. Svante Arrhenius,1883. He won the Nobel prize in 1903 for his work with electrolytic dissociation

      a. Acid- Substances that increase H+ concentrations in solution

      b. Base- Substances that increase OH- concentrations in solution.

 

    2. Brønsted-Lowry. Johannes Brønsted &  Thomas Lowry. Paper by Brønsted (1923)
      a. Acid- Substances that donate H+ to another

      b. Base- Substance that accepts H+ to another

            

         

        Many acid/base reactions can exist in equilibrium, so the products form acid/base pairing

      c. Conjugate Acid- An acid that was created from a base

      d. Conjugate Base- A base that was created from an acid

   

    3. G. N. Lewis. Lewis Acid/Base concept

      a. Acid- Substances that accept e- pairs. Often identified as an electrophile

      b. Base- Substances that donate e- pairs.  Often identified as a nucleophile

            

      c. Products of Lewis acid-base chemistry are called acid/base adducts. Created through coordinate covalent bonding.

 

    4. Naming Acids

      a. Binary acids- Acids containing a single anion bonded to a hydrogen.

           ex. HCl(aq), HBr(aq), HF(aq), HI(aq), H2S(aq)

        1. Naming- Use hydro- prefix and change anion name ending -ide to -ic acid

            ex. HCl - hydrochloric acid

                 HBr- hydrobromic acid

 

      b. Ternary acid- Acids containing polyatomic anions bonded to a hydrogen

           ex. H2SO4(aq), HNO3(aq), H3PO4(aq), HC2H3O2(aq)

        1. Naming- DO NOT Use hydro prefix.  Change anion endings -ate to -ic acid and -ite to -ous acid.

            ex. HNO3(aq) - nitric acid                H2SO4(aq) -sulfuric acid                   

                 HNO2(aq) -nitrous acid                H2SO3(aq) - sulfurous acid

 

 B. Self-ionization of water

     Ionization- covalent molecules which form ions by gaining or losing protons (H+)

     Self-ionization (auto-ionization) of water- Water molecules ionize to form hydroxide and hydronium ions

             

    1. Ion-Product Constant for water

          approximately 1 out of 550,000,000 molecules ionize

       So from the Equilibrium equation:        =  1.8 x 10-16

     

    -but [H2O] ~ 55.5 M, therefore:  Keq [H2O] = Kw  = [H3O+]  [OH-] = 1.0 x 10-14 M at 25oC. 

 

    - at 25oC,  [H3O+] =  [OH-] = 1.0 x 10-7 M

 

  C.     pH & pOH scales- measures the relative strength of an acid or base. 

            Paper by S.P.L. Sörenson, Biochemische Zeitschrift (1909)

        

    1. Identifying pH and pOH       

           pH = -log [H3O+]      and     pOH = -log [OH-]

       The letter p is derived from the German word potenz meaning power or exponent of, in this case, 10

 

Concentrations and pH of various solutions

Solution type     

[H3O+] M

[OH-] M

 pH Value

Acidic 

> 1.0 x 10-7

 < 1.0 x 10-7

< 7.00

 Neutral  

= 1.0 x 10-7

= 1.0 x 10-7

 = 7.00

Basic

  < 1.0 x 10-7  

 > 1.0 x 10-7

> 7.00

 

       - In any solution, the concentrations of  [H3O+] &  [OH-] are inversely proportional

         

[H3O+]  [OH-] = 1.0 x 10-14

-log [H3O+]  [OH-] = -log ( 1.0 x 10-14)

pH  + pOH  = 14

Animation:  Relating pH to [H+]

Practice: Assessing pH of common materials

   

    2. Measuring pH & Acid/Base Indicators

      a. pH meters- utilize electrical potentials within a solution to measure pH. 

 

      b. Indicators:  Acid-Base indicators are dyes that are themselves weak acids and bases.  However, the conjugate 

          acid-base forms of the    dye have different colors.  The actual chemical structures of the dyes is often quite complex; 

          however, we can use the generic symbol for the indicator as HIn.  The Brönsted-Lowry equation for the indicator is:

                                           

         -Adding acid (increasing [H3O+]) would shift the equilibrium back to HIn 

         -Adding base (decreasing [H3O+]) would shift the equilibrium to In-

 

Activity: Making your own pH indicator from cabbage. Anthocyanins in cabbage produces the purple color

 

  D. Strength of Acid or Base

      Dissociation constants (Ka & Kb)- measures the amount of ionization of an acid or base.

   

    1. For strong electrolytes: ( like strong acids)

             ex. 

 

             So    >> 1 (much larger than one)

                                                 

           So… this says that in solution, the prevalent forms are  H3O+  and Cl-

         

       a. Calculating pH.  Strong acids are defined as those that completely ionize (or nearly).

          The [ H3O+ ] is invariably equal to [ HA ].

          So if pH = -log [ H3O+ ] then for a strong acid the pH = -log [ HA ]

 

          ex. Find the pH of a 0.015 M HCl solution

             pH = -log [ H3O+ ]  =  -log [ HCl ]  =  -log (0.015)  =  1.82

          ** The number of sig. figs. in the concentration is equal to the number of decimal places right of the decimal

 

      2. Weak Electrolytes.  Do they strongly dissociate?

          ex.  Acetic Acid:      CH3COOH  +  H2O   --> CH3COO-    +  H3O+

                          

           but with weak electrolytes, [H2O] remains very large and almost unchanged. IGNORE IT

                                         

                Keq  [H2O]   =      =    1.74 x 10-5 at 25oC

                                                                     

         a. Calculating pH. Weak acids only partially ionize so [ H3O+ ] is not equal to [ HA ].

           -To calculate the pH of weak acid or base, you must consider the Ka for the acid & Kb for the base

 

 For    HA    +       B    -->        A-      +       HB+

When B is H2O:

  Keq [H2O]  = 

When HA is H2O :

  Keq [H2O] =  

                       

table: Strength of Acids/Bases with conjugates.

table: Ka & Kb values for common acids & bases

           

          ex. Calculate the pH of 0.015 M Acetic Acid (CH3COOH). Construct a table to include concentrations at two

            points: initial (pre-ionization, [ ]o) and final (post-ionization, [ ]i) along with changes in concentrations

         

CH3COOH  +  H2O   --> CH3COO-    +  H3O+
  CH3COOH  H2O CH3COO- H3O+
[ ]o  -initial conc. 0.015 - 0 0
D [ ] - change in conc. - x - + x + x
[ ]i -conc. after ionization 0.015 - x x x x

        

           Solution:   Ka  =  [CH3COO-] [H3O+]       ===>     1.74 x 10-5 =    (x  •  x)            

                                         [CH3COOH]                                                   (0.015 - x)

           --If the Ka << 1, assume the change in initial acid concentration is not significant, so  (0.015 - x)   ~ 0.015

 

             so  1.74 x 10-5 =    (x  •  x)     ===>   (1.74 x 10-5) (0.015)  =  x2   ===>  x = 5.11 x 10-4 M

                                                0.015                          

             pH  =  -log (5.11 x 10-4)  =  3.29

       

       b. 5% Rule- As long as a weak acid or base is less than 5% ionized, then the change in initial acid/base concentration can be

            ignored.  If the ionization is greater than 5%, then the change in initial acid/base concentration must be included.

 

       c. Henderson-Hasselbalch Equation.

        -re:  Ka is an equilibrium expression which relates the concentrations of unionized acid with concentrations of the conjugate

               base and hydronium.  This expression can be modified to calculate the pH of any acid/base solution knowing the

               concentrations of acid and conjugate base.

                  

Henderson-Hasselbalch Equation

 

 

 

Titrations of Amino Acids

         

                 

                 

            

What do  pKa1  & pKa2 mean?

          Where does the Isoelectric point exist?  What does it mean?

 

Solving for Isoelectric point

 

pI  =  pKa1  +  pKa2   

 

                             2

 

          For alanine, the pI equals

 

                   pI =  2.3 + 9.7  = 6

                                2

                  

          So, what does this mean?

 

 

 
 

   

          Calculate the pI for Glutamic Acid.