Acid and Base Equilibrium

 

Acid/Base Definitions

Conjugates

Strong/Weak Acids and Bases

Autoionization of Water

pH

 

5 out of 10 top 10 industrial chemicals are acids or bases.  Many reactions important in industry and in life processes involve acid-base reactions.  Our bodies maintain an acid concentration in our blood between 4 X 10-8 M and 5 X 10-8 M.  If the acid concentration is outside of this range, death can occur. 

 

We will see there is a balance of acid and base concentration in aqueous solutions.  Since all life that we know of uses water, this balance is important to understand.  The H+ ion participates in the equilibria of respiration, so changes in its concentration will affect the ability of hemoglobin to carry oxygen.

 

Acids and Bases can act as catalysts and so have a big effect on reaction rates.  Many enzymes activity vary with acid and base concentration.

 

Acid Properties – sour taste, reacts with many metals to produce a salt and H2

Base Properties – bitter taste, “soapy” feel, neutralize acids

 

Arrhenius Acid – Substance that produces H+ in water            

1)         HCl(g) ® H+(aq) + Cl-(aq)

           

            Base – Substance that produces OH- in water     

2)MOH(s) ® M+(aq) + OH-(aq)

 

Separately in 1923 Johannes Bronstead & Thomas Lowry broadened the definition of acids and bases.

 

            Acid – substance that can transfer H+ to a base, (H+ donor)

            Base – substance that can accept H+ from an acid (H+ acceptor)

 

3)         HCl(g) +           H2O(L)  ®          H3O+(aq) +        Cl-(aq)

Acid                Base               Hydronium ion

            Note H3O+ is a more accurate representation of an acid in water than H+.  Protonated water clusters such as H(H2O)2+ or H(H2O)4+ are even more accurate but also more complicated.

            Chemist use H+(aq) and H3O+(aq) interchangeably for an acid in water.

 

Bronstead Lowry definition not limited to aqueous solutions.

 

4)         HCl(g) + NH3 (g) ® NH4+Cl- (s)

acid     base

 

Ammonia in water is an example of a base as a proton acceptor.

 

5)         NH3 (aq)  +        H2O = NH4+  + OH-

base               acid

 

Note water reacts with NH3 like an acid and with HCl like a base.  An Amphoteric substance can act like both an acid and a base.

 

In the reverse of reaction 5 above, the NH4+ acts like an acid and the hydroxide (OH-) acts like a base.   The NH4+/ NH3 are called a conjugate pair with the NH4+ being the conjugate acid of NH3.  NH3 is the conjugate base of NH4+.  The difference between conjugates is one H+, with the acid having one more H+.

 

            NH3 (aq)  +        H2O = NH4+  + OH-

 

 

What is the conjugate base of  H2PO4-?__________________

 

What is the conjugate acid of H2PO4-?______________________

 

See table page 733 (MSJ) about acid / base strength

 

 

Conjugate Acid

Conjugate Base

 

Strong Acids

HCl

H2SO4

HNO3

Cl-

HSO4-

NO3-

 

Negligible Base

 

H3O+

H2O

 

 

Weak Acid

HSO4-

HF

HC2H3O2

H2PO4-

HPO4-2

SO4-2

F-

C2H3O2-

HPO4-2

PO4-3

 

Weak Base

 

H2O

OH-

 

 

Negligible Acid

OH-

H2

CH4

O2-

H-

CH3-

Strong Base

 

Acids stronger than hydronium (H3O+) doesn’t exist in water, they react with water to produce hydronium.  These reactions have a very large Kc.  And a base stronger that OH- doesn’t exist in water, they react with water to produce hydroxide.  For example, when sodium oxide is added to water, the following reaction happens,

 

            Na2O (s) + H2O (L) → 2 Na+ (aq) + 2 OH- (aq)

 

The conjugate base of a strong acid is a negligible base (e.g. Cl-) and the conjugate acid of a strong base is a negligible acid (e.g. conjugate acid of H- ­is H2, which is usually not considered an acid).

 

 

 

Autoionization of water

 

Water can act like an acid and a base

 

            Show Lewis structures of 2 waters exchanging a proton to produce hydronium and hydroxide

 

            2 H2O = H3O+ + OH-

 

            Kc = [H3O+] [OH-] / [H2O]2

 

            Kc [H2O]2  = Kw = [H3O+] [OH-]  This is called the ion product of water.

 

At 25°C, Kw = 1.0 X 10-14.

This equilibrium works at all different conc. of AB.

In pure water at 25°C,  [H+] = [OH-]

            Kw = x2            => x = 1.0 X 10-7 = [H+] = [OH-]

 

In 0.1 M HCl (a strong electrolyte meaning 100% ionization), [H+] = 0.1 M = 1. X 10-1.  What is the [OH-] ?

 

            Kw = 1.0 X 10-14 = .1 [OH-]    ,  [OH-] = 1.0 X 10-13 M

 

In 0.1 M NaOH (another strong electrolyte), [OH‑] = 0.1 M = 1. X 10-1.  What is the [H+] ?

 

            Kw = 1.0 X 10-14 = .1 [H+]      ,  [H+] = 1.0 X 10-13 M

 

And for laughs, consider a 1.0 X 10-3M HCl solution

            [H+] = 1 X 10-3 ,  [OH-] = 1X 10-10M

So [H+] varies enormously in water and it is usually small, e.g. in blood 4.0 X 10-8M < [H+] < 5 X 10-8 M.  To avoid writing all the exponents we use pH

 

            pH = - log [H+]

 

            pOH = - log [OH-] (used less frequently)

 

Remember

            Kw  = [H3O+] [OH-] 

 

            Log Kw = Log [H3O+] + Log [OH-]

 

             -Log Kw = - Log [H3O+] + - Log [OH-]

 

            14          =  pH  + pOH

 

            Using above numbers

           

[H+] / M

[OH-] / M

pH

pOH

10-7

10-7

7

7

.1 (10-1)

1 X 10-13

1

13

1 X 10-13

.1 (10-1)

13

1

1 X 10-3

1 X 10-10

3

10

 

Note the 2nd and 4th solutions have [H+] > [OH-], and have a pH < 7.  These are acidic solutions.  The third solution has a [OH-] > [H+] and a pH > 7.  This is a basic solution.