Bronsted-Lowry
Acid – Base Reactions
Chemistry
Bronsted – Lowry Acid
Defined as a molecule or ion that is a
hydrogen ion donor
 Also known as a proton donor because H+
is a proton.
 Examples
 HCl – a “monoprotic” acid
 H2SO4 – a “diprotic” acid
 H3PO4 – a “triprotic” acid
 These all dissociate in water to produce
H+ ions which can then react with a
Bronsted-Lowry Base

Bronsted-Lowry Base
Defined as a hydrogen ion acceptor.
 In an acid-base reaction the base
“accepts” the hydrogen ion from the acid.
 NH3 + H+  NH4+

NH3 is Bronsted-Lowry Base
Conjugate Acid
The compound that is formed when the
B/L base gains a proton (H+).
 NH4+ is the conjugate acid of NH3 ( a
base)
 NH4+ acts as an acid in the reverse
reaction.

Conjugate Base
The compound that is formed when a B/L
acid gives away it’s proton (H+)
 Cl- is the conjugate base of HCl (an acid)
 Cl- acts as a base in the reverse reaction

The conjugate acid and bases are
the acid and bases for the reverse
reaction!
Amphoteric Definition: A substance that can act as an
acid (proton donor) or a base (proton
acceptor)
 Water (H2O) is amphoteric.
Water acts as a
 HBr + H2O  Br- + H3O+

base!

NH3 + H2O  NH4+ + OHWater acts as an
acid!
Strong vs. Weak Acids and
Bases
Strong - Dissociate 100%
 The following are the strong acids, all the
rest are considered weak. (Memorize!)
 HCl – Hydrochloric acid
 HBr – Hydrobromic acid
 HI – Hydroiodic acid
 HNO3 – nitric acid
 H2SO4 – sulfuric acid
 HClO4 – perchloric acid

Strong Bases
Strong bases are soluble hydroxides such
as NaOH, KOH, and LiOH
 All the rest are considered weak.

Weak – Only dissociate < 10% , They
are equilibrium reactions.
 We use Equilibrium constants to compare
weak acid strength.
 Strong Acids and Bases have very large
equilibrium constants Keq
 Weak Acids and Bases have small
equilibrium constants (<< 1)

The equilibrium constant for an
acid is called Ka
 HA
H+ + A-
The equilibrium constant for a base
is called the Kb
 NH3
+ H2O
NH4+ + OH-
The larger the Ka value – the stronger the acid!
Examples

Acid
Conjugate Base

HNO3
NO3-1
_____________
 HCO3-1
2CO
_____________
3
 HPO42-
3PO
_____________
4
*Conjugate base is what is left after acid donates
it’s H+ ion

Base
Conjugate Acid

HCO3-1
H2CO3
______________

F
HF
______________

-1
H2O
 HPO42-
+
H
O
______________
3
1H
PO
2
4
______________
* Conjugate acid is what you get after the base
accepts the H+ ion.
Ka’s are used to predict which species will act
as an acid and which as a base in an acid
base reaction. It also allows you to predict
which way an equilibrium reaction is favored.
(with the products or reactants)
Remember:
 The reactant with the higher Ka acts as the
acid.
 Compare the acid and the conjugate acid
Ka’s, the one with the higher Ka wants to
dissociate more, so the equilibrium will favor
the reaction direction that allows it to
dissociate!

Example 1
NH4+ + OH-

NH3 + H2O

Compare Ka values to determine which acts as
an acid. Higher Ka value on chart is acid!
Write the products (remove H+ from acid and
add it to the base)
Compare the Ka values of the acid and
conjugate acid.
Higher Ka wants to dissociate more, so the
equilibrium will be favored away from that acid.
NH4+ has higher Ka so it wants to dissociate
more, so Reactant side is favored!




Example 2

HCO3- + HPO42-
Example 3

HS- + H2CO3