Download Arrhenius Theory The Arrhenius theory, named after Swedish

Arrhenius Theory
The Arrhenius theory, named after Swedish physicist Svante August Arrhenius, views
an acid as a substance that increases the concentration of the hydronium ion (H3O+) in
an aqueous solution and a base as a substance that increases the hydroxide ion (OH−)
concentration. Well-known acids include hydrochloric acid (HCl), sulfuric acid (H2SO4),
nitric acid (HNO3), and acetic acid (CH3COOH). Bases include such common
substances as sodium hydroxide (NaOH) and calcium hydroxide (Ca(OH)2). Another
common base is ammonia (NH3), which reacts with water to give a basic solution
according to the following balanced equation.
NH3(aq) + H2O NH4+(aq) + OH−(aq)
(Reaction occurs to small extent; hydroxide ion concentration is small but measurable.)
A large number of natural bases are known, including morphine, cocaine, nicotine, and
caffeine; many synthetic drugs are also bases. All of these contain a nitrogen atom
bonded to three other groups, and all behave similarly to ammonia in that they can react
with water to give a solution containing the hydroxide ion.
Amino acids, a very important class of compounds, are able to function both as acids
and as bases. Amino acid molecules contain both acidic (−COOH) and basic (−NH2)
sites. In an aqueous solution, amino acids exist in both the molecular form and the socalled "zwitterionic" form, H3N + CH2CO2−. In this structure the nitrogen atom bears a
positive charge, and the oxygen atom of the acid group bears a negative charge.
According to the Arrhenius theory, acid-base reactions involve the combination of the
hydrogen ion (H+) and the hydroxide ion to form water. An example is the reaction of
aqueous solutions of sodium hydroxide and hydrochloric acid.
HCl(aq) + NaOH(aq) NaCl(aq) + H2O
Brønsted-Lowry Theory
A somewhat more general acid-base theory, the Brønsted-Lowry theory, named after
Danish chemist Johannes Nicolaus Brønsted and English chemist Thomas Martin
Lowry, defines an acid as a proton donor and a base as a proton acceptor. In this
theory, the reaction of an acid and base is represented as an equilibrium reaction.
acid (1) + base (2) base (1) + acid (2)
( indicates that the products can re-form the reactants.)
Acid (1) and base (1) are called a conjugate acid-base pair, as are acid (2) and base
(2). The advantage of this theory is its predictive capacity. Whether the equilibrium lies
toward the reactants (reactant-favored) or the products (product-favored) is determined
by the relative strengths of the acids and bases.
The Brønsted-Lowry theory is often closely associated with the solvent water.
Dissolving an acid in water to form the hydronium ion and the anion of the acid is an
acid-base reaction. Acids are classified as strong or weak, depending on whether the
equilibrium favors the reactants or products. Hydrochloric acid, a strong acid, ionizes
completely in water to form the hydronium (H3O+) and chlorine (Cl−) ions in a productfavored reaction.
HCl(aq) + H2O H3O+(aq) + Cl−(aq)
Using the Brønsted-Lowry theory, the reaction of ammonia and hydrochloric acid in
water is represented by the following equation:
NH3(aq) + HCl(aq) NH4+(aq) + Cl−(aq)
Hydrochloric acid and the chlorine ion are one conjugate acid-base pair, and the
ammonium ion and ammonia are the other. The acid-base reaction is the transfer of the
hydrogen ion from the acid (HCl) to the base (NH3). The equilibrium favors the weaker
acid and base, in this case the products. Note that the hydroxide ion does not appear in
this equation, a point differentiating the Arrhenius and Brønsted-Lowry theories.
Lewis Theory
A still broader acid and base theory was proposed by American physical chemist Gilbert
Newton Lewis. In the Lewis theory, bases are defined as electron-pair donors and acids
as electron-pair acceptors. Acid-base reactions involve the combination of the Lewis
acid and base through sharing of the base’s electron pair.
Ammonia is an example of a Lewis base. A pair of electrons located on the nitrogen
atom may be used to form a chemical bond to a Lewis acid such as boron trifluoride
(BF3). (In the following equation, the colon represents an electron pair.)
H3N: + BF3 H3N→BF3
Ammonia, water, and many other Lewis bases react with metal ions to form a group of
species known as coordination compounds. The reaction to form these species is
another example of a Lewis acid-base reaction. For example, the light blue color of a
solution of Cu2+ ions in water is due to the [Cu(H2O)6]2+ ion. If ammonia is added to this
solution, the water molecules attached to copper are replaced by ammonia molecules,
and the beautiful deep blue ion [Cu(NH3)4]2+ is formed.
Acid–Base Neutralization Reactions
When a Strong Acid and a Strong Base are mixed in solution, a Neutralization Reaction
occurs. The products do not have characteristics of either acids or bases. Instead, a
Neutral Salt and water are formed. Look at the reaction below:
HCl + NaOH NaCl + H2O
The anion from the acid (Cl–) reacts with the cation from the base (Na+) to produce
Sodium Chloride Salt and water. A salt is defined as any compound formed whose
anion came from an acid and whose cation came from a base.
Diprotic Acid
Diprotic Acids contain two Hydrogen Ions per molecule capable of dissociating in water.
The dissociation does not happen all at once due to the two stages of dissociation
having different Ka values. The first dissociation will, in the case of Sulfuric Acid, occur
completely, but the second one will not.
H2SO4 H+(aq) + HSO4−(aq) Ka = 1 × 103
HSO4− H+(aq) + SO42−(aq) Ka = 1 × 10−2
During titration, the curve will clearly show two equivalence points for the acid. This
occurs because the two Hydrogen Ions do not leave the acid at the same time.
Diprotic Base
Diprotic Bases, such as Nicotine, can accept two Hydrogen Ions.
Nicotine Protonation with Acid
“N2+ will predominate in solutions more acidic than a pH of 3, the monocation N1+ will
predominate in solutions whose pH lies between 3 and 8, and the free base, N, will
predominate in solutions that are more basic than a pH of 8.”
(The Acid–Base Chemistry of Nicotine: Extensions, Analogies, and a Generalization)