Lecture No. 4. Chemical Reaction Types

Lecture No. 4. Chemical Reaction Types

There are five general reaction types we will consider.

A. Dissolution-Precipitation

Examples: NaCl (s)--> Na⁺ (aq) + Cl^– (aq)

CaCO₃ (s)-->Ca²⁺ (aq) + CO₃^2– (aq)

The equilibria governing these reactions are described by solubility products, K_sp.

B. Acid-Base

In aqueous systems, the Brönsted-Lowry definition of an acid as a hydrogen-ion donor is usually sufficient. A base then is a hydrogen-ion acceptor — OH^–, NH₃, etc. In aqueous systems, a base ends up liberating a hydroxide ion, either directly or indirectly. The Lewis definition is more general: an acid is an electron-pair acceptor, and a base is an electron-pair donor.

Boric acid is, strictly speaking, not a Brönsted-Lowry acid, but a Lewis acid.

Examples: HCl --> H⁺ + Cl^–

CO₃^2– + H₂O --> HCO₃^– + OH^–

HCO₃^– + H₂O --> H₂CO₃ + OH^–

"Carbonic acid" or "H₂CO₃", is a shorthand term for hydrated carbon dioxide. The H₂CO₃ molecule does exist under certain circumstances (in the gas phase), but the structure of the actual carbonic acid (hydrated CO₂) is more complicated, probably involving multiple water molecules.

NH₃ + H₂O--> NH₄⁺ + OH^–

Ammonia acts as a proton acceptor, releasing a hydroxide ion from water. (It is sometimes written as the ionization of NH₄OH, but there is no evidence for the existence of such a molecule.)

H₃BO₃-->B(OH)₄^– + H⁺

Boric acid’s first ionization is not the release of a proton from the boric acid molecule, but rather the stripping of a hydroxide ion from water. The end result is the same: production of a proton. But boric acid acts as an electron-pair acceptor in taking an electron pair from hydroxide.

C. Combinations of (A) and (B)

Example: Ca²⁺ + HCO₃^–--> CaCO₃ + H⁺

This is a combination of a dissolution-precipitation and an acid-base one.

HCO₃^– (aq)--> H⁺ (aq) + CO₃^2– (aq)

and

Ca²⁺ (aq)+ CO₃^2–(aq) --> CaCO₃ (s)

D. Oxidation-Reduction (Redox) Reactions

Example: 4 Fe₃O₄ +O₂ --> 6 Fe₂O₃

2 Fe₃O₄ + ¸ O₂ --> 3 Fe₂O₃

Looking at this reaction in detail, we have two half-reactions: an oxidation (loss of electrons)

2 Fe₃O₄ +H₂O --> 3 Fe₂O₃ + 2 H⁺ + 2 e^–

and a reduction (gain of electrons)

¸ O₂ + 2 e^– + 2 H⁺ --> H₂O

Adding these half-reaction equations together gives

2 Fe₃O₄ + H₂O + ¸ O₂ + 2 H⁺ + 2 e^– --> 3 Fe₂O₃ + H₂O + 2 H⁺ + 2 e^–

Canceling out the species that are on both sides (and therefore show no net change) gives us the original equation:

2 Fe₃O₄ + ¸ O₂ --> 3 Fe₂O₃

Another example is the reduction of Fe³⁺ by carbohydrates in acid solution:

4 Fe³⁺ + CH₂O + H₂O --> 4 Fe²⁺ + CO₂ + 4 H⁺

It is a combination of two half-reactions:

Fe³⁺ + e^– --> Fe²⁺

and

CH₂O + H₂O --> CO₂ + 4 H⁺ + 4 e^–

Note that this is also a combination reaction involving ionization of water (an acid-base reaction) and oxidation of carbohydrate.

E. Surface Reactions

These are reactions, usually involving solutes, that take place on a surface and involve surface sites of a mineral crystal.(The symbol: ¼ represents three horizontal lines, the typical representation of attachment to a surface binding site.)

Example: ¼FeOH + H⁺ --> ¼FeOH₂⁺

¼FeOH₂⁺ + F^- --> ¼FeF + H₂O

Here at a hydroxide group on the surface of an iron oxyhydroxide is protonated, after which a fluoride ion displaces the water.

A fairly common conventional terminology has

¼S_c signifying a cation-binding surface site, and ¼S_a signifying an anion-binding surface site. These sites can therefore form surface complexes like ¼S_c–Ca²⁺ and ¼S_a–OCO₂^2–.

Reading Assignment:

Elder, John F., 1988, Metal biogeochemistry in surface-water systems: A review of principles and concepts, USGS Circular 1013; posted outside my office.

To previous lecture: Lecture No. 3. Chemical Fundamentals II.

To next lecture: Lecture No. 5. Measuring Metals in Sediments

To list of lecture notes