Before you learn about the material in this section, you should probably master the acids and bases section.
Now it's time to refine our notion of what makes a compound acidic or basic. It turns out that things without the capability to donate protons (H+) or hydroxyl groups (OH-) can be quite acidic or basic.
In the section on acids and bases, you learned that acids are substances that elevate the concentration of protons in a solution, and a base reduces it. More specifically, Swiss chemist Svante Arrhenius formulated this definition:
An acid is a substance which dissociates in solution so that it liberates a proton (H+), and a base is a substance which dissociates in solution to liberate a hydroxyl ion (OH-).
An acid is a proton donor
A base is a hydroxyl donor
A little later, two chemists, Johannes Brønsted (Denmark) and Thomas Lowry (England) refined the definition of acids and bases by calling acids proton donors and bases proton acceptors.
Here are some examples. First, consider the strong acid, HCl
HCl ⟶ H+ + Cl-
When HCl dissociates, it "donates" a proton to the solution.
Now let's look at the strong base NaOH. When NaOH dissociates in solution, a free hydroxyl ion is liberated:
NaOH ⟶ Na+ + OH-
Now this hydroxyl is free to soak up a free proton in solution either because the solution is acidic
(a neutralization reaction) or one that is there just because of the auto-ionization of water:
OH- + H+ ⟶ H2O
A weak base like ammonia raises the hydroxyl ion concentration by accepting a proton from auto-ionized water:
H2O ⇌ H+ + OH-
NH3 + H+ ⇌ NH4+
So the Brønsted-Lowry definition of acids and bases is simple and encompasses most of the acids and bases we work with regularly. However, there are other compounds capable of changing pH that aren't so obvious, and another definition is required.
An acid is a proton donor and a base is a proton acceptor
We use the word "general" in science and math to mean "all encompassing." We always want to find the simplest rules and theories that fully describe all possibilities.
Not all substances need to directly donate a proton or hydroxyl group to a solution in order to reduce or increase its pH, so we need to refine our definition of acids a bases further. Gilbert Lewis did this in 1938.
Here is an example. Boron trifluoride (BF3) can act as a base in aqueous solution by first bonding an F- ion,
BF3 + F- ⟶ BF4-
Then we can think of BF4- working in aqueous solution like this:
BF4- + H2O ⟶ HBF4 + OH-
So BF4- acts as a base, in much the same way as the weak base ammonia (NH3. In forming BF4-, notice that BF3 has "accepted a pair of electrons from an F- ion, thus acquiring a full octet of electrons. The Lewis structure of BF3 looks like this:
When BF3 binds and F-, the Lewis structure looks like this:
The F- effectively donates a pair of electrons (outlined green dots) to BF3.
Lewis' definition of a acids and bases is that they accept or donate a pair of electrons, respectively. In the BF3 case, BF3 accepts a pair of electrons, and therefore acts as an acid, and F- acts as a base by donating a pair of electrons.
Lewis' definition is completely compatible with the neutralization reaction
H+ + OH- ⟶ H2O
if we look at where the electrons go:
Here the lone (acidic) proton accepts a pair of electrons from OH-, and therefore acts as an acid in the Lewis theory, and OH- is the electron pair donor – a Lewis base.
Lewis' theory is the most all-encompassing theory of what makes up an acid or base. It's not always easy to recognize components of a solution that will raise or lower pH if they aren't obvious H+ or OH- donating compounds, but the Lewis definition helps. The Lewis definition is also handy for acid and base solutions that aren't aqueous, and many such solutions exist.
In Lewis' theory of acids and bases,
An acid is an electron-pair acceptor.
A base is an electron-pair donor.
Like BF3, all of the elements under and including boron (group XIII) can act as Lewis bases. Likewise,
because ammonia is a Lewis base, so are all of the other group XV elements when they form compounds with three bonds, such as PH3 (phosphine).
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