Acid and base chemistry is found everywhere
from manufacturing to biochemistry, but it
gets somewhat convoluted when all the different
possibilities are considered.
There are bronsted-lowry bases, bronsted-lowry
acids, lewis acids, and lewis bases--yet how
can their inter and intramolecular reactions
be told apart?
A bronsted-lowry acid donates a positive hydrogen
ion, a bronsted-lowry base accepts a positive
hydrogen ion, a lewis acid accepts a negative
electron pair, and a lewis base donates a
negative electron pair.
Thus, both types of bases become more positive
after their reaction, and both types of acids
become more negative.
It’s important to think of magnets: opposite
charges enjoy one another’s company.
How these acids and bases are told apart actually
depends on the reactants involved, as molecules
like water can act as bronsted-lowry acids
or bronsted-lowry bases, depending on the
other reactants’ acidity or basicity.
But something which is more basic will ‘enjoy’
becoming more positive, something which is
more acidic will ‘enjoy’ becoming more
negative, and the strengths of the acids or
bases could be determined by how close to
neutral their resulting products, or ‘conjugates’
are--discussed later in this video.
Strong acids are acids which give up all their
hydrogen ions in a solution, as in the bronsted-lowry
definition.
Such acids have hydrogen bound to really electronegative
atoms--electron-loving atoms--so the atoms
don’t mind giving up positivity in the form
of a positive hydrogen ion.
There is a list of these, and there is also
a list of strong bases, which, in contrast,
have less-electronegative group 1 and 2 atoms
that easily hand over electrons to their OH
group and afterwards release all their basic
OH groups in a solution, but many substances
do not fit into these categories.
For substances which do not meet the definition
of strong acids or strong bases, some clues
can be drawn from their structure to determine
which type of acid or base one may be.
For example, does an atom in the molecule
have a pair or pairs of electrons which are
not bound to another atom?
If so, the atom with unpaired electrons may
be coaxed to give those up, as in the lewis
base definition.
However, halogens are very electronegative,
and will not behave like this with their unpaired
electrons.
Azo dyes rely on the donation of neighboring,
unpaired electrons when struck by a ray of
light in sending back the light, producing
the spectacular colors many of our clothes
are dyed with, and beta-carotene works the
same way, giving a brilliant orange color
to carrots and leaves in fall.
Thus, groups with unpaired electrons, such
as with the ‘O’ in an OH group or with
the ‘N’ in an NH3 group, will also share
electrons easily, make an aromatic ring or
conjugated system more basic.
Such movement of electrons, or ‘resonance’,
can also cause substances to be more acidic
if the hydrogens are attached to an atom which
is attached to a ‘withdrawing group’ with
pi bonds, as with a carbonyl group, where
a carboxylic acid may easily give up its positive
hydrogen ion and become stabilized by the
molecule’s ability to share electrons, allowing
the hydrogens to leave more easily.
Since the movement of electrons is able to
‘stabilize’ the structure, the effect
is either more basic or more acidic, depending
on the arrangement of the atoms.
Similarly, if the ‘conjugate’, or ‘product’
of acids or bases is more neutral, this helps
drive the reaction forward, comparable to
an exothermic reaction, which is more favorable
and spontaneous, and the strongest acids therefore
have the weakest conjugate bases, while the
strongest bases have the weakest conjugate
acids.
Once the acidity or basicity of the reactants
involved is determined, it is easier to figure
out if the reactants will act as lewis acids,
lewis bases, bronsted-lowry acids, or bronsted-lowry
bases.
