Five Key Factors That Influence Acidity
Before we get started , however , let’s quickly review the basics of
acidity and basicity. Here’s the condensed version:
1.
Bronsted acids are proton
donors, Lewis acids are electron pair acceptors. Converse: Bronsted
base = proton acceptor, Lewis base = electron pair donor.
2.
A conjugate base is what you obtain when you remove a proton
(H+) from a compound.
For instance, HO(-) is the conjugate base
of water. O(2-) is the conjugate base of HO(-).
Conversely, conjugate acids are what you obtain when you add a proton to a
compound. The conjugate acid of water is H3O(+).
3.
Quick quiz: is pH 1
acidic or basic? pKa
is similar to pH in that low (and even negative values) denote strong acids. That’s because pKa
is based on the equilibrium:
4.
According to this, anything
which stabilizes the conjugate base will increase the acidity.Therefore pKa is also a
measure of how stable the conjugate base is. Put another way, strong acids have
weak conjugate bases, and vice versa.
With that out of the
way, let’s get started.
Factor #1 – Charge.
Removal of a proton, H+ , decreases the formal charge on an atom or molecule by
one unit.
This is, of course, easiest to do when an atom bears a charge of +1
in the first place, and becomes progressively more difficult as the overall charge
becomes negative.
The acidity trends reflect this:
Note that once a
conjugate base (B-) is negative, a second deprotonation
will make the dianion (B 2-).
While far from impossible, forming the dianion can be
difficult due to the buildup of negative charge and the corresponding
electronic repulsions that result.
Factor #2 – The Role of the Atom
This point causes a
lot of confusion due to the presence of two seemingly conflicting trends.
Here’s the first
point: acidity increases as we go across a row in the periodic table. This makes sense,
right? It makes sense that HF is more electronegative than H2O, NH3, and CH4
due to the greater electronegativity of fluorine
versus oxygen, nitrogen, and carbon.
A fluorine
bearing a negative charge is a happy fluorine.
But here’s the
seemingly strange thing. HF itself is not a “strong” acid, at least not in the
sense that it ionizes completely in water. HF is a weaker acid than HCl, HBr, and HI. What’s going on
here?
You could make two
arguments for why this is.
The first reason has
to do with the shorter (and stronger) H-F bond as compared to the larger
hydrogen halides. The second has to do with thestability of the conjugate
base.
The fluoride anion, F(–) is a tiny and vicious little beast, with the smallest
ionic radius of any other ion bearing a single negative charge.
Its charge is therefore spread over a smaller
volume than those of the larger halides, which is energetically unfavorable:
for one thing, F(–) begs for solvation,
which will lead to a lower entropy term in the ΔG.
Note that this trend
also holds for H2O and H2S, with H2S being about 10 million times more acidic.
Factor #3 – Resonance.
A huge stabilizing
factor for a conjugate base is if the negative charge can be delocalized
through resonance. The classic examples are with phenol (C6H5OH) which is about
a million times more acidic than water, and with acetic acid (pKa of ~5).
Watch out though – it
isn’t enough for a π system to simply be adjacent to a proton – the
electrons of the conjugate base have to be in an orbital which allows for
effective overlap
Factor #4 – Inductive effects. Electronegative atoms can draw negative charge toward
themselves, which can lead to considerable stabilization of conjugate bases.
Check out these examples:
Predictably, this
effect is going to be related to two major factors: 1) the electronegativity
of the element (the more electronegative, the more acidic) and the distance
between the electronegative element and the negative charge.
Factor #5 – Orbitals. Again, the acidity relates nicely to the stability of the
conjugate base. And the stability of the conjugate base depends on how well it can accomodate its newfound pair
of electrons. In an effect akin to electronegativity,
the more s character in the orbital, the closer the electrons will be to the
nucleus, and the lower in energy (= stable! ) they will be.
Look at the difference
between the pKa of acetylene and alkanes
– 25! That’s 10 to the power of 25, as in, “100 times bigger than Avogadro’s
number”.
Just to give you an idea of scale.
That’s the amazing thing
about chemistry – the sheer range in the power of different phenomena is awe-inspiring.
There’s actually a
mnemonic I’ve found that can help you remember these effects. I can’t take
credit for it, but here it is:
Charge
Atom
Resonance
Dipole Induction
Orbitals
= CARDIO.
PSK CHAKRAVARTHY
No comments:
Post a Comment