Acidity and Basicity

Chapter 1: Structure Determines Properties

Acids and Bases

Your organic teachers are quite likely to ask you questions like identify the most acidic protons or the most basic site in a molecule. These facts can be important for determining where a molecule is likely to react when treated with a base or acid respectively. Many students can not do this efficiently. The following topics are covered here:

Definitions : Arrenhius, Bronsted, Lewis
Acidity : Ka and pKa
Common acids in organic chemistry
Basicity
Common bases in organic chemistry

Remember that acidity and basicity are the based on the same chemical reaction (but looking at it from opposite sides) and both happen simultaneously. In the following simple example the base, B, removes a proton from the acid, H-A:

Can you think of another type of reaction that involve opposites simultaneously ?

Definitions
There are three theories used to describe acids and bases :

	Acids	Bases
Arrenhius	Ionise to give H⁺in H₂O	Ionise to give HO^-in H₂O
Bronsted-Lowry	A proton donor	A proton acceptor
Lewis	An electron pair acceptor	An electron pair donor

Now, some terminology:

Look at this equation and see how it fits the Bronsted-Lowry and Lewis definitions.

Acidity
Here are some general guidelines of principles to look for that can help you address the issue of acidity:
First, consider the simplified general equation of a simple acid reaction:

The more stable the conjugate base, A^-, is then the more the equilibrium favours the product side (Ka > 1)
The more the equilibrium favours products, the more H⁺ there is....
The more H⁺ there is then the stronger H-A is as an acid....
So looking for factors that stabilise the conjugate base, A^-, gives us a "tool" for assessing acidity.
The larger Ka implies more dissociation of HA and so the stronger the acid.
The larger Ka is, the more negative the pKa so the lower the pKa, the stronger the acid.

Key factors that affect the stability of the conjugate base, A^-,


HF > H₂O > NH₃ > CH₄	Electronegativity. When comparing atoms within the same row of the periodic table, the more electronegative the anionic atom in the conjugate base, the better it is at accepting the negative charge.

HI > HBr > HCl > HF	Size. When comparing atoms within the same group of the periodic table, the larger the atom the weaker the H-X bond and the easier it is for the conjugate base to accommodate negative charge (lower charge density)

RCO₂H > ROH	Resonance. In the carboxylate ion, RCO₂^-the negative charge is delocalised across 2 electronegative oxygen atoms which makes it more stable than being localised on a specific atom as in the alkoxide, RO^-.

General acidity trend of common organic acids (a useful sequence to remember):

acidity of common organic functional groups

Basicity
A convenient way to look at basicity is based on electron pair availability.... the more available the electrons, the more readily they can be donated to form a new bond to the proton and, and therefore the stronger base.

Key factors that affect electron pair availability in a base, B


CH₃^-> NH₂^-> HO^- > F^-	Electronegativity. When comparing atoms within the same row of the periodic table, the more electronegative the atom donating the electrons is, the less willing it is to share those electrons with a proton, so the weaker the base.

F^-> Cl^- > Br^- > I^-	Size. When comparing atoms within the same group of the periodic table, the larger the atom the weaker the H-X bond and the lower the electron density making it a weaker base.

RO^- >RCO₂^-	Resonance. In the carboxylate ion, RCO₂^-the negative charge is delocalised across 2 electronegative atoms which makes it the electrons less available than when they localised on a specific atom as in the alkoxide, RO^-.

Common organic bases:

Note that the factors for acidity and basicity are just the reverse of each other.