Acid-base reactions are a key to understanding organic reactivity (see, for example, What makes a good leaving group?
Detailed information is given throughout Bruice, and collected in Appendix II. Qualitative acid-base behavior can be summarized thus:
- Brønsted acid strength of element-hydrogen bonds decreases moving horizontally on the periodic table (from fluorine to carbon or chlorine to phosphorus), but increases moving vertically, for example from fluorine to iodine. (Remember that the strength of a Brønsted base is proportional to the weakness of its conjugate acid!)
- Lewis base strength of lone pairs on atoms increases moving away from fluorine in any direction. See this summary.
- Anions are more basic than neutrals (for the same element only!); similarly, cations are more acidic (in a Lewis sense if there are no available protons).
- Anything with a lone pair can, in principle, be a base (unless it's already a cation like H3O+). More electronegative elements are LESS basic than less electronegative ones, all else being equal (that is, if they are both anions or both neutral). For example, NH3 is more basic than H2O; and NH2- is more basic than OH-, which is more basic than F-.
- On the other hand, going down within the same column drastically reduces basicity even though electronegativity also drops. Thus, PH3 is much less basic than NH3, and H2S is much less basic than H2O.
- Anything with a hydrogen bound to an electronegative atom can be a Bronsted acid ("electronegative atom" is, in practice, nitrogen, oxygen or sulfur, and of course the hydrohalic acids HF, HCl and HBr).
- In principle, anions such as OH- can be acids since they have H bonded to an electronegative atom. In practice, only a few anions can be easily deprotonated. They are mineral acid anions such as HSO4-, H2PO4-, HPO42- and HCO3-.
You need to know what is more acidic than what else. In rough order of decreasing Brønsted acidity:
|hydrohalic acids, mineral acids (except H2CO3, HCO3- and HPO42-) and protonated alcohols (ROH2+ and H3O+)
||pKa ~ -5 to +2
|carboxylic acids RCOOH and H2CO3
||pKa ~ 4-5
|phenols (aromatic alcohols, OH on a benzene ring), bicarbonate HCO3-, monohydrogen phosphate HPO42-, and ammonium cations R3NH+, R2NH2+, RNH3+, NH4+
||pKa ~ 8-10
|aliphatic alcohols ROH (including water)
||pKa ~ 15-18
|saturated hydrogens α to carbonyl groups, RCOCR'2-H|
aldehydes and ketones are more acidic than carboxylic acid derivatives
|pKa ~ 17-25|
|terminal alkynes ()
||pKa ~ 25
|amines R2NH, RNH2, NH3. The hydroxide anion OH- is also around here as a Brønsted acid.
||pKa ~ 35-40
|Vinylic hydrogens R2C=CR-H are slightly more acidic than alkyl hydrogens R3C-H, but in practice neither can be removed by most bases. Their conjugate bases will deprotonate OH- to give the oxide dianion O2-.
||pKa ~ 40-60
Once you know the order of acidity, you know the strength of the conjugate bases (conjugate bases of strong acids are weak; conjugate bases of weak acids are strong).
Remember that there are other factors, such as conjugation or electronegative substituents, that can make an acid much stronger. For example, trifluoroacetic acid is as strong as the mineral acids, with a pKa ~ 0.5. Acetaldehyde has a pKa of 17, in the alcohol range!
This class of organic acids--hydrogens α to carbonyl groups--is so important that it rates a whole chapter in Bruice, Chapter 18.
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