Let's take the $\ce{KOH}$ molecule into account. I know it is a base from literature, but how would one go about determining if a molecule is acidic or basic simply based on the structure of the molecule? How about amphoteric?

Also, I understand that there are two complementary systems - Lewis and Brønsted-Lowry theories. How do they work and how do they fit together?

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    $\begingroup$ I think this is a very good question. Many Bronsted acid and bases are easy to recognize because they have an obvious place to donate or accept a proton (e.g. $\ce{HCl}$ and $\ce{KOH}$ are easy) and some amphiprotic compounds are also easy (e.g. $\ce{H_2O}$ and $\ce{HCO_3^{-}}$). My 'feeling' is, but please let someone give you a more thorough answer, is that Bronsted-lowry acids and bases are easily recognized by proton donating/accepting sites, and that lewis acids/bases are more difficult to 'see coming' $\endgroup$
    – Michiel
    Commented Nov 3, 2013 at 8:28
  • $\begingroup$ What is unclear, with respect to Bronsted and Acid base and related articles? $\endgroup$
    – ssavec
    Commented Nov 3, 2013 at 8:55
  • $\begingroup$ Wouldn't it be as simple as saying that entities with low energy LUMOs tend to be acidic, while those with high energy HOMOs tend to be basic? $\endgroup$ Commented Nov 3, 2013 at 9:45
  • $\begingroup$ @NicolauSakerNeto it is not as simple as that since electron transfer reactions can be explained with HOMO/LUMO gaps as well. The relative energies of HOMOs and LUMOs of products and reactants leads to a qualitative (semi-quantitative?) explanation for acid/base behavior. $\endgroup$ Commented Nov 3, 2013 at 15:08
  • $\begingroup$ @bobthechemist Well yes, low/high LUMO/HOMO species can act in ways other than acids/bases, but I believe it's true that there is no such thing as a base without a high HOMO and an acid without a low LUMO (this is all in relative terms, of course)? Therefore, all bases would have (relatively) high HOMOs, but not all entities with high HOMOs act as bases (and the equivalent for acids). $\endgroup$ Commented Nov 3, 2013 at 15:25

3 Answers 3


It is sometimes challenging to determine if a molecule is going to be acidic or basic if the system in which it is reacting is not considered. An important point to consider when dealing with acids and bases is that acid/base strength is inherently tied to the solvent. For this answer, I'm going to limit the discussion to acids and bases in an aqueous environment.

It is helpful to consider the terms acid and base as a means to classify substances. This way, Chemists can explain chemical reactivity and structure-function relationships of substances. Very early classification systems depended on our senses (acids are sour, bases are slippery to the touch) and more recent classification systems utilize structural characterization tools such as NMR or crystallography. Many classification systems have been proposed over the years, and only a few of them have found sufficiently widespread use to end up in textbooks used in the standard Chemistry curriculum. Below are a few systems, taken from Miessler & Tarr's Inorganic Chemistry textbook. The 2nd through 4th entries are those in common use today.

  • NAME (YEAR) acid definition [example]; base definition [example]
  • Liebig (~1776) Acid: an oxide of N, P, S [$\ce{SO_3}$]; Base: Reacts with acid [$\ce{NaOH}$]
  • Arrhenius (1894) Acid: Forms hydronium ion [$\ce{HNO_3}$]; Base: Forms hydroxide ion [$\ce{NaOH}$]
  • Brønsted (1923) Acid: Proton donor [$\ce{HCl}$]; Base: Proton acceptor [$\ce{NaOH}$]
  • Lewis (1923) Acid: Electron-pair acceptor [$\ce{Ag^+}$]; Base: Electron-pair donor [$\ce{NH_3}$]
  • Ingold-Robinson (1932) Acid: Electrophile [$\ce{BF_3}$]; Base: Nucleophile [$\ce{NH_3}$]
  • Lux-Flood (1939) Acid: Oxide ion acceptor [$\ce{SiO_2}$]; Base: Oxide ion donor [$\ce{CaO}$]
  • Usanovich (1939) Acid: Electron acceptor [$\ce{Cl_2}$]; Base: Electron donor [$\ce{Na}$]
  • Solvent system (1950s) Acid: Solvent cation [$\ce{BrF_2^+}$]; Base: Solvent anion [$\ce{BrF_4^-}$]
  • Frontier Orbitals (1960s) Acid: LUMO of acceptor [$\ce{BrF_3}$]; Base: HOMO of donor [$\ce{NH_3}$]

I find the various acid/base systems very enlightening. Note how From Arrhenius through Lewis there was a broadening of the acid/base classification system; Arrhenius can't be used to describe non-aqueous acids and bases and Brønsted can't be used with aprotic substances. Yet, the Lewis definition incorporates the previous two (a Brønsted acid is also a Lewis acid; an Arrhenius base is also a Lewis base). One then may ask, what's up with the Lux-Flood definition then? This definition is counter to the trend of broadening the classification system, and yet it is useful in describing anhydrous solid-state chemistry and is used to describe geochemical reactions as well as the chemistry of high-temperature melts. My point being: classification of substances into acids and bases is only meaningful if it helps explain chemical phenomena.

Which brings us to the Usanovich definition, which essentially states that every reaction is an acid-base reaction. Such a broad definition is not overly helpful, and to those of us with an affinity towards electrochemistry (ahem), is somewhat arrogant :-)

I do believe that determining if a substance will behave as an acid or a base requires a bit of chemical intuition (or a Socratic method). For example, you may have performed an experiment in which $\ce{KOH}$ served as a base; you know that, like potassium, sodium is an alkali metal; therefore you presume that $\ce{NaOH}$ would be a base as well.

Back to the question at hand

So how do I suggest one use this information to predict whether a substance will behave as an acid or a base using its structure alone? Personally, I find the Lewis theory as the most useful classification system in answering this type of question. If the structure of a compound is set before me and I am to predict its acid/base chemistry, I will ask two questions:

  • Are there any lone pair electrons that can be donated?
  • Are there any electron-deficient atoms that could serve as electron pair acceptors?

If the answer to question 1 is yes, then the molecule can behave as a base. If the answer to question 2 is yes, then the molecule is an acid. If both are true, then I have an amphoteric substance.

In proofreading this answer, I realize I said I would restrict myself to aqueous systems, in which case using the Brønsted system may be more helpful. In this case, the questions become:

  • Is there a hydrogen that can be donated? (You'll be right more often than you are wrong if you rephrase this question as "Is there a hydrogen that is attached to something other than carbon?").
  • Is there a lone pair that can accept a proton?
  • $\begingroup$ Your date for Liebig's definition has to be wrong. Justus von Liebif was born in 1803. $\endgroup$
    – Ben Norris
    Commented Apr 3, 2014 at 1:07
  • $\begingroup$ @BenNorris Thanks for pointing this out. The dates are from Miessler and Tarr's Inorganic Chem text, and two entries are provided under the Liebig definition - the second entry has a date of 1838. The "1776" definition of an acid is the Oxide of N, P, S and the 1838 is "H replaceable by a metal". I will have to investigate why both are attributed to Liebig. $\endgroup$ Commented Apr 3, 2014 at 3:06
  • $\begingroup$ Shouldn't one consider first how the compound would react towards dissolution in water? For example sodium hydroxide is a salt and will produce sodium and hydroxide ions. $$\ce{NaOH <=> Na+ +\ ^{-}\!OH}$$ Hence only the hydroxide ion should be considered a base. (+1 for the extensive collection of definitions) $\endgroup$ Commented Apr 3, 2014 at 6:27
  • $\begingroup$ @Martin If we limit our conversation to aqueous solutions (which is not unreasonable) then yes, we must consider the compounds interaction with water, especially since it is a leveling solvent. In water, hydroxide is the strongest base, but not the only base. $\endgroup$ Commented Apr 3, 2014 at 14:04
  • $\begingroup$ I think it is not necessary to limit it to aqueous solutions. (It is mandatory up to and including Brønsted.) My statement holds for any solutions and probably even for the gas phase. If you are not dissolving sodium hydroxide, i.e. in toluene, then it will be only lying there, doing nothing. If you are considering Lewis definition (and after) for NaOH, then you already have an acid-base-pair. $\endgroup$ Commented Apr 4, 2014 at 3:20

How does one tell if a specific molecule is acidic or basic?

An acid (from the Latin acidus/acēre meaning sour) is a chemical substance whose aqueous solutions are characterized by a sour taste, the ability to turn blue litmus red, and the ability to react with bases and certain metals (like calcium) to form salts.

In chemistry, a base is a substance that, in aqueous solution, is slippery to the touch, tastes bitter, changes the colour of indicators (e.g., turns red litmus paper blue), reacts with acids to form salts, and promotes certain chemical reactions (base catalysis).

We do have $3$ different definitions from Arrhenius, Brønsted-Lowry, and Lewis. We would be in trouble, in some cases, where one definition accepts a molecule as acid or base and other not. So, we do have a common definition for acids and bases as mentioned above.

Coming to amphoteric substance, consider $\ce{H2O}$ which has no taste.



In organic chemistry, whether a compound is acidic or basic can be predicted by looking at the molecular structure and can quite easily be explained by applying the resonance theory. In essence, if the compound can ionize to form a proton and the resulting anion is stable, then the ionization equilibria will be shifted to the right and the compound is acidic. The more stable the anion, the more acidic the compound is.
Consider the following examples:

  1. Sulfuric acid $\ce{H2SO4 <=> 2H+ + SO4^2-}$ the sulfate anion is stabilized by 4 resonant structures, therefore sulfuric acid is very acidic.
  2. Nitric acid $\ce{HNO3 <=> H+ + NO3-}$ the nitrate anion is stabilized by 3 resonant structures, therefore nitric acid is acidic.
  3. Carboxylic acids $\ce{RCOOH <=> H+ + RCOO-}$ the carboxylate anion is stabilized by 2 resonant structures, therefore carboxylic acids are acidic.
  4. Phenols versus alcohols. $\ce{C6H5OH <=> H+ + C6H5O-}$ the phenoxide anion has 4 resonant structures while the alkoxide anion has none. Therefore, phenols are more likely to ionize to produce a proton than alcohols. That is, phenols are more acidic than alcohols.
  5. Ketones versus alcohols. $\ce{R1COCHR2R3 <=> H+ + R1COC-R2R3}$ the resulting anion has 2 resonant structures whereas alcohols have none. Therefore the α-hydrogen to the carbonyl group is acidic because the enoxide ion is more stable than the alkoxide ion (keto-enol tautomerism).
  6. Amides $\ce{RCONH2 <=> H+ + RCONH-}$ the resulting anion has 2 resonant structures, but they are not equivalent because one has the negative charge on the oxygen, the other one has it on the nitrogen. Therefore the anion is not particularly stable so amides are only weakly acidic. The lone pair of electrons on the nitrogen is only slightly available for bonding because they are withdrawn by the more electronegative oxygen. Therefore amides are only weakly basic. They are practically neutral.
  7. Sulfonamides $\ce{RS(O)2NH2 <=> H+ + RS(O)2NH-}$ the resulting anion has 3 resonant structures, 2 of which have the negative charge on the more electronegative oxygen atom. Therefore sulfonamides are a little more acidic than amides and can form salt with metals (e.g. silver sulfadiazine, sodium sulfacetamide).
  8. Urea $\ce{H2NCONH2}$ the lone pair of electrons on one nitrogen are withdrawn by the more electronegative oxygen (see amide above), but the lone pair on the other nitrogen is available for bonding. Therefore urea is basic.
  9. Imides $\ce{R1CONHCOR2 <=> H+ + R1CON-COR2}$ the resulting anion has 3 resonant structures, 2 of which have the negative charge on the more electronegative oxygen, therefore imides are acidic (for example phthalimide).
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    $\begingroup$ -1 for suggesting that carbonyl compounds (pKa ~20) are more acidic than alcohols (pKa ~15); it's simply not true. Also it's misleading to suggest that the two amine groups in urea are any different. Lastly, resonance is great, but hardly the only deciding factor. HI is very acidic, yet it is difficult to argue that the iodide ion is resonance-stabilised in any manner. $\endgroup$ Commented Mar 13, 2019 at 20:33
  • $\begingroup$ I was not saying all ketones are more acidic than alcohols. I am just saying that the alpha hydrogen to the carbonyl gp is acidic because of keto-enol tautomerism. I was not saying that the 2 amino groups in urea are different, I was just saying that because there are still electrons available even oxygen has suppressed them somewhat. $\endgroup$
    – Isaac Lai
    Commented Mar 13, 2019 at 21:57
  • $\begingroup$ There's a large gap in the ability to handle Lewis acids with this description. $\endgroup$
    – Zhe
    Commented Mar 13, 2019 at 22:41
  • $\begingroup$ The resonance theory serves well in explaining the acid-base nature of pharmaceuticals. Lewis acids and bases are too toxic to be used in medicine, Therefore it is rarely used to explain acid-base properties of drugs. $\endgroup$
    – Isaac Lai
    Commented Mar 14, 2019 at 1:23
  • $\begingroup$ I haven't come across any FDA approved drugs that cannot be explained by the Bronsted-LowryTheory. If you know of any, please let me know. $\endgroup$
    – Isaac Lai
    Commented Mar 14, 2019 at 1:38

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