Acid and Base
o
Arrhenius acid - Any substance that releases H+ ions when placed in water.
e.g.
o
Arrhenius base - Any substance that releases OH- ions when placed in water.
e.g.
Arrhenius's theory helped to explain why acids have similar characteristics, since they all give H+ ions when they
dissolve in water. It also explained why acids are neutralized by bases and why bases are neutralized by acids;
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The H+ ions from acids combine with the OH ions from bases to form water:
Though the Arrhenius theory helped to explain more about acids and bases, there were still several drawbacks to this
theory.
1.
2.
3.
The theory can only classify substances when they are dissolved in water since the definitions are based
upon the dissociation of compounds in water.
It does not explain why some compounds containing hydrogen such as HCl dissolve in water to give acidic
solutions and why others such as CH4 do not.
The theory can only classify substances as bases if they contain the OH - ion and cannot explain why some
compounds that don't contain the OH- such as Na2CO3 have base-like characteristics.
To extend the Arrhenius theory a little further, consider the formation of water from the combination of an H + ion
and an OH- ion:
This reaction is actually reversible, represented by the forward/backward arrow in the following reaction:
Based on the fact that the above reaction is reversible, we can conclude the following operational definitions for
acids and bases.
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Acid - Any substance that increases the concentration of the H+ ion when it dissolves in water.
Base - Any substance that increase the concentration of the OH- ion when it dissolves in water.
NOTE: A common shorthand notation for the concentration of a substance is placing it in brackets:
[H+] = concentration of the H+ ion
Now, substances not containing H+ or OH- ions can be classified as acids or bases if they alter the [H +] or [OH-]
when they dissolve in water.
e.g. CO2 cannot dissociate to give H+ but it does increases [H+] when it is dissolved in water.
e.g. CaO cannot dissociate to give OH- but it does increase [OH-] when it dissolves in water.
In 1923, Johannes Brønsted and Thomas Lowry separately proposed a new set of definitions for acids and bases
which are known as either Brønsted acids and bases or Brønsted-Lowry acids and bases.
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Brønsted Acid - Any substance that can donate a proton, H+ ion to a base.
aka: hydrogen-ion donors or proton donors

Brønsted Base - Any substance that can accept a proton, H+ ion from an acid.
aka: hydrogen-ion acceptor or proton acceptor
In the above reaction, the H from HCl is donated to H2O which accepts the H to form H3O+, leaving a Cl- ion.
The dissociation of water can be represented as follows:
The Brønsted-Lowry model of acids and bases brings rise to the concept of conjugate acid-base pairs. The part of
the acid remaining when an acid donates a H+ ion is called the conjugate base. The acid formed when a base accepts
a H+ ion is called the conjugate acid. For the generic acid HA:
For the generic base A-:
More examples of conjugate acid-base pairs:
In the following reactions, it is shown how H2PO4- and H2O can act as both acids and bases. Such compounds are
said to be amphoteric.
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Strong acids have weak conjugate bases.
Strong bases have weak conjugate acids.
Water has the tendency to equalize the strengths of all strong acids and strong bases, regardless of the strength of the
acid itself. This is known as the leveling effect. Acids are limited to the strength of the H 3O+ ions that they form
when they lose H+ ions when they dissolve in water. Likewise, bases are limited to the strength of the OH- ions that
they form when they gain H+ ions when they dissolve in water.
The relative strength of an acid is described as an acid-dissociation equilibrium constant.
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Acid-dissociation equilibrium constant (Ka) - For the generic acid reaction with water:
The acid-dissociation equilibrium constant is the mathematical product of the equilibrium concentrations of the
products of this reaction divided by the equilibrium concentration of the original acid:
Strong acids dissociate almost completely in water and therefore have relatively large Ka values. Weak acids,
however, dissociate only slightly in water and therefore have relatively small K a values. To distinguish between
strong and weak acids, the following guidelines are used:
Advantages to the Brønsted-Lowry model of acids and bases
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Acids and bases can now be ions or neutral molecules.
Acids and bases can now be any molecule with at least one pair of nonbonding electrons.
It explains the role of water in acid-base reactions; H2O accepts H+ ions from acids to form H3O+ ions.
It can be applied to solutions with solvents other than water and even in reactions that occur in the gas or
solid phases.
It relates acids and bases to each other with conjugate acid-base pairs and can explain their relative
strengths.
It can explain the relative strengths of pairs of acids or pairs of bases.
It can explain the leveling effect of water.