- A Brønsted-Lowry acid is a molecule that donates H+ when it reacts.
- A Brønsted-Lowry base is a molecule that accepts H+ when it reacts.
- A conjugate acid-base pair differ only by having (acid) or not having (base) an H+
- A Lewis acid is a molecule that accepts a pair of electrons when it reacts.
- A Lewis base is a molecule that donates a pair of electrons when it reacts.
- An Arrhenius acid is a molecule that forms H3O+ when it reacts with water.
- An Arrhenius base is a molecule that forms OH− when it reacts with water.
- Strong acids react essentially to completion with water.
- Strong bases dissociate essentially to completion in water.
- Reactants are favoured at equilibrium when weak acids or bases react with water.
- An amphoteric species is one that can act as either an acid or a base.
at 298 K
A Brønsted-Lowry acid is a molecule that donates H+ when it reacts, while a Brønsted-Lowry base is a molecule that accepts H+ when it reacts. The reaction below shows how a generic acid, HA, and base, B, react:
Notice that when the acid (HA) reacts, the A loses its bond to hydrogen but keeps the two electrons from the bond to form A− on the product side. B, conversely, uses its lone pair of electrons on the reactant side to make a new bond with the hydrogen that was just released by HA. This forms HB+ on the product side. At a minimum, every Brønsted-Lowry acid must contain an H (so that it can donate H+) and every Brønsted-Lowry base must have a lone pair of electrons (which it uses to make a new bond to H when accepting H+). Notice that just like the reactants, our two products also fulfill these requirements: A− has a lone pair of electrons and could thus act as a Brønsted-Lowry base while HB+ has a hydrogen and could thus act as a Brønsted-Lowry acid. When a Brønsted-Lowry acid (e.g. HA) reacts, it always forms a base (e.g. A−) as a product. This pair of molecules (e.g. HA and A−) that differ only by an H+ are called a conjugate acid-base pair. In the example above, HB+ and B are also a conjugate acid-base pair. Try the activity below to practice identifying conjugate acid-base pairs.
The Brønsted-Lowry definition is not the only way to classify an acid or a base. For example, it is also possible to classify some molecules as Lewis acids/bases or Arrhenius acids/bases. Lewis acids accept a pair of electrons when they react while Lewis bases donate a pair of electrons when they react. Let's consider the same generic reaction that we used to introduce Brønsted-Lowry acids and bases:
Notice that here A in HA accepts the electrons from the bond to H to form A− when it reacts, so HA is classified as a Lewis acid. B, however, donates its lone pair of electrons to make a new bond with H+, so B is classified as a Lewis base. The Lewis acid/base definition is broader than the Brønsted-Lowry definition, such that all Brønsted-Lowry acids/bases are Lewis acids/bases, but not all Lewis acids/bases are Brønsted-Lowry acids/bases (see Venn diagram below). Conversely, the Arrhenius acid/base definition is a narrower definition than the Brønsted-Lowry definition since it classifies acids/bases based only on how they react with water. An Arrhenius acid reacts with water to form H3O+ while an Arrhenius base reacts with water to form HO-.
For each reaction below, classify the specified molecule as an Arrhenius acid/base, Brønsted-Lowry acid/base, and/or Lewis acid/base.
The Arrhenius definition of an acid or base focuses on which products form when the acid/base reacts with water. We can further classify acids and bases based on their strength by also considering how much they react with water; that is, we consider the equilibrium constant, K, of the reaction between the acid/base and water. For an acid reacting with water, the equilibrium constant is called the acid dissociation constant, Ka.
Recall that pure liquids, such as water here, are not included when constructing equilibrium expressions. When the acid reacts only a small amount with water such that the reactants are favoured at equilibrium (Ka < 1) we classify the acid as weak. If the acid reacts essentially completely with water such that the products are strongly favoured at equilibrium (Ka >> 1) we classify the acid as strong. Likewise, the equilibrium constant for a base reacting with water is called the base dissociation constant, Kb, and a base is classified as weak when Kb < 1 or strong when Kb >> 1.
A handful of strong acids and strong bases are so prevalent that it is helpful to memorize that they are strong. The strong acids that you should be familiar with are:
- HCl
- HBr
- HI
- HNO3
- H2SO4
- HClO4
The strong bases that you should be familiar with are:
- MOH where M is an alkali metal from Group 1 of the periodic table (e.g. Na+, Li+, K+)
- M(OH)2 where M is an alkaline Earth metal from Group 2 of the periodic table (e.g. Mg2+, Ca2+)
Water, H2O, is amphoteric, which means that it can react as an acid or a base. In fact, water will react with itself where one molecule will serve as the acid and the other as the base, as in the reaction shown below:
This reaction is called the self-ionization of water since it involves water reacting to make two ions. The equilibrium constant for this reaction is labelled as Kw, and has a value of Kw = 1.0 x 10−14 at 298 K. Kw is very helpful in converting between the Ka of an acid and the Kb of its conjugate base. To demonstrate this, consider the generic conjugate acid-base pair HA and A− and how they each react with water.
Notice that if we multiply together the Ka for the weak acid, HA, and Kb for its conjugate base, A−, then we get Kw.
This relationship that is very helpful in converting between Ka and Kb for a conjugate acid-base pair. Since
is always equal to Kw, this also gives us the insight that the larger a Ka value is (i.e. the stronger an acid is) the smaller the Kb value of the conjugate base will be (i.e. the weaker the conjugate base will be).
The values of Ka and Kb can be very large or very small, and are thus often reported using scientific notation. To discuss the relative strengths of acids and bases without resorting to scientific notation, Ka and Kb values are also often converted into a logarithmic scale using and
.
Do weaker acids and weaker bases have larger or smaller pKa and pKb values?
Weak acids and weak bases have small postive values: Ka<1 and Kb<1, respectively. The logarithm of a number that is less than 1 is negative, so for weak acids logKa is negative and for weak bases logKb is negative. As an acid (base) gets weaker, the Ka (Kb) value gets smaller and thus the logKa (logKb) value also gets smaller (i.e. more negative). Because pKa and pKb are defined as the negative logarithms of Ka and Kb, respectively, this means that as acids and bases get weaker the pKa and pKb values get larger. Conversely, smaller pKa and pKb values correspond to stronger acids and bases, repsectively.
Just like Ka and Kb values, the concentrations of H3O+ and HO− in solution can vary greatly. Again, defining a logarithmic scale for these values as and
is helpful to compress the range of these values. At 298 K,
and thus by taking the negative logarithm of both sides here
we can see that at this temperature
The values and relationships introduced on this page are summarized below as they will be very useful in solving problems as you proceed through the rest of the sections on acids and bases.
Interactive: