Brønsted-Lowry Acids and Bases

    A Bronsted-Lowry acid is defined as anything that releases H1+ ions; a Bronsted-Lowry base is defined as anything that accepts H1+ ions. This definition includes all Arrhenius acids and bases but, as we will soon see, it is a bit more general. The Bronsted-Lowry concept is based on the transfer of a proton from one substance to another. The ionization of HCl in water can be viewed as a Bronsted-Lowry acid-base reaction with HCl behaving as the acid (H1+ ion donor) and water serving as the base (H1+ ion acceptor). Just as we saw with Arrhenius acids, the acidic hydrogens are usually written at the beginning of the chemical formula of a Bronsted-Lowry acid. So, for instance, we can tell from the formulas that HCl, HNO3, and HCH3O2 are monoprotic acids, each liberating one H1+ ion per acid molecule, whereas H2SO4 is diprotic, and H3PO4 is triprotic.

    The reaction of HCl with water can be called either an Arrhenius acid-base reaction or a Bronsted-Lowry acid-base reaction. There are many acid-base reactions, however, for which the Arrhenius definition is inappropriate. Let's consider the gas phase reaction of HCl with ammonia, NH3, which occurs as shown in the chemical equation below. The arrows represent the movement of electron pairs as bonds are formed and broken.


    In this reaction the ammonia behaves as the proton acceptor, a Bronsted-Lowry base, while the HCl serves as the proton donor, a Bronsted-Lowry acid. Notice that a pair of nonbonding electrons on the base is used to form a covalent bond with the hydrogen of the acid. The covalent bond between the hydrogen and the chlorine atom is broken, and this pair of electrons becomes nonbonding on the chlorine atom. To function as a Bronsted-Lowry base a substance must have a pair of nonbonding electrons that can be used to form a bond to H1+. Be careful though, because the presence of nonbonding electrons is not always apparent. When we write chemical formulas, the nonbonding electrons are not shown explicitly so it may not be obvious that systems such as PH3, F1-, or amines are bases.