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4.1: Bronsted-Lowry acids and bases

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    We’ll begin our discussion of acid-base chemistry with a couple of essential definitions. The first of these was proposed in 1923 by the Danish chemist Johannes Brønsted and the English chemist Thomas Lowry, and has come to be known as the Brønsted-Lowry definition of acidity and basicity.  An acid, by the Brønsted-Lowry definition, is a species which acts as a proton donor (i.e., it gives away an H+), while a base is a proton (H+) acceptor.  One of the most familiar examples of a Brønsted-Lowry acid-base reaction is between hydrochloric acid and hydroxide ion:

    image001.png

    In this reaction, a proton is transferred from HCl (the acid, or proton donor) to hydroxide ion (the base, or proton acceptor). As we learned in the previous chapter, curved arrows depict the movement of electrons in this bond-breaking and bond-forming process.

    After a Brønsted-Lowry acid donates a proton, what remains is called the conjugate base.  Chloride ion is thus the conjugate base of hydrochloric acid. Conversely, when a Brønsted-Lowry base accepts a proton it is converted into its conjugate acid form: water is thus the conjugate acid of hydroxide ion.

    Here is an organic acid-base reaction between acetic acid and methylamine:

    image002.png

    In the reverse of this reaction, acetate ion is the base and methylammonium ion (protonated methylamine) is the acid.

    image003.png

    For now, let’s just consider one common property of bases: in order to act as a base, a molecule must have a reactive pair of electrons.  In all of the acid-base reactions we’ll see in this chapter, the basic species has an atom with a lone pair of electrons. When methylamine acts as a base, for example, the lone pair of electrons on the nitrogen atom is used to form a new bond to a proton.  A negative charge often (but not always!) indicates that a structure (in this case, an anion) is likely to act as a base.

    Clearly, methylammonium ion cannot act as a base – it does not have a reactive pair of electrons with which to accept a proton.

    image004.png

    In summary,

    • A Brønsted-Lowry acid is a proton (hydrogen ion) donor.
    • A Brønsted-Lowry base is a proton (hydrogen ion) acceptor.

    When a Brønsted acid HA dissociates in water, it increases the concentration of hydrogen ions in the solution, H+; conversely, Brønsted bases dissociate by taking a proton from the solvent (water) to generate OH-.

    • Acid dissociation

    \[HA_{(aq)} \rightleftharpoons A^-_{(aq)} + H^+_{(aq)}\]

    • Acid Ionization Constant:

    \[K_a=\dfrac{[A^-][H^+]}{[HA]}\]

    • Base dissociation:

    \[B_{(aq)} + H_2O_{(l)} \rightleftharpoons HB^+_{(aq)} + OH^-_{(aq)}\]

    • Base Ionization Constant

    \[K_b = \dfrac{[HB^+][OH^-]}{[B]}\]

    The determination of a substance as a Brønsted-Lowry acid or base can only be done by observing the reaction. In the case of the H2O it is a base in the first case and an acid in the second case.

    Water does not need to be involved in a Bronsted-Lowry reaction.  In general, for an acid HA and a base Z, we have

    \[ HA + Z \rightleftharpoons A^- + HZ^+ \]

    • A Donates H to form HZ+.
    • Z Accepts H from A which forms HZ+
    • A becomes conjugate base of HA and in the reverse reaction it accepts a H from HZ to recreate HA in order to remain in equilibrium
    • HZ+ becomes a conjugate acid of Z and in the reverse reaction it donates a H to A recreating Z in order to remain in equilibrium

    4.1: Bronsted-Lowry acids and bases is shared under a not declared license and was authored, remixed, and/or curated by LibreTexts.

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