# 6.2: 6.2 Acids, Bases, and Salts

#### Beyond Water: Structure of Acids and Bases

To define acids and bases, we’ll use the definitions below:

Acid – A substance that can, to at least some extent, donate a proton into solution.

Base – A compound that can, to at least some extent, accept a proton from solution.

Acids and bases are generally categorized into either strong acids or weak acids. Although this is a continuum rather than a hard definition, it is a useful one for understanding aqueous chemistry. Also, the categories are more often applied to acids than bases. So, some examples of strong acids and weak acids are shown below:

 Strong Acids Weak Acids Nitric Acid, HNO3 Acetic Acid, HC2H3O2 Sulfuric Acid, H2SO4 Ammonium ion, NH4+ Hydrochloric Acid, HCl ...many others

For these three strong acids, we will assume that there is 100% complete dissociation into proton and conjugate base.

For all weak acids, we cannot assume that there is 100% dissociation. Instead, there will be some partitioning between the acid and base forms.

We use the equilibrium constant to quantify the tendency of an acid to dissociate, and we call that equilibrium constant Ka.  A weak acid (the form with the proton) and its conjugate base (the form without the proton) can have very different properties. So, it is important that we can quantify the amount or concentration of the compound in both the protonated and unprotonated forms.

Some examples that we’ll show in class include:

Hypochlorous Acid, HOCl:

$\ce{ HOCl (aq) \rightleftharpoons H^{+} (aq) + OCl^{-} (aq)} \nonumber$

$K_a=\dfrac{[\ce{H^{+}}][\ce{OCl^{-}}]}{[\ce{HOCl}]} \nonumber$

Bicarbonate Ion, HCO3-

$\ce{ HCO3^{-} (aq) \rightleftharpoons H^{+} (aq) + CO3^{-2} (aq)} \nonumber$

$K_a=\dfrac{[\ce{H^{+}}][\ce{CO3^{-2}}]}{[\ce{HCO3^{-}}]} \nonumber$

Acetic Acid, CH3COOH

$\ce{ CH3COOH (aq) \rightleftharpoons H^{+} (aq) + CH3COO^{-} (aq)} \nonumber$

$K_a=\dfrac{[\ce{H^{+}}][\ce{CH3COO^{-}}]}{[\ce{CH3COOH}]} \nonumber$

Ammonium Ion, NH4+

$\ce{ NH4^{+} (aq) \rightleftharpoons H^{+} (aq) + NH3(aq)} \nonumber$

$K_a=\dfrac{[\ce{H^{+}}][\ce{NH3}]}{[\ce{NH4^{+}}]} \nonumber$

In these Ka equations I have left off the phase designations for now, since we will be working primarily in the aqueous phase.

Because the numeric values of most acid dissociation equilibrium constantsare <0.1, we usually tabulate them as pKa values, or, -logKa. For example, if Ka = 1×10-6, pKa = 6.

#### Salts of Acids and Bases

Species that participate in acid/base reactions don’t always start out as acids or bases. For instance, sodium can be combined with phosphate in dry chemical form to make sodium phosphate, Na3PO4(s). If we dissolve it in water at some concentration below its solubility, it will generally all dissolve, right?

$\ce{Na3PO4}(s)⟶\ce{3Na^{+}}(aq)+\ce{PO4^{-3}}(aq)$

But, phosphoric acid is a weak acid. In solution, it behaves according to the reactions below:

$\ce{ H3PO4 (aq) \rightleftharpoons H^{+}(aq) + H2PO4^{-} (aq) \rightleftharpoons 2H^{+}(aq) + HPO4^{-2}(aq) \rightleftharpoons 3H^{+}(aq) + PO4^{-3}(aq)} \nonumber$

In our system of interest here we have water, some of which has dissociated into H+ and OH-.

$\ce{ H2O (l) \rightleftharpoons H^{+} (aq) + OH^{-} (aq)} \nonumber$

So, some of that PO43- from the sodium phosphate salt will react with H+ back up the phosphoric acid dissociation reaction pathway:

$\ce{ H^{+}(aq) + PO4^{-3}(aq) \rightleftharpoons HPO4^{-2} (aq) + H^{+}(aq) \rightleftharpoons H2PO4^{-}(aq) + H^{+}(aq) \rightleftharpoons H3PO4 (aq)} \nonumber$

In essence, we have created a weak acid system by adding the sodium salt of the acid.

Common clues that you are dealing with a salt:

The cation is sodium or potassium, and its associated anion is a component of a weak acid/base system.

The anion is chloride, and its associated cation is a component of a weak acid/base system

For the problems we will be doing, we will usually be working with salts in which the cations are sodium or potassium. These are monovalent cations, so their salt compounds are very soluble in water and will result in that salt completely dissolving. Sometimes calcium and magnesium salts will come into play, as well, and we’ll also assume they completely dissolve. If the salt is an anion, it will usually be chloride that we are looking at. Chloride salts of acids and bases are highly soluble in water and completely dissociate.

Learn to recognize the salts of acids and basis, so that you can determine when a compound will completely dissociate into its constituent ions. This will make it much easier to solve acid and base equilibria.