Now that we’ve defined acids and bases, let’s discuss how they worktogether in reactions. Look at the generic acid-base reaction below:
HX(aq) + H2O(l) 
X-(aq) + H3O(aq)
When the forward reaction occurs, HX donatesa proton to water (so it acts as the base) to form hydronium. When thereverse reaction occurs, the hydronium ion acts as the acid, donating aproton to the X-. Together, HX and X- are said tobe conjugate acid-base pairs. Conjugate acid-base pairs are compoundsthat differ by the presence of one proton, or H+. All acidshave a conjugate base, which is formed when their proton has beendonated; likewise, all bases have a conjugate acid, formed after theyhave accepted a proton.
Example
Apply the appropriate acid-base theory tofirst identify the acid and base reacting and then identify theconjugate acid-base pairs in the examples below:
Explanation
- In this first reaction, we see that HNO3 gives a proton to water, which then forms a hydronium ion. This makes HNO3 the acid in the forward reaction, and water acts as the base. HNO3’s conjugate base is
, and water’s conjugate acid is the hydronium ion, or 
- Here
donates the proton to water, so in the forward reaction it acts as the acid, and water is still the base. 
’s conjugate base is NH3, and water’s conjugate acid is again the hydronium ion, H3O+.
Relative Strengths of Acids and Bases
Certain acids are stronger than other acids, andsome bases are stronger than others. What this means is that some acidsare better at donating a proton, and some bases are better protonacceptors. A
strong acid or base dissociates or ionizes completely in aqueous solution. A
weak acid or base does not completely ionize.
Strong Acids
There are six strong acids that you’ll need to memorize for the
SAT II Chemistry test:
- Hydrohalic acids: HCl, HBr, HI
- Nitric acid: HNO3
- Sulfuric acid: H2SO4
- Perchloric acid: HClO4
Let’s take a closer look at how acids differin strength by focusing on perchloric acid. In general, the
greater thenumber of oxygen atoms in a polyatomic ion, the stronger the acid.
zn1ewK u;Gd:z So HClO4 is stronger than HClO3, which is stronger than HClO2,which is stronger than HClO. (Perchloric acid is the strongest amongthe six, but all the other oxyacids of chlorine are not consideredstrong acids.) Now think about why, as you take away oxygens, thestrength of the acid decreases. The hydrogen (proton) to be removed isbonded to an oxygen atom. The oxygens are highly electronegative andare pulling the bonded pair of electrons
away from the site where the hydrogen is bonded, thus making it easier to remove the H+. As the number of oxygen decreases, the molecule becomes less polar, and the H+ is harder to remove.
Strong Bases
There are fewer strong bases to memorize forthe exam. These are hydroxides (—OH), oxides of 1A and 2A metals(except Mg and Be), H-, and
.Remember that the stronger the acid, the weaker its conjugate base, andthe converse is also true. The figure below illustrates the relativestrengths of some common conjugate acid-base pairs.
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The pH Scale
As you know, water can act as either a proton donor (in the form of the hydronium ion, H3O+) or a proton acceptor (as OH-).In solution, a water molecule can even donate a proton to or accept aproton from another water molecule, and this process is called
autoionization:
2H2O
H3O+ + OH-
Since this reaction takes place in equilibrium, we can write an equilibrium expression,
Keq, for it:
Keq = [H3O+][OH-]
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And since this expression refers specifically to the ionization of water, we can write the equilibrium expression as
Kw. At 25ºC, the value of
Kw, which is known as the
ion-product constant, is 1
10-14. This means that the [H3O+] = [OH-] and each is equal to 1
10-7. When the concentrations of H+ and OH- are equal in a solution, the solution is said to be neutral. In acidic solutions, the concentration of H+ is higher than that of OH-, and in basic solutions, the concentration of OH- is greater than that of H+.
The
pH of a solution is calculated asthe negative logarithm in base 10 of the hydronium ion concentration—itis an expression of the molar concentration of H+ ions in solution:
pH = -log [H+] or -log [H3O+]
A solution like the equilibrium expression for water, which is neutral at standard temperature, would have a pH of
pH = -log [1
10-7] = -(-7.00) = 7.00
So as you can see, neutral solutions have apH of 7. If the solution contains more hydronium ions than this neutralsolution ([H+] > 1
10-7),the pH will be less than 7.00, and the solution will be acidic; if thesolution contains more hydroxide ions than this neutral solution ([OH-] > 1
10-7), the pH will be greater than 7.00, and the solution will be basic.
Similarly, the pOH of a solution is calculated as the negative logarithm in base 10 of the hydroxide ion concentration:
pOH = -log [OH-]
and pH and pOH are related to each other by the equation
pH + pOH = 14
Since you won’t be allowed to have acalculator for the SAT II Chemistry test, you can use the followingequation if you need to calculate the hydronium ion concentration of asolution:
[H3O+] = 10-pH
Now try a problem: What is the pH of a solution at 25ºC in which [OH-] = 1.0
10-5
M?
Explanation
The fact that this solution is at 25ºC tells us that we should use the
Kw relationships. If the [OH-] = 1.0
10-5 M, then pOH = 5. You know that 1.0
10-5 is the same as plain old 10-5. The log of 10-5 is -5 (simply use the exponent when a number, any number, is written as 10power, so the “negative” of the log is equal to -(-5), or simply 5. Now, if the pOH is 5, then the pH is 9 since pH + pOH = 14.
