1. …all you need to “get” for the test… In 20 minutes!

2. AcidBase  Produces hydronium in aqueous (water) solutions (Arrhenius)  Donates hydrogen ions to another species (Bronsted-Lowry)  Taste sour  pH < 7  Turns litmus (and many other indicators red)  Produces hydroxide in aqueous (water) solutions (Arrhenius)  Receives hydrogen ions from acid (Bronsted- Lowry)  Taste bitter; feel slippery  pH > 7  Turns litmus (and many other indicators blue)

3. A compound’s ability to behave as an acid is that’s compound’s ability to “donate” hydrogen ions (protons).  “Strong” acids release those ions VERY readily and completely  For example CH 4 is NOT an acid—at all!  That donation is represented thusly:  H 2 SO 4 + H 2 O HSO 4 1- + H 3 O 1+ (1 st ionization)  HSO 4 1- + H 2 O SO 4 2- + H 3 O 1+ (2 nd ionization)

4.  HSO 4 1- + H 2 O SO 4 2- + H 3 O 1+  What you should notice:  HSO 4 1- becomes SO 4 2- ; therefore, (donates H 1+ )  in the reverse, SO 4 2- becomes HSO 4 1- (receives H + )  H 2 O becomes H 3 O 1+ ; therefore, (receives H + )  In the reverse, H 3 O 1+ becomes H 2 O (donates H + )  Translation: for weak ionizations and/or dilute solutions, that are reversible (in equilibrium), acids become conjugate bases, and, conversely, bases become conjugate acids.

5.  HF + H 2 OH 3 O + + F -  NH 4 + + OH - NH 3 + H 2 O  CO 3 2- + H 2 OHCO 3 - + OH -

6.  Hydronium ions in the presence of hydroxide ions can form water!  Of course, the leftovers ions form a “salt”.  For example:  HCl (aq) + NaOH (aq) H 2 O (l) + Na + (aq) Cl - (aq)  Because both the acid and the base are “strong”, the resulting hydronium and hydroxide concentration are equal.  The resulting pH is neutral. The “salt” is sodium chloride.  Another example:  HSO 4 - + NaOHH 2 O (l) + Na + + SO 4 2- + OH -  The resulting solution is still basic.

7.  The actual measurements of concentration result in the calculation of pH.  Pure water is defined by equal concentrations of hydrogen ions and hydroxide ions.  [H 3 O + ] = [OH - ] = 1 x 10 -7 M  [H 3 O + ] x [OH - ] = 1 x 10 -14 (memorize these numbers)

8.  Using the logarithmic function of those concentrations, we get the pH scale:  Water has a pH of 7  pH = -log [H 3 O + ]  Higher concentrations of hydronium means a smaller log!  2.34 x 10 -4 [H 3 O + ] = 3.63  Smaller concentrations mean higher logs!  2.34 x 10 -10 [H 3 O + ] = 9.63

9.  Because a species is only an acid or a base in water, the concentrations of these ions are related:  [H 3 O + ] [OH - ] = 1 x 10 -14  Which means that as one concentration increases, the other decreases…. (don’t forget the constant.)  One can also take the pOH of the hydroxide concentration.  Interestingly, pH + pOH = 14

12.  Buffer- a solution that resists changes in pH when limited amounts of acid OR base are added.  Ions of “weak” acids and bases, by definition, mean ions that are available to receive or to donate hydrogen ions &/or hydroxide ions.  CO 2(g) + H 2 O (l) H 2 CO 3 (aq) H + (aq) + HCO 3 - (aq)

13.  Compile 3 questions to ask/clarify/review:  1.  2.  3.