1. 21 Titration: Reactions of Acids and Bases

2. The Self-ionization of Water  In pure water at 25 o C, both H 3 O + and OH- ions are found at concentrations of 1.0 x 10 -7 M  K w is called the ion-product constant K w = [H 3 O + ] [OH - ]  In pure water at 25 o C, K w = 1.0 x 10 -14

3. The Self-ionization of Water  K w is very useful because it applies not only to pure water, but to every water solution at 25 o C, even acidic or basic solutions.  Thus, if a solution has an H 3 O + concentration of 1.0 x 10 -2 M, then the OH - concentration must be 1.0 x 10 -12 M. i.e. [OH - ] = K w / [H 3 O + ]

4. pH  In 1909, Danish biochemist Soren Sorensen proposed a simple way to express the concentrations of H 3 O + ions based on logarithms. This scale is known as the pH scale.  The pH of a solution is -1 times the Log of the H 3 O + concentration in moles per Liter.  pH = -log[H 3 O + ]

5. pH = -log[H 3 O + ]  Because the pH scale is a logarithmic scale, each one-unit change in pH represents a 10-fold change in the concentration of H 3 O + ions.  The pH of a solution can be measured by using acid-base indicators such as litmus paper, or by using an electronic device called a pH meter.

6. pH[H 3 O + ] Concentration.1 M HCl010 0 M 110 -1 M Lemon juice210 -2 M 310 -3 M Banana410 -4 M Coffee510 -5 M Saliva610 -6 M Pure water710 -7 M Blood810 -8 M 910 -9 M Borax1010 -10 M Lime water1110 -11 M 1210 -12 M Bleach1310 -13 M 1.0 M NaOH1410 -14 M pH Scale

7. Buffers  A buffer is a mixture that is able to release or absorb H + ions, keeping a solution’s pH constant.  Most common buffers are mixtures of weak acids with their conjugate bases. i.e. a buffer of acetic acid and the acetate anion keeps the pH near 4.74  The amount of acid or base that a buffer can neutralize is called the buffer capacity. [See examples on page 634]

8. Acid-Base Titration  The concentration of a weak acid or a weak base can be easily calculated from the results of a procedure called an acid-base titration.  An acid-base titration is a carefully controlled neutralization reaction.

9. Acid-Base Titration  To conduct a titration, a standard solution is slowly added to the unknown solution until neutralization is complete – called the equivalence point.  The point at which the indicator changes color is called the end point of the titration.

10. Performing a Titration  To run a titration, the standard solution is slowly added to the unknown solution.  As the two solutions mix, the base in one solution neutralizes the acid in the other solution, a reaction that runs nearly to completion.

11. Performing a Titration  The reaction between acetic acid and sodium hydroxide is: HC 2 H 3 O 2 + NaOH H 2 O + NaC 2 H 3 O 2 The point at which the indicator changes color is called the end point of the titration.  Total # mol of H + = Total # mol of OH - This equation is key to calculating the concentration of an acid or base using data from a titration.

12. Performing a Titration  The reaction between oxalic acid and sodium hydroxide from Lab #57 is: H 2 C 2 O 4 + 2 NaOH 2H 2 O + Na 2 C 2 O 4 Note that it takes 2 mol of base to neutralize each 1 mol of acid. This equation is key to calculating the concentration of an acid or base using data from a titration.

13. Titration Strong Acid / Strong Base HCl + NaOH H 2 O + NaCl Weak Acid/ Strong Base HC 2 H 3 O 2 + NaOH H 2 O + NaC 2 H 3 O 2 Weak Base / Strong Acid NH 3 + HCl NH 4 Cl [Fig 19-14, 19-15, & 19-16 on pgs. 640-642]