1. Titration
In this technique, an acid (or base) solution of known concentration is slowly added to a base (or acid) solution of unknown concentration.
A pH meter or indicators are used to determine when the solution has reached the equivalence point: The amount of acid equals that exactly neutralizes
the base.

2. Titration of a Strong Acid with a Strong Base
From the start of the titration to near the equivalence point, the pH goes up slowly.
Just before (and after) the equivalence point, the pH rises rapidly.
At the equivalence point, pH = 7.
As more base is added, the pH again levels off.

3. Titration of a Strong Base with a Strong Acid
It looks like you “flipped over” the strong acid being titrated by a strong base.
Start with a high pH (basic solution); the pH = 7 at the equivalence point; low pH to end.

4. Titration of a Weak Acid with a Strong Base
Use Ka to find initial pH.
Find the pH in the “buffer region” using stoichiometry followed by the Henderson–Hasselbalch equation.
At the equivalence point the pH is >7. Use the conjugate base of the weak acid to determine the pH.
As more base is added, the pH levels off. This is exactly the same as for strong acids.

5. Ways That a Weak Acid Titration Differs from a Strong Acid Titration
A solution of weak acid has a higher initial pH than a strong acid.
The pH change near the equivalence point is smaller for a weak acid. (This is at least partly due to the buffer region.)
The pH at the equivalence point is greater than 7 for a weak acid.

6. Use of Indicators
Indicators are weak acids that have a different color than their conjugate base form.
Each indicator has its own pH range over which it changes color.
An indicator can be used to find the equivalence point in a titration as long as it changes color in the small volume change region where the pH rapidly changes.

7. Indicator Choice Can Be Critical!

8. Indicators summary

9. 1. strong acid and strong base

10. 2. Weak acid and strong base

11. 3. Strong acid and weak base

12. 4. Weak acid and weak base

13. 8.5 Acid Deposition
Rain is naturally acidic because of dissolved CO2 and has a pH of Acid deposition has a pH below 5.6.
Acid deposition is formed when nitrogen or sulfur oxides dissolve in water to form HNO3, HNO2, H2SO4, and H2SO3.

14. Causes of acid deposition
All rain water is naturally acidic due to the presence of dissolved carbon dioxide (carbonic acid).
Acid rain refers to rain water with a pH below 5.6 due to the additional presence of acidic pollutants.
Acid deposition = all of the precipitation
Wet acid deposition- rain, snow, sleet, hail, fog, mist, dew fall to ground as aqueous precipitates
Dry acid deposition = acidifying particles or gases fall to the ground as dust or smoke and later dissolve in water to form acids.

15. Sulfur oxides
Produced from burning fossil fuels, particularly coal and heavy oil in power plants. Also smelting metals.

16. Nitrogen oxides
Produced mainly from internal combustion engines. Burning the fuel provides the heat energy necessary for nitrogen and oxygen in air to combine.

18. Impact on materials
(write some equations, list the effects, and the impact)

19. What happens to plants?

20. What happens to water?

21. What happens to human health?

22. List some things that are being done to help reduce these problems.