1. Acids, Bases, and pH
Chapters 14/15

2. Ba(s) + H2SO4(aq) → BaSO4(s) + H2(g)
Acids
1. Aqueous solutions of acids have a sour taste.
2. Acids change the color of acid-base indicators.
3. Some acids react with active metals and release hydrogen gas, H2.
Ba(s) + H2SO4(aq) → BaSO4(s) + H2(g)
4. Acids react with bases to produce salts and water.
5. Acids conduct electric current.

3. Properties of Acids

4. Some Common Industrial Acids
Sulfuric Acid
Sulfuric acid is the most commonly produced industrial chemical in the world.
Nitric Acid
Phosphoric Acid
Hydrochloric Acid
Concentrated solutions of hydrochloric acid are commonly referred to as muriatic acid.
Acetic Acid
Pure acetic acid is a clear, colorless, and pungent-smelling liquid known as glacial acetic acid.

5. Bases 1. Aqueous solutions of bases taste bitter.
2. Bases change the color of acid-base indicators.
3. Dilute aqueous solutions of bases feel slippery.
4. Bases react with acids to produce salts and water.
5. Bases conduct electric current.

6. Properties of Bases

7. Arrhenius Acids and Bases
An Arrhenius acid is a chemical compound that increases the concentration of hydrogen ions, H+, in aqueous solution.
An Arrhenius base is a substance that increases the concentration of hydroxide ions, OH−, in aqueous solution.

8. Strength of Acids and Bases
A strong acid is one that ionizes completely in aqueous solution.
a strong acid is a strong electrolyte
HClO4, HCl, HNO3, HClO, HBr, HI, H2SO4
A weak acid releases few hydrogen ions in aqueous solution.
hydronium ions, anions, and dissolved acid molecules in aqueous solution
HCN
Organic acids (—COOH), such as acetic acid

9. Strength and Weakness of Acids and Bases

10. Brønsted-Lowry Acids and Bases
A Brønsted-Lowry acid is a molecule or ion that is a proton donor.
Hydrogen chloride acts as a Brønsted-Lowry acid when it reacts with ammonia. HCl + NH3 → NH Cl-
A Brønsted-Lowry base is a molecule or ion that is a proton acceptor.
Ammonia accepts a proton from the hydrochloric acid. It acts as a Brønsted-Lowry base. HCl + NH3 → NH Cl-

11. Brønsted-Lowry Acids and Bases

12. pH = −log [H3O+] = −log(1 × 10−7) = −(−7.0) = 7.0
The pH Scale
The pH of a solution is defined as the negative of the common logarithm of the hydronium ion concentration, [H3O+].
pH = −log [H3O+]
example: a neutral solution has a [H3O+] = 1×10−7
The logarithm of 1×10−7 is −7.0.
pH = −log [H3O+] = −log(1 × 10−7) = −(−7.0) = 7.0

13. The pH Scale
The pOH of a solution is defined as the negative of the common logarithm of the hydroxide ion concentration, [OH−].
pOH = −log [OH–]
example: a neutral solution has a [OH–] = 1×10−7
The pH = 7.0.
pH + pOH = 14.0

14. The pH Scale

15. Comparing pH and pOH