1. Acids
Bases

2. Acids Acids were first recognized as a
distinct class of compounds because
of the common properties of their
aqueous solutions.
Aqueous solutions have a sour taste (many are corrosive
and poisons).
Acids change the color of acid-base indicators.
Some acids react with active metals to release hydrogen
gas, H2.

3. Common Industrial Acids
Acids react with bases to
produce salts and water.
Some acids conduct an electric current.
Common Industrial Acids
Sulfuric Acid
Used in petroleum refining, automobile
batteries, and used as a water-removing
agent (can cause serious burns).

4. Common Industrial Acids
Nitric Acid
Used in making explosives (stains
skin and causes serious burns).
Phosphoric Acid
Used for manufacturing fertilizers and
animal feed and for flavoring beverages.

5. Common Industrial Acids
Hydrochloric Acid
Used in food processing.
Acetic Acid
Used in the manufacture of
plastics and as a fungicide.

6. Bases Aqueous solutions of bases have a bitter taste.
Bases change the color of acid-base indicators.
Dilute aqueous solutions of bases feel slippery.
Bases react with acids to produce salts and water.
Bases conduct electric current.

8. Acid-Base Theories Arrhenius Acid
An Arrhenius acid is a chemical
compound that increases the concentration
of hydrogen ions, H+, in aqueous solution.
Examples: HCl, H2SO4, HI, HF

9. Arrhenius Base An Arrhenius base is a substance that increases the
concentration of hydroxide, OH-, in aqueous solution.
Examples: NaOH, LiOH, Ca(OH)2

10. Bronsted-Lowry Acids and Bases
Acid-Base Theories
Bronsted-Lowry Acids and Bases
A Brønsted-Lowry acid is a molecule or ion that is a
proton donor (H+).
A Brønsted-Lowry base is a molecule or ion, that is a
proton acceptor.
In a Brønsted-Lowry acid-base reaction, protons are
transferred from one reactant (the acid) to another (the base).
HCO3 + HOH → H2CO3 + OH-
Base
Acid

11. Acid-Base Theories Lewis Acids and Bases
A Lewis acid is an atom, ion, or molecule that accepts an
electron pair to form a covalent bond.
A Lewis base is an atom, ion, or molecule that donates
an electron pair to form a covalent bond.
A Lewis acid-base reaction is the formation of one or
more covalent bonds between an electron-pair donor and
an electron-pair acceptor.

13. Concepts of pH

14. Hydronium and Hydroxide Ions
Pure water is a weak electrolyte and undergoes
self-ionization as seen below.
Remember ionization is the process of adding or removing electrons from an atom or molecule, which gives the atom or molecule a net charge.
In the self-ionization of water, two water molecules
produce a hydronium ion and a hydroxide ion by
transfer of a proton.
hydronium ion
hydroxide ion
(Acid)
(Base)

15. Concentrations of hydronium ion (H3O+) and
hydroxide ion (OH-) are 1.0 x 10-7 mol/L in
water at 25 oC.
The product of [H3O+][OH-] = 1.0 x and
the product is called the ionization constant
of water, Kw.
Note: [ ] means molar concentration (mol/L).

16. Neutral, Acidic, and Basic
Solutions
Any solution in which [H3O+] = [OH-] is neutral.
Any solution in which [H3O+] is greater than [OH-] is
acidic. Remember acids increase the concentration of H3O+ in aqueous solutions.
Any solution in which [OH-] is greater than [H3O+] is
basic. Remember bases increase the concentration of OH- in aqueous solutions.

17. pH Instead of expressing acidity or basicity in
terms of the concentration of H3O+ or OH-, a
quantity called pH is used.

18. pH Scale The pH scale commonly has values from 0 to 14 with a
pH value of 7 considered neutral.
pH values less than 7 are acidic and pH value greater
than 7 are basic.

19. Calculating pH pH indicates the hydronium ion (H3O+) concentration
of a solution.
The pH of a solution is defined as the negative of the
common logarithm of the hydronium ion concentration.
pH = -log[H3O+] or pH = -log[H+]
pOH (hydroxide ion concentration)
pOH = -log[OH-]
pH + pOH = 14.0

20. NO MORE CHEMISTRY NOTES!!
HALLELUJAH!!!
NO MORE CHEMISTRY NOTES!!