1. Unit 11: Acids and Bases

2. Unit Overview…
We will learn about Acids and Bases, two important types of compounds in chemistry
Learn the distinct properties of Acids and Bases
Understand the pH scale, and how we can use indicators to detect pH in an Acid-Base solution
Examine important acid-base chemical reactions

3. Arrhenius’ Theory of Acids and Bases
Arrhenius Acid
Any substance that dissolves in water to produce hydrogen ions (H+ ) or Hydronium ions (H3O+)
Ex: Hydrochloric acid, HCl (a strong acid).
HCl (g) H+ (aq) + Cl- (aq)
H2O (l) + H+ (aq)  H3O+ (aq)
ACIDIC solutions are formed when an acid transfers a proton to water.

4. Hydronium Ion
When an H+ ion interacts with the electrons of the oxygen in a water molecule, an H3O+ ion called the hydronium ion is formed.

5. Arrhenius’ Theory of Acids and Bases
Arrhenius Base
Any substance that dissolves in water to produce hydroxide ions (OH- )
Ex: Sodium Hydroxide, NaOH (a strong base) or Ammonia NH3
NaOH (s) Na+ (aq) + OH- (aq)
NH3 (l) + H2O (l)  NH4+ (aq) + OH- (aq)
*Important Note – most bases end in “OH” but if there is CARBON attached to the compound then it is NOT a base

6. Properties of Acids
They’re electrolytes – chemicals that break up into ions in water, so they conduct electricity in water
Strong acid: good conductor
Weak acid: poor conductor
Sour Taste
React with most metals to produce H2 (g)
React with bases to form Water and salt
React with indicators to change color
pH values are from 0 to 6
see Reference Table M for color changes

7. Properties of Bases Also electrolytes Bitter Taste
Feel slippery (like soap)
React with acids to form water and salt
React with indicators to change color
pH values are 8 to 14
See Reference Table M for color changes

8. Properties of Salts Also electrolytes
Ionic Compounds – containing one metal ion (+) and one non-metal ion (-), or a polyatomic ion (except OH-)
Ex) NaCl, MgCl2, or Ca3(PO4)2
React with indicators to change color
pH values are around 7 (neutral)
See Reference Table M for color changes

9. pH Scale
A scale, called the pH scale, has been developed to express the concentration of hydrogen ions [H+] as a number from 0 to 14.
A pH of 0 is strongly acidic
A pH of 7 is neutral
A pH of 14 is strongly basic
(High concentration of H+)
Equal Concentration of H+ and OH- ions
(High concentration of OH-)

10. pH Scale

11. Acid – Base Indicators See Table M
An indicator is a chemical that changes color when it gains or loses a “proton” (H+ ions).
There are many indicators listed on Reference Table M, and the color changes associated with certain pH values are given.
See Table M
We can use multiple indicators
to find the approximate pH
of a substance

12. Table K and Table L
These two tables list the most common acids and bases you will see in this course.
Table K: The top 4 acids are strong acids
The last 2 are weak acids.
Table L: The top 3 bases are strong bases
The last 1 is a weak base.
Table L
Table K

13. Acidity vs. Alkalinity
Acidity – is a measure of the strength of an acid related to the concentration of H+ ions.
Alkalinity – is a measure of the strength of a base related to the concentration of OH- ions
High Acidity = pH less than 7
High Alkalinity = pH more than 7

14. Strong Acids
Strong acids: ionize (dissolve) completely, which means if 100 molecules are placed in water all 100 will break up into ions
Good Conductors!
Examples:
HCl  H+ + Cl-
H2SO4  2H+ + SO4-2

15. Weak Acids
Weak acids: ionize (dissolve) only slightly, If 100 molecules are placed in water then only a small percentage will break up into ions.
Bad Conductors…
Example: acetic acid “vinegar”
CH3COOH  CH3COO- + H+
*NOTE: if an organic compound ends in COOH it’s a weak acid!!!

16. Strong Bases Strong bases: ionize completely Good Conductors!
Example: Sodium Hydroxide
NaOH  Na+ +OH-

17. Weak Bases
Weak base: ionizes slightly (won’t dissolve completely in water)
Bad Conductors…
Example: ammonia
NH3 + H2O  NH4+ + OH-

18. Naming Acids

19. Binary acids –two elements - H+ *
- Previous Naming System
- New Naming System
hydrogen _____ ide
 hydro_____ic acid
Hydro_____ic acid
Hydrochloric acid
Hydrofluoric acid
Hydrosulfuric acid
Hydrobromic acid
Hydrogen _____ ide
Hydrogen chloride
Hydrogen fluoride
Hydrogen sulfide
Hydrogen bromide

20. Ternary acids: H+ & polyatomic ion
- Previous Naming System
- New Naming System
Hydrogen _____ate 
_________ic acid
Hydrogen _____ate
Hydrogen sulfate
Hydrogen chlorate
Hydrogen nitrate
Hydrogen phosphate
_________ic acid
Sulfuric acid
chloric acid
nitric acid
Phosphoric acid

21. More Ternary Acids – “-ite Ions”
- Previous Naming System
- New Naming System
Hydrogen _____ite 
_______ous acid
_______ous acid
Nitrous acid
Sulfurous acid
Chlorous acid
Hydrogen _____ite
Hydrogen nitrite
Hydrogen sulfite
Hydrogen chlorite

22. Naming Bases Ca(OH)2  Calcium Hydroxide
No change in how we name these…
Full name of the metal atom +
Hydroxide
Ex)
Ca(OH)2  Calcium Hydroxide

23. Acid and Base Reactions
You will need to be familiar with 2 types of reactions involving acids & bases.
1st reactions between any acid and certain metals.
2nd reactions combine acids & bases to produce salt and water (neutralization).

24. Acid-Metal Reactions
Most metals will corrode when exposed to certain acids.
When this happens H2 gas is produced.
Example: Ni + 2HCl  NiCl2 + H2 or
2Al + 6HCl  2AlCl3 + 3H2
Only precious metals won’t corrode like copper, gold, and silver because they aren’t reactive (Low on Table J)

25. Neutralization Reaction
In a neutralization reaction an Acid will react with a Base to produce salt and water.
Word equation
Formula equation
Net ionic equation
A net ionic equation has only the ions that take part in the reaction. The ions that don’t change are removed, called spectator ions.
Hydrochloric Acid + Sodium Hydroxide yields Water + Sodium Chloride
HCl NaOH  H2O NaCl
H+(aq) OH-(aq)  H2O(l)

26. Try This One H2SO4 + 2NaOH 2H2O + Na2SO4
Sulfuric acid + sodium hydroxide yields water + sodium sulfate
H2SO NaOH
2H2O Na2SO4
2H+ + SO Na+ + 2OH-
2H2O + 2Na+ + SO4-2
H+ + OH-
H2O

27. Finding Concentration of H+ in Acids
A 1.0 M HCl “monoprotic acid”
HCl  H+ + Cl-
Produces 1 H+ so the concentration of H+ is 1.0 M
Written like this: [H+] = 1.0 M
A 1.0 M H2SO4 “diprotic acid”
H2SO4  2H+ + SO4-2
Produces 2 H+ so the concentration of H+ is 2 x 1.0 M
Written like this: [H+] = 2.0 M

28. Finding Concentration of OH- in Bases
A 1.0 M KOH
KOH  K+ + OH-
Produces 1 OH- so the concentration of OH- is 1.0 M.
Written like this: [OH-] = 1.0 M
A 1.0 M Mg(OH)2
Mg(OH)2  Mg+2 + 2OH-
Produces 2 OH- so the concentration of OH- is 2(1.0 M).
[OH-] = 2.0 M

29. Titration
An Acid-Base titration is a technique used find the concentration of an acid or a base by neutralizing it.
During a titration you add set volumes of a base to an acid until it is neutralized.
Using the Acid-Base titration formula listed on Table T you can solve for your unknown concentration.

30. Acid- Base Titrations MaVa = MbVb (Table T) Ma = molarity of H+
Va = volume of acid
Mb = molarity of OH-
Vb = volume of base

31. Titrations (Neutralization) Problems
Ex. What volume of 0.50M HCl is required to neutralize 100mL of 2.0M NaOH?
MaVa = MbVb
Ma = 0.5 M
Va = ??? mL
Mb = 2.0 M
Vb = 100 mL
0.5(x) = 2.0(100)
X = 400mL HCl

32. Bronsted-Lowry Acid-Base Theory
Acids: chemicals that act as proton donors in water ( H+ + H2O  H3O+)
Bases: Chemicals that act as proton acceptors in water (H2O  H+ + OH-)
Water: Can act as an acid or a base because it can either donate or accept “protons” depending on the situation

33. Bronsted-Lowry Conjugate Acids/Bases
A conjugate acid - what is left of a Bronsted-Lowry Base after accepting a proton
A conjugate base - what is left of aa Bronsted-Lowry Acid after donating a proton
Ex)
HNO3 + H2O  NO3- + H3O+
Conj. Base
Conj. Acid
Base
Acid