1. Acids, Bases, and Salts

2. Acid/Base Theory Theory Acid Definition Base Definition Arrhenius
Substance that releases a H+ ion in water
Substance that releases an OH- ion in water
Bronsted-Lowry
Substance that donates a proton (H+)
Substance that accepts a proton (H+)
- H+ ion called a proton because no electron, just one proton left (show on board)

3. Classifying acids, bases, salts
donates a proton (H+)
Bases
accepts a proton (H+)
Salts
dissolve in water; bound ionically
Is not classified as an acid or base using the definitions above

4. Acids, Bases and Salts
Examples:
HCl
NaOH
NaCl
H2SO4
Ca(OH)2
H2O

5. Multiple Hydrogen Atoms
Monoprotic acids have one Hydrogen
Ex: HCl, HBr, HI
Diprotic acids have two hydrogen
Ex: H2SO3, H2SO4, H2CO3
Triprotic acids have three hydrogen
Ex: H3PO4, H3PO3

6. HCl(aq) + H2O(l) → H3O+(aq) + Cl-(aq)
Bases without OH-
Example:
HCl(aq) + H2O(l) → H3O+(aq) + Cl-(aq)
Is HCl an acid or a base? Why?
Is H2O an acid or a base? Why?
- make sure they know what the hydronium ion is

7. H2O (l) + NH3(l) → NH4+(aq) + OH-(aq)
Bases without OH-
Example:
H2O (l) + NH3(l) → NH4+(aq) + OH-(aq)
Is H2O an acid or a base? Why?
Is NH3 an acid or a base? Why?

8. Autoionization of Water
H2O(l) + H2O(l) → H3O+(aq) + OH-(aq)
Is H2O(l) the acid or base?
Kw=?

9. Naming Acids Hydrogen bonded with an element: Examples HBr H2S
hydro__________ic acid
Examples
HBr H2S
hydrobromic acid hydrosulfuric acid

10. Naming Acids Hydrogen bonded with a polyatomic ion: Examples
______________ic acid
Examples
HNO H2SO4
nitric acid sulfuric acid

11. Naming Acids Acids with a different number of oxygens: Formula Name
HNO3
nitric acid
HNO2
nitrous acid
HNO
hyponitrous acid
HNO4
pernitric acid

12. Naming Acids Example - naming acids: HClO3 HClO2 HClO HClO4
- homework: acids, bases, salts ws

13. Electrolytes Acids, Bases, and Salts dissociate in water
Hydration Simulation
Acids, Bases, and Salts
dissociate in water
(break apart into ions)
Ions conduct electricity
Substances can be:
strong electrolytes
weak electrolytes
nonelectrolytes

14. Strong vs Weak Strong Acids HCl HNO3 HI HClO3 HClO4 H2SO4 Strong Bases
Group 1A metal with hydroxides
Heavy group 2A metals(Ca, Sr, Ba) with hydroxide.

15. Completely breaks into ions Partially breaks into ions
Electrolytes
Type of Electrolyte
Solubility
Ions
Conductor of
electricity?
Strong electrolyte
Dissolves in water
Completely breaks into ions
Very good
conductor
Weak electrolyte
Partially breaks into ions
Poor
Nonelectrolyte
No ions No
conduction
since ions have charges, they have the ability to “transport” electrons, thus they can conduct electricity
demo: light bulb connected with wires to graphite electrodes

16. Other Properties of Acids and Bases
Metals in Acids
Recall the zinc in HCl demo
What was produced?
Taste
Acids taste __________
Bases taste _________

17. Determining Acidity Acidity measured using the pH scale
- pH scale starts at 0 and ends at 14. closer to 0, acidic. closer to 14, basic. 7 is neutral solution.

18. pH = -log [H+] or pH = -log [H3O+]
Determining Acidity pH
measure of the H+ or H3O+ ions in solution
pH = -log [H+] or pH = -log [H3O+]
pOH
measure of the OH- ions in solution
pOH = -log [OH-]
pH + pOH = 14
pH: 0 to 14, closer to zero means more acidic
pOH: 0 to 14, closer to zero means more basic
Scales are exact opposites of one another

19. Calculating pH and pOH Example
Find the pH and the pOH for the following:
A 3.0 x 10-4 M solution of HCl?
show on board (balanced equation, M of H+, pH equation)
lower pH means stronger acids, means strong electrolyte (same for high pH and strong bases)

20. Calculating pH and pOH Example
Find the pH and the pOH for the following:
A 1.9 x 10-9 M solution of NaOH?
show on board (balanced equation, M of H+, pH equation)
lower pH means stronger acids, means strong electrolyte (same for high pH and strong bases)

21. Calculating pH and pOH Example
Find the pH and the pOH for the following:
a) 5.4 x 10-2 M solution of KOH?
b) 1.0 M solution of HBr?
show on board (balanced equation, M of H+, pH equation)
lower pH means stronger acids, means strong electrolyte (same for high pH and strong bases)
pOH = 1.3
pH = 12.7
pOH = 14
pH = 0

22. Neutralization Reactions
General equation:
Acid + Base → Salt + Water
Example:
HCl + NaOH →
HCl + OH- → Cl- + H2O
NaCl + H2O
another type of double replacement reaction
remember, salts are electrolytes and dissolve in water: form ions
Acid
Base
Conjugate Acid
Conjugate Base

23. Neutralization Reactions
Example:
List which is the acid, the base, the conjugate acid, and the conjugate base.
HNO3 + OH- →
H2O + NO3-

24. Titrations Titration Neutralization reaction occurs
The reaction between an acid and a base used to find the concentration of the unknown acid or base.
Neutralization reaction occurs
An indicator is used to find the “end point”
Equivalence point (i.e. the “end point”)
when the number of moles of the acid equals the number of moles of the base
titration demo: prep for lab tomorrow
show how calculations can be done to find unknown M of base
titration lab

25. Titrations
Use the equation below when calculating the unknown concentration of the acid or the base
MAVA = MBVB
MA = molarity (concentration) of acid
VA = volume of acid used during titration
MB = molarity (concentration) of base
VB = volume of base used during titration

26. Titrations
Example:
It takes mL of NaOH to reach the end point with mL of M HCl. What is the concentration (molarity) of NaOH?

27. Titrations
Example:
It takes mL of NaOH to reach the end point with mL of 1.2 M HCl. What is the concentration (molarity) of NaOH?

28. Indicators Indicator pH range for color change 0.5-1.25 3.25-4.25 5-6
Methyl violet
Bromphenol blue
Methyl red
5-6
Litmus
6-7.5
Phenolphthalein
8-10
Alizarin yellow
indicator is chosen depending on what pH it changes color and on how dramatic that color change is (the more dramatic, the easier it is to see)
phenolphthalein is the most commonly used in chemistry

29. Natural Indicators
Red Cabbage

30. Natural Indicators Hydrangea Acidic Soil
Alkaline (Basic) Soil or Neutral Soil