1. Acids & Bases

2. • Some are electrolytes
Properties of Acids
Sour taste
Change color of acid-base indicators (red in pH paper)
Some react with active metals to produce hydrogen gas
Ba(s) + H2SO4(aq) BaSO4(s) + H2(g)
Some react with bases to neutralize and form salt and water
H2SO4 (aq) + 2NaOH(aq) Na2SO4 (aq) + 2H2O(l)
• Some are electrolytes

3. Examples of Acids Lemons and oranges - citric acid
Vinegar - 5% by mass acetic acid
Pop and fertilizer - phosphoric acid

4. • Some are electrolytes
Properties of Bases
Bitter taste
Change color of acid-base indicators (blue in pH paper)
Dilute aqueous solutions feel slippery
Ex. Soap
• Some react with acids to neutralize and form salt and water
• Some are electrolytes

5. Examples of Bases Soap - NaOH Household cleaners - NH3
Antacids - Ca(OH)2, Mg(OH)2

6. Arrhenius Acids
Acids that increase the concentration of hydronium (H3O+) in aqueous solutions
HNO3(aq) + H2O(l) H3O+(aq) + NO3-(aq)
H+ + NO3- + H2O
acid

7. Why do acids produce H3O+?
H+ is extremely attracted to the unshared pair of electrons on the water molecule so it donates itself to this molecule where it becomes covalently bonded. The ion formed is known as the hydronium ion (H3O+) H+

8. NaOH(s) Na+(aq) + OH-(aq)
Arrenius Bases
Bases that increase the concentration of hydroxide ions (OH-) in aqueous solutions
NaOH(s) Na+(aq) + OH-(aq)
H2O

9. Strength of Acids & Bases
Strong acids & bases completely ionize in aqueous solutions
H2SO4 + H2O H3O+ + HSO4-
NaOH Na+ + OH-
Strong acids & bases are strong electrolytes
A list of strong acids & bases can be found on pg

10. Weak acids & bases only partially break down into ions when in aqueous solutions
HCN + H2O H3O+ + CN-
NH3 + H2O NH4+ + OH-
• Weak acids & bases are weak electrolytes
A list of weak acids & bases can be found on pg

11. H2O(l) + H2O(l) H3O+(aq) + OH-(aq)
Why can we drink H2O?
Water self ionizes to form equal concentrations of H3O+ and OH-
H2O(l) + H2O(l) H3O+(aq) + OH-(aq)
A substance is considered “neutral” when [H3O+] = [OH-]
[H3O+] concentration = 1.0 x 10-7M
[OH-] concentration = 1.0 x 10-7 M

12. 1.0 x 10-14 M2 = ionization constant for H2O (Kw)
When [H3O+] = [OH-]
If [H3O+] > 1.0 x 10-7 M, the solution is acidic
If [OH-] > 1.0 x 10-7 M, the solution is basic
To find the concentration of [H3O+] or [OH-] in acidic or basic solutions, the following equation can be used:
1.0 x M2 = [H3O+] [OH-]
1.0 x M2 = ionization constant for H2O (Kw)

13. Sample Problem
A 1.0 x 10-4 M solution on HNO3 has been prepared for laboratory use.
a. Calculate the [H3O+] of this solution
b. Calculate the [OH-] of this solution
c. Is this solution acidic or basic? Why?
d. Substitute H2SO4 as the acid. How would the calculations change?

14. Sample Problem
An aqueous 3.8 x 10-3 M NaOH solution has been prepared for laboratory use.
a. Calculate the [H3O+] of this solution
b. Calculate the [OH-] of this solution
c. Is this solution acidic or basic?
Why?
d. Substitute Ca(OH)2 as the base. How would the calculations change?

15. Practice Problems
Complete practice problems on pg. 484 #1-4

16. pH = -log [H3O+] [H3O+] = antilog (-pH)
The pH scale
The pH scale measures the power of the hydronium ion [H3O+] in a solution
The scale typically goes from 1-14 (although it can extend below or above it under extreme conditions)
The following equations can be used to determine the pH or [H3O+] of a solution:
pH = -log [H3O+] [H3O+] = antilog (-pH)
[H3O+] = 1 x 10-pH

17. pH > 7 basic
pH = 7 neutral
pH < 7 acidic

18. pOH = -log [OH-] [OH-] = antilog (-pOH)
The pOH scale
The pOH scale measures the power of the hydroxide ion [OH-] in a solution
The scale typically goes from 1-14 (although it can extend below or above it under extreme conditions)
The following equations can be used to determine the pOH or [OH-] of a solution:
pOH = -log [OH-] [OH-] = antilog (-pOH)
[OH-] = 1 x 10-pOH

19. pH + pOH = 14

20. Sample Problems
Calculate the pH of each of the following. Classify as acidic or basic.
1.3 x 10-5 M NaOH
1.0 x 10-4 M HCl

21. Sample Problems
What is the [H3O+] for each of the following? Classify as acidic or basic.
pH = 5.8
b. pOH = 8.9

22. Sample Problems
What is the [OH-] for each of the following? Classify as acidic or basic.
[H3O+] = 9.5 x M
pOH = 1.3

23. Practice Problems Complete practice problems on pg. 487 #1

24. Strong Acid-Base Neutralization
When equal parts of acid and base are present, neutralization occurs where a salt and water are formed
HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l)

25. Sample Problems H2CO3 + Sr(OH)2 HClO4 + NaOH HBr + Ba(OH)2
NaHCO3 + H2SO4

26. Titrations
When you have a solution with an unknown concentration, you can find it by reacting it completely with a solution of known concentration
This process is known as titrating
To perform a titration, an instrument called a buret can be used to precisely measure amounts of solution, drop by drop

28. Titration Termonology
Equivalence point - the point at which the known and unknown concentration solutions are present in chemically equivalent amounts
moles of acid = moles of base
Indicator - a weak acid or base that is added to the solution with the unknown concentration before a titration so that it will change color or “indicate” when in a certain pH range (table 16-6 on pg. 495 in your text will show various indicators and their color ranges)

29. Phenolpthalein is clear at pH<8,
End point - the point during a titration where an indicator changes color
The 2 most common indicators we will use in our chemistry class will be:
Phenolphthalein - turns very pale pink at a pH of 8-10
Bromothymol blue - turns pale green at a pH of
Phenolpthalein is clear at pH<8,
pale pink at pH 8-10 and
magenta at pH >10
Bromothymol blue

30. Practice Titration for an unknown acid
1. Titrate 5.0 of mL of unknown HCl into a 250 mL erlenmeyer flask - *remember to document the starting amount and ending amount of acid on the buret to prevent error
2. Add 2 drops of indicator (phenolphthalein) to the flask - the color of the solution should be clear
3. Titrate with .5M NaOH, continuously swirling the flask, until the solution turns very pale pink for 30 seconds - *remember to document the starting amount and ending amount of base on the buret
4. Mathematically determine the concentration of the unknown HCl solution by using the following equation:

31. Titration Equation MAVA = MBVB MA = molarity (mol/L) of acid
VA = volume in L of acid
MB = molarity (mol/L) of base
VB = volume in L of base
molesA = molesB
5. After calculating the molarity of the unknown acid experimentally, get the theoretical molarity and calculate % error

32. Practice titration for an unknown base
1. Titrate 5.0 of mL of unknown NaOH into a 250 mL erlenmeyer flask - *remember to document the starting amount and ending amount of base on the buret to prevent error
2. Add 2 drops of indicator (phenolphthalein) to the flask - the color of the solution should be magenta
3. Titrate with .5M HCl, continuously swirling the flask, until the solution turns very pale pink for 30 seconds - *remember to document the starting amount and ending amount of acid on the buret
4. Mathematically determine the concentration of the unknown NaOH solution by using MAVA = MBVB
5. After calculating the molarity of the unknown base experimentally, get the theoretical molarity and calculate % error