1. The Chemistry of Acids and Bases
To play the movies and simulations included, view the presentation in Slide Show Mode.

2. Acid and Bases They are everywhere
Acid and Bases They are everywhere.. In your food, in the air, in your house EVEN IN YOU!!!!!

3. Acid and Bases

4. Acid and Bases

5. Some Properties of Acids
Produce H+ (as H3O+) ions in water (the hydronium ion is a hydrogen ion attached to a water molecule)
Taste sour
Corrode metals
Form Electrolytes (conduct electricity) when dissolved in water
React with bases to form a salt and water
pH is less than 7
Turns blue litmus paper to red “Blue to Red A-CID”
Feel sticky

6. Some Properties of Bases
Produce OH- ions in water
Taste bitter, chalky
Are electrolytes in water
Feel soapy, slippery
React with acids to form salts and water
pH greater than 7
Turns red litmus paper to blue “Basic Blue”

7. Some Common Acids and Bases
HCl hydrochloric acid stomach acid
HC2H3O2 acetic acid vinegar
H3C6H5O7 citric acid citrus fruits
KOH potassium hydroxide liquid soap
Mg(OH)2 magnesium hydroxide Milk of magnesia
Al(OH)3 aluminum hydroxide antacid

8. Acid/Base definitions
Definition #1: Arrhenius (original)
Acids – produce H+ ions (or hydronium ions H3O+)
Bases – produce OH- ions
(problem: some bases don’t have hydroxide ions!)

9. HCl + H2O  H3O+ + Cl– – + Acids form hydronium ions (H3O+) acid
Arrhenius - In aqueous solution…
Acids form hydronium ions (H3O+)
HCl + H2O  H3O+ + Cl– H Cl O – +
acid

10. NH3 + H2O  NH4+ + OH- – + Bases form hydroxide ions (OH-) base
Arrhenius - In aqueous solution…
Bases form hydroxide ions (OH-)
NH3 + H2O  NH4+ + OH- H N O – +
base

11. Acid/Base Definitions
Definition #2: Brønsted – Lowry
Acids – proton donor
Bases – proton acceptor
A “proton” is really just a hydrogen atom that has lost its electron!

12. A Brønsted-Lowry acid is a proton donor
A Brønsted-Lowry base is a proton acceptor
conjugate acid
conjugate base
base
acid
Conjugate What??!! We’ll go over this later, but think about what it might mean…..

13. Summary of Definitions
Arrhenius acid: Produce H+ ion
Arrhenius base: Produce OH- ion
Bronsted-Lowery Acid: Proton donor
** Bronsted Lowery Base: Proton acceptor
**More Broad definition than Arrhenius**

14. How do you Identify an Acid or a Base?
The chemical formula of an acid usually begins with an H
HCl (hydrochloric acid)
HNO3 (nitric acid)
HC2H3O2 (acetic acid)
The chemical formula of a base usually includes an –OH group
NaOH (sodium hydroxide)
Mg(OH)2 (magnesium hydroxide)
NOT TRUE for the common base, ammonia NH3

15. Acid Nomenclature Flowchart

16. Acid Nomenclature Review
Binary: No Oxygen
Ternary: w/Oxygen
He –ate –ic and got stomach –ite -ous

17. Acid Nomenclature Review
HBr (aq)
H2CO3
H2SO3
 hydrobromic acid
 carbonic acid
 sulfurous acid

18. Name ‘Em!
HI (aq)
HCl (aq)
H2SO3
HNO3
HIO4

19. Naming and Writing Formulas for Bases
Bases are usually IONIC: metal + hydroxide
Remember: NO PREFIXES WITH IONICS
Example:
NaOH
Sodium hydroxide
Mg(OH)2
- Magnesium hydroxide
Most common non-hydroxide Base:
Ammonia, NH3

20. Conjugate Acid/Base Pairs
A conjugate base is what is left over after an acid donates a proton. A conjugate acid is what is formed when a base accepts a proton.
HCl + H2O  Cl– + H3O+
acid base conjugate conjugate
base acid

21. Conjugate Acid/Base Pairs
HA B A BH
-H+
+H+
acid = proton donor
base = proton acceptor

22. Conjugate Acid/Base Pairs
Identify the acid, base, conjugate acid, conjugate base, and conjugate acid-base pairs:
HC2H3O2(aq) + H2O(l)  C2H3O2–(aq) + H3O+(aq)
acid
base
conjugate base
conjugate acid
conjugate acid-base pairs
OH –(aq) + HCO3–(aq)  CO32–(aq) + H2O(l)
base
acid
conjugate base
conjugate acid
conjugate acid-base pairs

23. You Try (a) HF(aq) + SO32–(aq)  F–(aq) + HSO3–(aq) (b)
CO32–(aq) + HC2H3O2(aq)  C2H3O2–(aq) + HCO3–(aq)
(c)
H3PO4(aq) + OCl –(aq)  H2PO4–(aq) + HOCl(aq)

24. Strength of an Acid or Base
Strong Acids/Bases
100% ionized in water
strong electrolyte
Remember: any solution
that contains ions is an
electrolyte: conducts
Electricity - +
Strong Acids
HCl
HNO3
H2SO4
HBr HI
HClO4
Strong Bases
NaOH
KOH
Ca(OH)2
Ba(OH)2

25. Weak Acids/Bases
do not ionize (dissociate) completely; only a small percentage of the original acid or base molecules will dissociate. The majority stay intact (do NOT ionize) and therefore don’t make many ions.
Weak acids and bases are also weak electrolytes, meaning they only conduct electricity a little.
Weak Acids HF
CH3COOH
H3PO4
H2CO3
HCN
Weak Base
NH3

26. Acid Strength is described using the Acid Dissociation Constant Ka
Acid dissociation constant, Ka= concentration of products
concentration of reactants
The more ionization occurs, the larger Ka, and the stronger the acid!
HA + H2O
H3O+ + A-
Ka =
[H3O+] [A-]
[HA]

27. Acid Equilibrium Constants are greater than 1(›1) for strong acids and less than 1 (‹1) for weak acids

28. Equilibrium Constants for Weak Bases Kb = concentration of products concentration of reactants
Weak base has Kb < 1

29. Ionization Constants for Acids/Bases
Conjugate
Bases
Increase strength
Increase strength

30. Strength of Acids and Bases
The terms weak and strong are used to describe the ability of an acid or base to ionize (produce H+ or OH- ions). A strong acid ionizes completely; a weak acid only partially ionizes.
Dilute and concentrated are terms used to describe how much of an acid or base (concentration) is present in a solution.

31. from a high concentration of a weak acid.
Acids and Bases
Strength Is NOT Concentration
pH of a solution tells us the concentration of H+ ions in solution. A high concentration of H+ ions can result from
a strong acid OR
from a high concentration of a weak acid.
So… you really need Ka to know how strong an acid is.

32. The pH scale is a way of expressing the strength of an acidic or basic solution. Under 7 = acidic = neutral Over 7 = basic

33. pH of Common Substances

34. [ ] means Concentration (Molarity)
Calculating pH
pH = - log [H+]
[ ] means Concentration (Molarity)
Example: If [H+] = 1 X M pH = - log (1 X 10-10)
pH = - (- 10)
pH = 10
Example: If [H+] = 1.8 X 10-5 M pH = - log (1.8 X 10-5)
pH = - (- 4.74)
pH = 4.74

35. Try These! Find the pH of these:
1) A 0.15 M solution of Hydrochloric acid
2) A 3.00 X 10-7 M solution of Nitric acid

36. pH calculations – Solving for H+
If the pH of Coke is 3.12, [H+] = ???
Because pH = - log [H+] then
- pH = log [H+]
Take antilog (10x) of both sides and get
10-pH = [H+]
[H+] = = 7.6 x 10-4 M
*** to find antilog on your calculator, look for “Shift” or “2nd function” and then the log button or the 10x button, where x is -pH

37. pH calculations – Solving for H+
A solution has a pH of What is the Molarity of hydrogen ions in the solution?
pH = - log [H+]
8.5 = - log [H+]
-8.5 = log [H+]
Antilog -8.5 = antilog (log [H+])
= [H+]
3.16 X 10-9 = [H+]

38. pH testing There are several ways to test pH
Blue litmus paper (red = acid)
Red litmus paper (blue = basic)
pH paper (multi-colored)
pH meter (7 is neutral, <7 acid, >7 base)
Universal indicator (multi-colored)
Indicators like phenolphthalein: basic is pink (backwards!) and acidic is clear.

39. pOH Since acids and bases are opposites, pH and pOH are opposites!
pOH looks at the perspective of a base
pOH = - log [OH-]
If you know either pH or pOH, you can find the other easily!
pH + pOH = 14

40. Relationship between pH, pOH, [H+] and [OH-]

41. [H3O+], [OH-] and pH What is the pH of a 0.0010 M NaOH solution?
Be careful! NaOH is a base and produces OH- ions.
[OH-] = (or 1.0 X 10-3 M)
pOH = - log
pOH = 3
pH = 14 – 3 = 11

42. The pH of rainwater collected in a certain region of the northeastern United States on a particular day was What is the H+ ion concentration of the rainwater?
The OH- ion concentration of a blood sample is 2.5 x 10-7 M. What is the pH of the blood?

43. Paper testing Paper tests like litmus paper and pH paper
Put a stirring rod into the solution and stir.
Take the stirring rod out, and place a drop of the solution from the end of the stirring rod onto a piece of the paper
Read and record the color change. Note what the color indicates.
You should only use a small portion of the paper.

44. pH meter Tests the voltage of the electrolyte (ionic solution)
Converts the voltage to pH
Very cheap, accurate
Must be calibrated with a buffer solution

45. pH indicators
Indicators are dyes that can be added that will change color in the presence of an acid or base.
Some indicators only work in a specific range of pH
Phenolphthalein is the most common indicator

46. Acids and Bases

47. Relation of Ka, Kb, [H3O+] and pH

48. ACID-BASE REACTIONS; a type of Double-Replacement Reaction
Acid Base → Water and salt
HCl + NaOH → H2O + NaCl
Acid-Base reactions are called neutralization reactions because the products are water and a salt (ionic compound).

49. You Try…
Write balanced chemical equations for these neutralization reactions. Under each compound give the correct formula or IUPAC name.
iron(II) hydroxide + phosphoric acid
Ba(OH)2(aq) + HCl(aq)
calcium hydroxide + nitric acid
Al(OH)3(aq) + H2SO4(aq)
ammonium hydroxide + hydrosulfuric acid
KOH(aq) + HClO2(aq)