1. The Chemistry of Acids and Bases
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2. Acid and Bases

3. Acid and Bases

4. 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
Electrolytes
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”

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

6. Some Common Bases NaOH sodium hydroxide lye
KOH potassium hydroxide liquid soap
Ba(OH)2 barium hydroxide stabilizer for plastics
Mg(OH)2 magnesium hydroxide “MOM” Milk of magnesia
Al(OH)3 aluminum hydroxide Maalox (antacid)

7. Acid/Base definitions
Definition #1: Arrhenius (traditional)
Acids – produce H+ ions (or hydronium ions H3O+)
acids can be classified as monoprotic, diprotic or triprotic
Bases – produce OH- ions
problem: some bases don’t have hydroxide ions
Nice, simple definitions but has limitations…
restricted to aqueous solutions
couldn’t explain basic properties of NH3

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

9. A Brønsted-Lowry acid is a proton donor
A Brønsted-Lowry base is a proton acceptor
conjugate acid
conjugate base
base
acid

10. Conjugate Pairs

11. Bronsted-Lowry A-B
Because many A/B reactions are reversible, it is useful to identify conjugates (opposites) that will react in the reverse reaction
conjugate A-B pairs = 2 formulas in an equation whose formulas differ by a H+
strong A-B are not equilibrium expressions, but all other A-B are reversible

12. Acids & Base Definitions
Definition #3 – Lewis
Lewis acid - accepts an electron pair
Lewis base - donates an electron pair

13. Lewis Acid/Base Reaction

14. Lewis Acid-Base Interactions in Biology
The heme group in hemoglobin can interact with O2 and CO.
The Fe ion in hemoglobin is a Lewis acid
O2 and CO can act as Lewis bases
Heme group

15. Brønsted-Lowry vs. Lewis
All B/L bases are Lewis bases BUT, by definition, a B/L base cannot donate its electrons to anything but a proton (H+)
While B/L is most useful for our purposes, Lewis allows us to treat a wider variety of reactions (even if no H+ transfer occurs) as A/B reactions

16. Strong Acids/Bases
The strength of an acid (or base) is determined by the amount of IONIZATION (Not pH value).
HCl, HBr, HI, HNO3, H2SO4 HClO3 and HClO4 are the strong acids.

17. Strong Acids/Bases The strong bases are: alkali-metal hydroxides and
alkaline earth metal hydroxides.
Ex. NaOH, KOH, CsOH, Ca(OH)2, etc.
Note: The stronger the acid/base, the weaker its conjugate pair.

18. The pH scale is a way of expressing the strength of acids and bases
The pH scale is a way of expressing the strength of acids and bases. Instead of using very small numbers, we just use the NEGATIVE power of 10 on the Molarity of the H+ (or OH-) ion. Under 7 = acid = neutral Over 7 = base

19. pH of Common Substances

20. pH and acidity
Acidity or Acid Strength depends on Hydronium Ion Concentration [H3O+]
The pH system is a logarithmic representation of the Hydrogen Ion concentration (or OH-) as a means of avoiding using large numbers and powers.
pH = - log [H3O+]
pOH = - log [OH-]
In pure water [H3O+] = 1 x 10-7 mol / L (at 25oC)
 pH = - log(1 x 10-7) = - (0 - 7) = 7
pH range of solutions:
pH < 7 (Acidic) [H3O+] > 1 x 10-7 m / L
pH > 7 (Basic) [H3O+] < 1 x 10-7 m / L

21. 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

22. pOH Since acids and bases are opposites, pH and pOH are opposites!
pOH does not really exist, but it is useful for changing bases to pH.
pOH looks at the perspective of a base
pOH = - log [OH-]
Since pH and pOH are on opposite ends,
pH + pOH = 14

23. pH
[H+]
[OH-]
pOH

24. 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
Once the drops are added, the sample is ruined
Some dyes are natural, like radish skin or red cabbage

25. Indicators

27. Wait, water can go both ways?
amphoteric substances can behave as either an acid or base depending on what they react with.
water and anions with protons (H+) attached are most common amphoterics

28. Autoionization of Water
H2O + H2O OH- + H3O+
@ equilibrium Keq = [H3O+] [OH-] = so Keq becomes Kw and equals 1.00 x at 25 oC
In a neutral solution [H3O+] = [OH-]

29. Weak Acids/Bases Only partially ionize (less than 100%) in water.
Acids (symbol=HA)
@ equilibrium Keq = [H3O+] [OH-]/[HA] = so Keq becomes Ka
The bigger the Ka the stronger the acid
Polyprotic acids have more than 1 H+ available to ionize
Have multiple Ka values (Ka1 Ka2 Ka3)
Always easier to remove 1st H+ than 2nd
If Ka2 is different than Ka1 by 103 (or more) consider only Ka1 to get pH

30. Weak Acids/Bases Bases (symbol=B) Relationship between Ka and Kb
2 general categories
Contain neutral substances that have an atom with non-bonding electrons Ex. NH3  NH4+
Anions of weak acids Ex. ClO-  HClO
@ equilibrium Keq = [HB+] [OH-]/[B] = so Keq becomes Kb
Relationship between Ka and Kb
Ka x Kb = Kw
pKa + pKb = pKw = 14.00

31. Equilibrium Constants for Weak Acids
Weak acid has Ka < 1
Leads to small [H3O+] and a pH of 2 - 7

32. Equilibrium Constants for Weak Bases
Weak base has Kb < 1
Leads to small [OH-] and a pH of 7-12

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

34. Equilibria Involving A Weak Acid
You have 1.00 M HA. Calc. the equilibrium concs. of HA, H3O+, A-, and the pH.
HA + H2O H3O+ + A-
Step 1. Define equilibrium concs. in ICE table [HA] [H3O+] [A-]
initial
change
equilib
-x +x +x
1.00-x x x

35. Equilibria Involving A Weak Acid
Step 2. Solve equilibrium (Ka) expression
This is a quadratic. Solve using quadratic formula.
because Ka for most weak acids is less than 10-3, 1-x is about equal to 1, so Ka = x2/1.00

36. Equilibria Involving A Weak Acid
So the Ka expression
becomes
Now we can more easily solve this approximate expression.

37. Equilibria Involving A Weak Acid
Step 3. Calculate the pH
x = [H3O+] = [A-] = 4.2 x 10-3 M
pH = - log [H3O+] = -log (4.2 x 10-3) = 2.37

38. Equilibria Involving A Weak Base
You have M of B. Calc. the pH.
B + H2O OH- + HB+
Kb = 1.8 x 10-5
Step 1. Define equilibrium concs. in ICE table
[B] [HB+] [OH-]
initial
change
equilib
-x +x +x
x x x

39. Equilibria Involving A Weak Base
Step 2. Solve the equilibrium expression
b/c Kb for most weak bases is less than 10-3, x is about equal to 0.010, so Kb = x2/0.010
x = [OH-] = [HB+] = 4.2 x 10-4 M and [B] = x 10-4 ≈ M

40. Equilibria Involving A Weak Base
Step 3. Calculate pH
[OH-] = 4.2 x 10-4 M
so pOH = - log [OH-] = 3.37
Because pH + pOH = 14,
pH = 10.63

41. A-B properties of salt solutions
for the most part anions are slightly basic (because they attract protons) and cations are slightly acidic (because they can donate protons)
Anions
Look at strength of the acid to which it is a conjugate
If acid is strong, anion won’t really take H+ from water
If acid is weak, anion will react to a small extent ( pH)
If amphoteric need to look at magnitudes of Ka and Kb
Cations
Polyatomic cations can be the conjugate acid of a weak base
Usually react to make H3O+ (  pH)

42. A/B Behavior & Chemical Structure
Polarity
if H is the positive end = acid
if H is negative end (NaH) = base
if non-polar, neither acid or base
Strength of bond
Look at what H is bonded to
How easy is it to dissociate into ions?
Stability of conjugate: the greater the stability of the conjugate, the stronger the acid

43. A/B Behavior & Chemical Structure
Binary Acids
1. HX bond strength is most important
2. acidity increases down a group
- larger elements
- longer and weaker bonds
3. in same row, look at bond polarity
- across from L to Right increases EN
- Across from L to R increases acidity

44. A/B Behavior & Chemical Structure
Oxyacids (have O in compound)
1. Have same # of OH groups and the same # of O atoms, acid strength  with  in EN of center atom
2. Have same center atom (Y) acid strength  as the # of O atoms attached to Y 
3. in a series of oxyacids, acidity  with as the oxidation # of the center atom 

45. A/B Behavior & Chemical Structure
Carboxylic Acids (have –COOH in compound)
H on OH group is ionized
the other O atom draws elctron density away from OH bond increasing polarity
the conj. base (COO-) can exhibit resonance making it more stable
Acid strength of COOH group increases as the # of electronegative atoms in the acid increase.