1. ACIDS AND BASES

2. Properties of Acids Acids are proton (hydrogen ion, H+) donors
Acids have a pH lower than 7
Acids taste sour
Acids effect indicators
Blue litmus turns red
Methyl orange turns red
Acids react with active metals, producing H2
Acids react with carbonates
Acids neutralize bases

3. Acids are Proton (H+ ion) Donors
Strong acids are assumed to be 100% ionized in solution (good H+ donors).
HCl
H2SO4
HNO3
Weak acids are usually less than 5% ionized in solution (poor H+ donors).
H3PO4
HC2H3O2
Organic acids

4. Acids Have a pH less than 7

5. Acids Effect Indicators
Blue litmus paper turns red in contact with an acid.
Methyl orange turns red with addition of an acid

6. Acids React with Active Metals
Acids react with active metals to form salts and hydrogen gas.
Mg + 2HCl  MgCl2 + H2(g)
Zn + 2HCl  ZnCl2 + H2(g)
Mg + H2SO4  MgSO4 + H2(g)

7. Acids React with Carbonates
2HC2H3O2 + Na2CO3
2 NaC2H3O2 + H2O + CO2

8. Effects of Acid Rain on Marble (calcium carbonate)
George Washington:
BEFORE
George Washington:
AFTER

9. Acids Neutralize Bases
Neutralization reactions ALWAYS produce a salt and water.
HCl + NaOH  NaCl + H2O
H2SO4 + 2NaOH  Na2SO4 + 2H2O
2HNO3 + Mg(OH)2  Mg(NO3)2 + 2H2O

10. Properties of Bases Bases are proton (hydrogen ion, H+) acceptors
Bases have a pH greater than 7
Bases taste bitter
Bases effect indicators
Red litmus turns blue
Phenolphthalein turns purple
Solutions of bases feel slippery
Bases neutralize acids

11. Bases are Proton (H+ ion) Acceptors
Sodium hydroxide (lye), NaOH
Potassium hydroxide, KOH
Magnesium hydroxide, Mg(OH)2
Calcium hydroxide (lime), Ca(OH)2
OH- (hydroxide) in base combines with H+ in acids to form water
H+ + OH-  H2O

12. Bases have a pH greater than 7

13. Bases Effect Indicators
Red litmus paper turns blue in contact with a base.
Phenolphthalein turns bright pink in a base.

14. Bases Neutralize Acids
Milk of Magnesia contains magnesium hydroxide, Mg(OH)2, which neutralizes stomach acid, HCl.
2 HCl + Mg(OH)2
MgCl2 + 2 H2O

15. Titrations
Main Idea: Titrations are an application of acid-base neutralization reactions that require the use of an indicator.

16. NaOH (aq) + HCl (aq)  NaCl (aq) + H2O (l)
Stoichiometry
The stoichiometry of an acid-base neutralization reaction is the same as that of any other reaction that occurs in solution (they are double displacement reactions, after all).
For example, in the reaction of sodium hydroxide and hydrogen chloride, 1 mol of NaOH neutralizes 1 mol of HCl:
NaOH (aq) + HCl (aq)  NaCl (aq) + H2O (l)
Stoichiometry provides the basis for a procedure called titration, which is used to determine the concentrations of acidic and basic solutions.

17. Titration
Titration is a method for determining the concentration of a solution by reacting a known volume of that solution with a solution of known concentration.
If you wish to find the concentration of an acid solution, you would titrate the acid solution with a solution of a base of known concentration.
You could also titrate a base of unknown concentration with an acid of known concentration.

18. In the titration of an acid by a base, the pH meter measures the pH of the acid solution in the beaker as a solution of a base with a known concentration is added from the buret.

19. How is an acid-base titration performed?
The figure on the previous slide illustrates one type of setup for the titration procedure outlined on the next slide.
In the procedure pictured on Slide 4, a pH meter is used to monitor the change in pH as the titration progresses.

20. Titration Procedure
A measured volume of an acidic or basic solution of unknown concentration is placed in a beaker. The electrodes of a pH meter are immersed in this solution, and the initial pH of the solution is read and recorded.
A buret is filled with the titrating solution of known concentration. This is called the standard solution, or titrant.
Measured volumes of the standard solution are added slowly and mixed into the solution in the beaker. The pH is read and recorded after each addition. This process continues until the reaction reaches the equivalence point, which is the point at which moles of H+ ion from the acid equal moles of OH- ion from the base.

21. In the titration of a strong acid by a strong base, a steep rise in the pH of the acid solution indicates that all of the H+ ions from the acid have been neutralized by the OH- ions of the base. The point at which the curve flexes is the equivalence point of the titration. Bromthymol blue is an indicator that changes color at this equivalence point. Notice that phenolphthalein and methyl red don’t match the exact equivalence point, but the slope is so steep that it doesn’t matter.

22. Strong-Strong Titration
The previous slide shows how the pH of the solution changes during the titration of 50.0 mL of M HCl, a strong acid with M NaOH, a strong base.
The inital pH of the M HCl is 1.00.
As NaOH is added, the acid is neutralized and the solution’s pH increases gradually.
When nearly all of the H+ ions from the acid have been used up, the pH increases dramatically with the addition of an exceedingly small volume of NaOH.
This abrupt change in pH occurs at the equivalence point of the titration.
Beyond the equivalence point, the addition of more NaOH again results in the gradual increase in pH.

23. Strong Acid – Strong Base Titrations
You can calculate the volume of base needed to reach the equivalence point using the formula: V·M = V·M.
There are three situations in which you determine pH.
Initial strong acid concentration (this is simply the –log[H+] which is based on the [Acid].)
Equivalence point (or endpoint) when moles of OH- = moles of H+. The pH is 7 (due to the auto-ionization of water.)
Before and after the endpoint (calculate excess moles of H+ or OH-, divide by the total volume, and calculate the pH based on this value.)

24. Indicators and Titration End Point
Many indicators used for titrations are weak acids.
Each has its own particular pH or pH ranges over which it changes color.
The point at which the indicator used in a titration changes color is called the end point of the titration.
It is important to choose an indicator for a titration that will change color at the equivalence point of the titration.
Remember that the role of the indicator is to indicate to you, by means of a color change, that just enough of the titrating solution has been added to neutralize the unknown solution.
Equivalence point ≠ End point!
BUT for strong-strong titrations, the pH change is so steep and so large, that the are approximately equal.

25. Acid-Base Indicators
Chemists often use a chemical dye rather than a pH meter to detect the equivalence point of an acid-base titration.
Chemical dyes whose colors are affected by acidic and basic solutions are called acid-base indicators.
Many natural substances act as indicators.
If you use lemon juice in your tea, you might have noticed that the brown color of tea gets lighter when lemon juice is added.
Tea contains compounds called polyphenols that have slightly ionizable hydrogen atoms and therefore are weak acids.
Adding acid in the form of lemon juice to a cup of tea lessens the degree of ionization, and the color of the un-ionized polyphenols becomes more apparent.
Chemists have several choices in selecting indicators.
Bromthymol blue is a good choice for the titration of a strong acid with a strong base, and phenolphthalein changes color at the equivalence point of a titration of a weak acid with a strong base.

26. The equivalence point here is not at a pH of 7
The equivalence point here is not at a pH of 7. Phenolphthalein is an indicator that changes color at this equivalence point. Notice that the starting pH is different and the region of change is smaller.

27. Strong-Weak Titrations
You might think that all titrations must have an equivalence point at pH 7 because that is the point at which concentrations of hydrogen ions and hydroxide ions are equal and the solution is neutral.
This is not the case – some titrations have equivalence points at pH values less than 7, and some have equivalence points at pH values greater than 7.
These differences occur because of reactions between the newly formed salts and water – salt hydrolysis.
Some salts are basic (weak acid, strong base) and some salts are acidic (strong acid, weak base).
The previous slide shows that the equivalence point for the titration of HPr (a weak acid) with NaOH (a strong base) lies at pH 8.80.

28. Titration of a weak acid with a strong base
Such as CH3COOH and NaOH
Equivalence point occurs at a pH >7
Should use an indicator that changes at a pH above 7, such as phenolphthalein.
Notice that starting pH is higher and
the Equivalence Point pH is higher.

29. Why is the Equivalence point at a pH >7?
The conjugate base of a weak acid is a strong base.
CH3COOH = Weak acid
CH3COO- = Strong Base
Once all of the acid has been converted to CH3COO-, it starts to take the H+ from H2O in solution, creating more OH- ions thus making the pH >7

30. Weak Acid – Strong Base Titrations
When a weak acid (such as HC2H3O2) is neutralized by a strong base (such as NaOH), the graph varies in two ways:
the equivalence point is not at pH = 7 and
a buffer region exists as you approach the endpoint.
You can still calculate the volume of base needed to reach the equivalence point using the formula: V·M = V·M. Weak acids require the same amount of base for neutralization as strong acids because they dissociate as they are neutralized

31. Weak Acid – Strong Base Titrations
Buffer regions
Halfway to equivalence point
pH = pKa

32. There are five situations in which you need to be able to calculate the pH.
1. Initial weak acid concentration (this is an ICE box calculation.) The shortcut can be used here.
2. Equivalence point (endpoint) is when all of the weak acid has been neutralized and turned into the conjugate base (C2H3O2- in this case.) This is a hydrolysis problem. Calculate the [C2H3O2-] and then do an ICE box problem knowing that Kb = . Calculate the [OH-], the pOH, and then the pH.
3. Halfway to the equivalence point (as in a half-titration) the pH = pKa. This is because at this point, there is a perfect buffer as the [HA] = [A-]. At this point, you can determine the Ka of an unknown weak acid… very useful.

33. 4. Before and after the half-way point, the pH can be calculated using the Henderson-Hasselbach equation (or an ICE box, if you want.) Use stoichiometry to determine the [HA] and [A-] pH = pKa + log
5. Finally, after the equivalence point, the pH depends on the excess strong base that has been added. As in the strong acid-strong base titration, calculate excess moles of OH-, divide by the total volume, and calculate the pOH and then pH based on this value. The effect on the pH by the A- is negligible compared to the excess OH-.

34. Weak Base – Strong Acid Titrations
When a sample of a weak base is titrated with a strong acid, the curve resembles an inverted Weak Acid – Strong Base titration curve.
Note that the pH at the equivalence point is less than 7.
An indicator such as phenolphthalein that changes at pH of 9 would change when only 6 mL of acid had been added even though the equivalence point is reached at around 11 mL.
The acid-base indicator must be chosen with a Ka near to the [H+] of the equivalence point; that is the pKa of the indicator must match the pH of the equivalence point.

35. Weak Diprotic Acid – Strong Base Titrations
When a weak diprotic acid (examples: H2C2O4 or H2CO3) is titrated, there are two equivalence points.
The curve is not as distinct because of the various proton donors and proton acceptors in solution.

36. Titration with an Indicator

37. What’s the Point of a Titration Again?
To find the unknown concentration of an acid or a base.
So you perform the actual titration noting the volume you started with and how much volume of the titrant you added and then...
Math! (Oh no! Not math! Anything but math!)

38. Titration Calculations: An Example
The balanced equation of a titration reaction is the key calculating the unknown molarity. For example, sulfuric acid is titrated with sodium hydroxide according to this equation:
H2SO4 (aq) + 2 NaOH (aq)  Na2SO4 (aq) + 2 H2O (l)
Calculate the moles of NaOH in the standard from the titration data: molarity of the base (MB) and the volume of the base (VB). In other words, MB VB = (mol/L)(L) = mol NaOH in standard
From the equation, you know that the mole ratio of NaOH to H2SO4 is 2:1. Two moles of NaOH are required to neutralize 1 mol of H2SO4. mol H2SO4 titrated = mol NaOH in standard x (1 mol H2SO4 / 2 mol NaOH)
MA represents the molarity of the acid and VA represents the volume of the acid in liters. MA = mol H2SO4 titrated/VA

39. In Short Form... mol acid = mol base (mol acid/mol base)
MAVA = MBVB (mol acid/mol base)
This is the mole ratio
Does this make sense? Let’s find out using the definition of molarity (mol/L) and dimensional analysis...
(mol acid/L acid)(L acid) = (mol base/L base)(L base) (mol acid/mol base)
mol acid = mol base (mol acid/mol base)
mol acid = mol acid