1. Operational Definition - (Properties in which they differ)
ACIDS AND BASES
What are they?
There are different definitions for acids and bases dependent on the circumstances.
Operational Definition - (Properties in which they differ)

2. Acids
1. react with carbonates to produce carbon dioxide gas.
2. change pink phenolphthalein to colourless
3. make litmus paper red
4. make bromothymol blue turn yellow
5. taste sour
6. have a pH below 7
7. react with most metals and produce hydrogen gas

3. Bases
1. don’t react with carbonates to produce carbon dioxide gas.
2. change colourless phenolphthalein to pink
3. make red litmus paper turn blue
4. make bromthymol blue turn blue
5. taste bitter
6. have a pH above 7

4. A lemon tastes sour. Is it an acid or a base?

5. 7.0
Is this solution an acid or a base?
1.8
Acid
pH meter

6. Conceptual Definitions of Acids and Bases
A chemist named Arrhenius recognized acids were molecular compounds and as such didn't conduct electricity as liquids since they didn't release ions. But he observed that when acids were combined with water they did conduct electricity.
He also observed bases conducted electricity as liquids and when they were combined with water.
Bases were ionic but acids weren't.

7. Arrhenius’s Definition of Acids and Bases
Acids are substances which react in water and produce hydronium ions.
HCl(g) + H > H301+(aq) + Cl1-(aq)
Bases are substances which react with water and produce hydroxide ions.
NH3(g) + H > NH41+(aq)+ OH1-(aq)

8. This concept has its limitations however
This concept has its limitations however. Can’t substances be classified as acids or bases without the involvement of water?

9. Bronstead's and Lowry's Definition of Acids and Bases
Acids are substances which donate protons and bases are substances which accept protons.
In the examples above HCl(g) is an acid because it donates protons to H2O molecules and NH3 is a base because it accepts protons from H2O molecules.

10. Conjugate Acid - Base Pairs - When using the Bronsted concept for acids and bases it is convenient to consider all acid - base reactions as reversible equilibria.
For instance when sulfurous acid, H2SO3 reacts with water the following equilibrium is established:
acid base acid base
H2SO H2O H HSO31-
conjugate pair
conjugate pair

11. is the proton acceptor so it’s the base. In
conjugate pair
H2SO H2O H HSO31-
acid base acid base
conjugate pair
In the forward direction the H2SO3 is the proton donor so it’s the acid and the H2O
is the proton acceptor so it’s the base. In
the reverse direction the H301+ is the proton donor so it’s the acid and the HSO31- is the proton acceptor so it’s a base.

12. conjugate pair
acid base acid base
H2SO H2O H HSO31-
conjugate pair
When looking at both forward and reverse reactions it is easy to pick out a pair of molecules which differ by a single proton (H atom without its electron). These pairs are called conjugate acid-base pairs.

13. Why do acids of equal concentration have different levels of conductivity?
Some acids are stronger than others.
Why?

14. Strength of Acids and Bases is determined by the degree to which a substance produces ions in solution.
A strong acid or base is a substance which completely ionizes. In other words if 100 molecules of a strong acid like HCl are placed in water all 100 of them will react with H2O producing 100 H3O1+ ions and 100 Cl1- ions.
Weak acids and bases only partially ionize.
Strong Acid - the reaction below goes to completion.
HCl(g) + H > H301+(aq) + Cl1-(aq)

15. Weak Acid - the reaction occurs to a limited extent
Weak Acid - the reaction occurs to a limited extent. In the example below if 100 acetic acid molecules are placed in water only a few of them will successfully react with water molecules producing hydronium ions. Most CH3COOH molecules remain intact.
CH3COOH + H H301+(aq)+ CH3COO1-(aq)

16. Strong Acids in order of decreasing strength are HClO4, HI, HBr, H2SO4, HCl, HNO3
Acid strength has to do with the ease with which an acid can lose a proton. If the binary acid strengths (HI, HBr, HCl) are compared it can be seen that HI is the strongest acid of this group because its iodide ion is the largest of the group so the force between the hydrogen ion and the iodide ion is the weakest so it loses its proton most easily.

17. Force is strongest since the ions are closest
Br1-
Remember the weaker the force the stronger the acid
H1+
Force is weakest since the ions are furthest
I1-
H1+

18. Strong Bases include hydroxides of group 1A and Ca2+, Ba2+, and Sr2+.
A table with the remaining moderate and weak bases can be found in text. As with acids the weaker the bonds, the stronger the base since liberation of OH1- ions is easiest when the bonds are weakest.

19. Polyprotic Acids donate protons in steps
Polyprotic Acids donate protons in steps. For instance carbonic acid, H2CO3 has two protons to donate and it does this in two steps:
step 1
H2CO3 + H20 HCO H301+
step 2
HCO H20 CO H301+
note: The arrows are constructed in this manner to show the reverse reaction has a greater tendency than the forward reaction.

20. Amphoteric (Amphiprotic) Substances can behave as both acids or bases dependent on the circumstances. Water molecules, for instance, can sometimes except protons and behave as bases or donate protons and behave as acids.
HBr(g) + H2O H301+(aq) Br1-(aq)
base
NH3(g) H2O 0H1-(aq) NH41+(aq)
acid

21. The pH Scale
Aqueous solutions can be classified as acidic, basic, or neutral. This classification scheme is based on the quantities of 2 ions, hydronium ion, H301+ and hydroxide ion, OH1-. Where do these ions come from in solutions of pure water?
Water molecules in motion will randomly collide with one another. When this happens occasionally a hydrogen nucleus from one molecule will be transferred from one molecule to the other. This can be illustrated.

22. H2O + H20 H3O1+(aq) + OH1-(aq) H20 H1+(aq) + OH1-(aq)
Hydronium ion (H3O1+)
Hydroxide ion (OH1-)
H2O + H H3O1+(aq) + OH1-(aq)
Which is usually shortened to:
H H1+(aq) + OH1-(aq)

23. In mathematical terms
pH = -log[H1+]
so if in an aqueous solution the
[H1+] = 2.4 x 10-8, the pH is
7.62
Remember the whole number portion of a pH doesn’t count as a significant digit (SD), just like in the number 2.4 x 10-8 the exponent -8 doesn’t count as a SD.

24. If the pH of a solution is 1.45 find the [H1+].
[H1+] = 10-pH.
[H1+] =
3.5 x 10-2 mol/L

25. pH = -log[H1+] pH = -log 6.2 x 10-2 pH = 1.21
If the [H1+] of a solution is
6.2 x 10-2 mol/L find the pH.
pH = -log[H1+]
pH = -log 6.2 x 10-2
pH = 1.21

26. pH Scale
Increasing basicity
Increasing acidity N e u t r a l
How much more acidic is pH 1 than pH 5?
x’s
What pH is 1000x’s more acidic than pH 2?
Since this is a logarithmic scale pH 9 is 10x’s more basic than pH 8, pH 12 is 1000 x’s more basic than pH 9
pH of -1

27. Acid
Base
Titrations

28. Acids and bases, when combined in equal quantities, neutralize each other forming salt and water.1+ 1- 1+ 1- 1+ 1- 1+ 1-
HCl + NaOH
HOH + NaCl 1+ 2- 3+ 1- 1+ 1- 3+ 2- 3
H2SO4 + Al(OH)3 2 6
HOH + Al(SO4) 2 3
This neutralization can be used to determine the concentrations or molar masses of unknowns.
If the right indicator is placed in an acid or a base it will turn colour at the instant of neutralization.

29. Estimating Volumes Estimate - 0.77 mL 0.80 0.70 0.00 0.72 0.74 1.00
0.76
0.78
2.00

30. Measurements with burets must have 2 decimal places.
Estimating Volumes
Measurements with burets must have 2 decimal places.
Don’t record 1.1 instead record
1.10
The extra zero tells us the measuring instrument measures to the nearest tenth of a mL.
The last digit of any measurement is an estimated value. mL
0.00
1.00
2.00

31. When the solution turns pink the number
# of mol of acid = # of mol of base
equivalence pt.
Titrant - NaOH solution
22.52 mL
4.21 mL
18.31 mL
Measured vinegar solution
HC2H3O2
few drops phenolphthalein