1. Acids and Bases

2. Properties of Acids
In aqueous solutions, they conduct electricity They are ______________
Metals above H2 on Table J will react with acids to produce hydrogen gas and a salt
Will Mg react with HCl? _________
Will Cu react with HCl? _________
Acids cause color changes in acid-base indicators
Acids react with bases to form a salt + H2O _________
sour taste
pH levels __________
Common examples:___________________________
Table K

3. Properties of Bases
In aqueous solutions, bases conduct electricity  They are also ______________
Bases cause color changes in acid-base indicators.
Bases react with acids to form a salt + H2O!! ____________________________
bitter taste
pH levels ___________
Common examples: ____________________
Table L

4. Properties of Salts
Like acids and bases, salts are also _________________
Salts are formed as a product of neutralization reactions
Neutral  most have a pH of _______
Salts are ionic compounds
( _______+ _______)
Examples:

5. The Arrhenius Theory Svante Arrhenius
Based his ideas on the fact that aqueous solutions of acids and bases are electrolytes
said that the properties of acids and bases are because of their H+ and OH- ions

6. Arrhenius Acids
An acid is a substance that has H and releases (yields) H+ ions as the only positive ions in solution
Adding an acid to an aqueous solution ___________ the concentration of H+ ions in the solution
The H+ ions attach to H2O, forming hydronium ions _________
Examples: ___________________
Can be monoprotic or diprotic________

7. Arrhenius Bases
A base is a substance that has OH and releases (yields) OH- ions as the only negative ions in an aqueous solution
Adding a base to an aqueous solution __________ the concentration of OH- ions in the solution
Examples: __________________

8. Naming Acids, Bases and Salts
Bases and salts are ionic, and are named the usual way. Binary acids: hydro________ic acid Ternary acids: _________-ic acid (if poly is –ate) _______________ous acid (if poly is –ite)

9. The Bronsted-Lowry Theory
Arrhenius’ theory had some limitations, so they expanded the definitions of acids and bases
An acid is an H+ donor
A base is an H+ acceptor
H+ is simply a proton which interacts strongly with nonbonding e- on the oxygen of a water molecule
Interaction results in a hydronium ion H3O+

10. Bronsted-Lowry
Acids and bases can be defined in terms of ability to donate protons
Acid: proton donor
Base: proton acceptor

11. Bronsted- Lowry HCl(g) + H2O(l)  H3O+(aq) + Cl- (aq) EXAMPLE
HCl: _______ H2O: _______

12. Comparing the two theories
NH3(aq) + H2O(l)  NH4+(aq) + OH- (aq)
According to Arrhenius, NH3 is a base because adding it to water increases the hydroxide ion concentration
Bronsted-Lowry classifies it as a base because it accepts a proton from H2O, and the water acts as an acid because it donates a proton

13. Conjugate Acid-Base Pairs
In any acid-base equilibrium, both the forward and reverse reactions involve proton transfers Example: HX(aq) + H2O(l)  X- + H3O+ (aq) Forward reaction, HX donates a proton, it is an acid and H2O is the base Reverse reaction, H3O+ is the acid and X- is the base An acid and base that differ only in presence or absence of a proton is called a Conjugate acid-base pair

14. Conjugate Acid-Base Pairs
HNO2(aq) + H2O(l)  NO2- (aq) + H3O+(aq) What are the conjugate acid base pairs in the above reaction?

15. Amphoteric Substances
Substances that can act as an acid or base
Ex: H2O
With ammonia it’s an acid
With HCl it’s a base

16. Auto-ionization of Water
Because water is amphoteric, it can actually donate a proton to another water and vice-versa This is called auto-ionization No individual molecule remains ionized for long, and at room temperature, only about 1 in 109 molecules is ionized at any given moment

17. Auto-ionization of Water
Because it is an equilibrium process, we can write an equilibrium expression Keq= [H3O+][OH-] [H2O]2 This becomes the ion product constant for water (Kw) Kw at room temperature = 1 x [H+] = 1x10-7 [OH-] = 1x10-7

18. Auto-ionization of Water
A solution in which [H+] = [OH-] is considered neutral
[H+] > [OH-] acid
[H+] < [OH-] base
When the concentration of one increases, the concentration of the other decreases so their product is always
1 x 10-14

19. Calculations
Calculate the [H+] in a solution in which [OH-] is M [H+] = 1 x = 1 x = 1.0 x 10-12M [OH-] .010 M
Is the solution acidic or basic?

20. Calculations
Calculate the [H+] in a solution in which the [OH-] is 2.0 x 10-9 [H+] = 1 x = 1 x = 5.0 x 10-6 M [OH-] 2.0 x 10-9
Is the solution acidic or basic?

21. The pH scale
The concentration of [H+] in aqueous solution is usually very small, so the concentration is expressed as pH ( the negative logarithm in base 10) of [H+] pH = -log [H+] Calculate the pH of a solution with [H+] concentration of 1.0 x 10-7 M -log(1.0 x 10-7 M) = -(-7) = 7

22. Calculations with pH (regents)
Examples: Determine the pH given the [H+] concentration
1. If [H+] = 1 x pH = ___________
2. If [H+] = 1 x pH = ___________
3. If [H+] = 1 x pH = ___________
4. What is the pH of a M HCl solution? _________
5. What is the pH of a M solution? _________
6. What is the pH of a M solution? _________

23. Calculations with pH
A sample of lemon juice has [H+] = 3.8 x 10-4 M. What is the pH?
What is the pH of a window cleaning solution with an [H+] = 5.3 x 10-9
Calculate the pH of a solution in which [OH-] is M
Calculate the pH a solution in which the [OH-] is 2.0 x 10-9
*rules for sig figs and logs: #decimal places = # sig figs in original #

24. Calculations with pH
A sample of apple juice has a pH of Calculate [H+]
pH= -log[H+] =3.76
Thus : log [H+] = -3.76
Use antilog or INV
Ans: 1.7 x M

25. Calculations with pH
A solution formed by dissolving an antacid tablet has a pH of 9.18.
Calculate the [H+]
Calculate the [OH-]

26. Other “p” values
pOH = -log[OH-] pH + pOH = -log Kw = 14.00

27. Calculations with pH
Strong acids are acids which completely ionize, and there are relatively few.
HCl, HBr, HI,HClO3, HNO3 and HClO4
are STRONG monoprotic acids
H2SO4 is a STRONG diprotic acid.
Memorize the STRONG acids
[H+] concentrations are assumed to be = to the concentration of the solution because complete ionization

28. Calculations with pH
What is the pH of a 0.040M solution of HNO3?

29. Strength of Acids and Bases
The principles of equilibrium apply to weak acids and bases because they only partially dissociate into ions.
The dissociation of acetic acid into acetate ions and hydronium ions is shown as:
Chapter 8
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30. Strength of Acids and Bases
The Equilibrium Constant, Ka
Weak acids dissociate much less than 100% and have an equilibrium constant called an acid dissociation constant, Ka.
The strength of a weak acid can be determined from the Ka value. The larger the Ka value, the stronger the acid (the more protons dissociated).
Refer to the table distributed in class
Chapter 8
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31. Weak acids and bases
This table shows Ka values for substances acting as weak acids.
Chapter 8
© 2011 Pearson Education, Inc.

32. Strength of acids and bases
Using Ka to calculate pH
Using the value of Ka and the initial concentration of the solution, we can determine the [H+]
There are four steps to solving a Ka problem- RICE
(reaction, initial, change, equilibrium)

33. Ka Problems
Calculate the pH of a 0.20 M solution of HF. We need to find H+ concentration before we find pH Reaction: HF(aq)  H+(aq) + F- (aq)
Initial
0.20 X x
Change
-xM
+xM
Equilibrium
(0.20-x)M xM

34. Ka Problems Substitute values from table into equilibrium expression
Because Ka is very small, we can presume that x is very small compared to M, therefore we say 0.20M-x ~ 0.20
Ka = (x)(x) = 3.5 x 10-4
0.20
Solve x = 8.37 x 10-3
pH = - log [H+]
pH= -log 8.37 x 10-3
pH = 2.078

35. Ka problems The Ka of HCN is 4.9 x 10-10
Calculate the pH of a 0.20 M solution

36. Ka problems
Diprotic and polyprotic acids will have two or more Ka because each proton takes more work. The subsequent Ka values will be smaller.

37. The Meaning of pH (regents)
pH measures the H+ ion concentration in a solution. If…..
[H+] = [OH-] ______________________
[H+] > [OH-] ______________________
[H+] < [OH-] ______________________
An increasing pH means the H+ ion concentration is _________
A decreasing pH means the H+ ion concentration is _________
The pH scale is logarithmic
Each change of a single pH unit signifies a TENFOLD change in the concentration of H+ ions
A solution with a pH of 4 is 10x more acidic than a pH of ___
A solution with a pH of 7 is 100x more acidic than a pH of __

38. Neutralization Reactions
When an acid reacts with a base to produce salt and water
It occurs when there are the same amount of H+ ions as OH- ions
General formula: acid + base  water + salt

39. Writing Neutralization Reactions
Take the H+ from the acid and the OH- from the base and combine them to form H2O
Then combine the other ions to form the salt
Remember to look at oxidation numbers!!!!
Make sure the equation is balanced
When a solution is neutral, the moles of H+ = moles OH-
Ex: HCl + KOH  ___________________
Ex: H2SO4 + NaOH  ___________________

40. Acid-Base Titration
A laboratory technique used to achieve neutralization between an acid and a base
It is used to find the unknown molarity of an acid or a base by slowly adding measured volumes of an acid or base of known molarity until neutralization occurs
End point: The point when the indicator changes color
nHMAVA = nOHMBVB

41. Examples
How many milliliters of 4.00M NaOH are required to exactly neutralize 50.0mL of a 2.00M solution of HNO3?
If it takes 55mL of 0.1M NaOH solution to neutralize 450mL of an H2SO4 solution of unknown concentration, what’s the molarity of the acid?

42. Table M Examples You put bromcresol green
in a solution  it turns blue
Then you put bromthymol
blue in the same solution 
it turns yellow
What is the pH of the solution? _____________
Is the solution acidic or basic? _____________
A solution turns yellow with thymol blue and blue with bromthymol blue.
What is the pH of th\e solution? ____________
Is the solution acidic or basic? ____________