1.  1. General Definitions:
 Acid: any substance which when dissolved
into the water produces hydrogen ions [H+].
 Base: any substance which when dissolved
into the water produces hydroxide ions [OH-
].
 2. Water dissociation: H2O(l) → H+
(aq) + OH-
(aq)
 equilibrium constant, KW = [H+][OH-] / [H2O]
 Value for Kw = [H+][OH-] = 1.0 x 10-14

2.  Note: The reverse reaction, H+
(aq) + OH-
(aq) → H2O(l) is not equal to 1 x 10-14
 [H+] for pure water = 1 x 10-7
[OH-] for pure water = 1 x 10-7
 3. Definitions of acidic, basic, and neutral
solutions based on [H+]
 acidic: if [H+] is greater than 1 x 10-7 M
 basic: if [H+] is less than1 x 10-7 M
 neutral: if [H+] if equal to 1 x 10-7 M

3.  Example 1: What is the [H+] of a sample of
lake water with [OH-] of 4.0 x 10-9 M?
Is the lake acidic, basic, or neutral?
 Solution: [H+] = 1 x 10-14 / 4 x 10-9 = 2.5 x
10-6 M
 Therefore the lake is slightly acidic

4.  pH is a measurement of the H+ concentration
in a liquid.
 relationship between [H+] and pH
 pH = -log10[H+]

5.  Acids
 1. Strong Acids:
 A substance is strong acid when dissolved
into the water or other solvent completely
dissociate into proton( H+) and an anion.
 example: HN03 dissociates completely in
water to form H+ and N03
1-.
 The reaction is
 HNO3(aq) → H+
(aq) + N03
1-
(aq)

6.  example: HN03 dissociates completely in
water to form H+ and N03
1-.
 The reaction is
 HNO3(aq) → H+
(aq) + N03
1-
(aq)
 A 0.01 M solution of nitric acid contains 0.01
M of H+ and 0.01 M N03
- ions and almost no
HN03 molecules. The pH of the solution
would be 2.0.

7.  1.HCl
2.H2SO4
3.HNO3
4.HClO4
5.HBr
6.HI

8.  Note: when a strong acid dissociates only one
H+ ion is removed. H2S04 dissociates giving
H+ and HS04
- ions( first ionization).
 H2SO4 → H+ + HSO4
1-
 A 0.01 M solution of sulfuric acid would
contain 0.01 M H+ and 0.01 M HSO4
1-
(bisulfate or hydrogen sulfate ion).
 Because HS04
- is still having proton, it is also
an acid and can dissociate into H+ and SO4
2-(
Second ionization)

9.  2. Weak acids:
 a weak acid only partially dissociates in water
or other solvents to give H+ and the anion
 for example, HF dissociates in water to give
H+ and F-. It is a weak acid. with a
dissociation equation that is :

10.  There are only 6 strong acids, the remainder
of the acids therefore are considered weak
acids.

11. Strong Bases
 Dissociate 100% into the cation and OH-
(hydroxide ion).
 example: NaOH(aq) → Na+
(aq) + OH-
(aq)
 a. 0.010 M NaOH solution will contain 0.010
M OH- ions (as well as 0.010 M Na+ ions) and
have a pH of 12.
 The strong bases are the hydroxides of
Groups I and II.

12.  Note: the hydroxides of Group II metals
produce 2 mol of OH- ions for every mole of
base that dissociates. These hydroxides are
not very soluble, but what amount that does
dissolve completely dissociates into ions.
 exampIe: Ba(OH)2(aq) → Ba2+
(aq) + 2OH-
(aq)
 a. 0.000100 M Ba(OH)2 solution will be
0.000200 M in OH- ions (as well as 0.00100
M in Ba2+ ions) and will have a pH of 10.3.


13. Weak Bases
 What compounds are considered to be weak
bases?
 Most weak bases are anions of weak acids.
 Weak bases do not furnish OH- ions by
dissociation. They react with water to furnish
the OH- ions.
 Note that like weak acids, this reaction is
shown to be at equilibrium, unlike the
dissociation of a strong base which is shown
to go to completion.

14.  When a weak base reacts with water the OH-
comes from the water and the remaining H+
attaches itsef to the weak base, giving a weak
acid as one of the products. You may think of
it as a two-step reaction similar to the
hydrolysis of water by cations to give acid
solutions.

15.  examples:
 NH3(aq) + H2O(aq) → NH4
+
(aq) + OH-(aq)
 methylamine: CH3NH2(aq) + H20(l) → CH3NH3
+
(aq) +
OH-
(aq)
 acetate ion: C2H3O2
-
(aq) + H2O(aq) → HC2H302(aq) +
OH-
(aq)
 General reaction:
weak base(aq) + H2O(aq) → weak acid(aq) + OH-
(aq)

16.  Since the reaction does not go to completion
relatively few OH- ions are formed.

17. Acid-Base Properties of Salt Solutions
 A salt : an ionic compound made of a cation
and an anion, other than hydroxide or the
product besides water of a neutralization
reaction.
 determining acidity or basicity of a salt
solution:

18. 1.split the salt into cation and anion
2. add OH- to the cation
a. if you obtain a strong base. the cation is
neutral
b. if you get a weak base, the cation is acidic
3. Add H+ to the anion
a.if you obtain a strong acid, the anion is neutral
b. if you obtain a weak acid. the anion is basic
 Salt solutions are neutral if both ions are neutral
 ( BPH WEEKEND)

19.  Salt solutions are acidic if one ion is neutral
and the other is acidic
 Salt solutions are basic if one of the ions is
basic and the other is neutral.
 The acidity or basicity of a salt made of one
acidic ion and one basic ion cannot be
determined without further information.

20. Examples: determine if the following solutions
are acidic, basic, or neutral
 KC2H3O2
 NaHPO4
 Cu(NO3)2
 LiHS
 KClO4
 NH4Cl

21. Acid-Base Reactions
 Strong acid + strong base: HCl + NaOH → NaCl
+ H2O
 net ionic reaction: H+ + OH- → H2O
 Strong acid + weak base:
•example: write the net ionic equation for the
reaction between hydrochloric acid, HCl, and
aqueous ammonia, NH3. What is the pH of the
resulting solution?

22.  when solution gets neutralized?
 During this process, indicators will be used.
 Indicators are chemical compounds that turn
different colors when they're in solutions with
different pH's.
 Litmus, for example, is red in acid solutions and
blue in basic solutions.
 Phenolphthalein is clear in acid solutions and
pink in basic solutions.
 .The basic equation for titration or
neutralization is:
 M1V1 = M2V2


23. •M1 stands for the molarity of the acid
•V1 stands for the volume of the acid you use
•M2 stands for the molarity of the base
•V2 stands for the volume of the base you use
Example: If it takes 55 mL of 0.1 M NaOH solution
to neutralize 450 mL of a HCl solution of
unknown concentration, what's the molarity of
the acid?
 M1, in our equation, stands for the molarity of
the acid.

24.  Since that's what we're trying to find, we'll call that
X.
 V1 stands for the volume of the acid we use. Since
HCl is an acid, the volume of acid is 450 mL
 M2 stands for the molarity of the base. Since NaOH
is a base, the molarity was stated in the problem to
be 0.1 M
 V2 stands for the volume of the base. The problem
says that we used 55 mL of base, so that's M2.

25.  Now, all we need to do is plug it into the
equation:
 (X)(450 mL) = (0.1 M)(55 mL)
X = 0.12 M

26. Buffers solutions
 solutions that don't change pH very much when
you add acid or base solutions to it.
 For example, if you were to add a little bit of HCl
to a glass of water, the pH might change from 7
to 3.
 If you had the same amount of buffer solution,
the pH might change from 7 to 6.8.

27.  Buffers are formed :
• a weak acid + its conjugate .
Example1: acetic acid+ sodium acetate.
→ acidic buffer
Example 2: a weak base+ its conjugate acid.
→ basic buffer

28. Weak acid-strong base titrations
Example:
Titration curve for the titration of vinegar with
NaOH. pH at end point- approximately 8.5 ;
species present- H2O and NaC2H3O2 and
appropriate indicator-phenolphthalein

30.  Note: no matter what type of titration you do,
at the equivalence (end) point the number of
moles of H+ is equivalent to the number of
moles of OH-.
 This applies whether you have weak or strong
acids and/or bases.
 Problems:
l. Citric acid (C6H807) contains a mole of
ionizable H+/mole of citric acid. A sample
containing citric acid has a mass of 1.286 g.

31.  The sample is dissolved in 100.0 mL of water.
The solution is titrated with 0.0150 M of
NaOH. If 14.93 mL of the base are required to
neutralize the acid. then what is the mass
percent of citric acid in the sample?

32. Models of acids
• Arrhenius Model
 The basis for the model is the action in water
 The Arrhenius definition:
 acids are compounds that give off H+ ions in
water
 bases are compounds that give off OH- ions
in water.

33.  As a result, all acids have hydrogen atoms on
them that are ready to go jumping off in water.
 Most common acids have the letter H in the
beginning of the formula, with the exception of
acetic acid.
 Arrhenius and Bronsted-Lowry definitions are
for most purposes identical. When you see the
formula of a base, it's got "OH" in it.
 The one exception to this is ammonia,
NH3. (NH3 combines with water to form
NH4OH, which is really the thing that's basic in
ammonia.

34. Strong base + weak acid:
•example: write the net ionic equation for the
reaction between citric acid (H3C6H507) and
sodium hydroxide. What is the pH of the
resulting solution?
Titrations
 Titration : method used in order to determine
the concentration of an acidic solution( or basic
solution) by adding amount of base( or acid)
that you know the concentration.

35.  You have an acidic solution and you want to
figure out the molarity. You can't do that
directly, because you can't count acid
molecules. You can, however, make a basic
solution with a concentration that you already
know. If you keep adding base to the acid,
eventually all of the acid molecules will be
neutralized and the solution will turn from an
acid to a base.

36.  If you know how many base molecules you
added to the solution before the solution gets
neutralized (and you will, because you'll add
the solution drop-by-drop), you can figure
out how much acid was in the solution in the
first place.

37.  Indicators: chemical compounds that turn
different colors when they're in solutions with
different pH's.
 Litmus, for example, is red in acid solutions
and blue in basic solutions.
 Phenolphthalein is clear in acid solutions and
pink in basic solutions.

38.  The basic equation for titration or
neutralization is:
 M1V1 = M2V2
 M1 stands for the molarity of the acid
 V1 stands for the volume of the acid you use
 M2 stands for the molarity of the base
 V2 stands for the volume of the base you use

39.  Example: If it takes 55 mL of 0.1 M NaOH
solution to neutralize 450 mL of a HCl
solution of unknown concentration, what's
the molarity of the acid?
 M1, in our equation, stands for the molarity of
the acid. Since that's what we're trying to
find, we'll call that X.
 V1 stands for the volume of the acid we
use. Since HCl is an acid, the volume of acid
is 450 mL

40.  M2 stands for the molarity of the base. Since
NaOH is a base, the molarity was stated in the
problem to be 0.1 M
 V2 stands for the volume of the base. The
problem says that we used 55 mL of base, so
that's M2.
 Now, all we need to do is plug it into the
equation:
 (X)(450 mL) = (0.1 M)(55 mL)
X = 0.12 M

41.  Buffers solutions
 Buffers are solutions that don't change pH
very much when you add acid or base
solutions to it.
 For example, if you were to add a little bit of
HCl to a glass of water, the pH might change
from 7 to 3.
 If you had the same amount of buffer
solution, the pH might change from 7 to 6.8.


42.  Buffers are formed when you have a weak
acid and its conjugate base present in the
same place.
 If you wanted to make an acidic buffer, you'd
place some acetic acid into a container with
some sodium acetate.
 If you want a basic buffer, just put a weak
base into a container with it's conjugate
acid. Our blood is a buffered solution.

43.  If it wasn't, our pH would be go way down
every time we had a soda and way up
whenever we took some Tums.

44. Weak acid-strong base titrations

45.  l. Arrhenius Model
 The basis for the model is the action in water
 The Arrhenius definition of acids says that
they're compounds that give off H+ ions in
water and that bases are compounds that
give off OH- ions in water.
 These definitions are the same. Basically, if
you've got something that can give off H+ in
water, it's an acid. As a result, all acids have
hydrogen atoms on them that are ready to go
jumping off in water.

46.  As a result, all acids have hydrogen atoms on
them that are ready to go jumping off in
water.
 Most common acids have the letter H in the
beginning of the formula, with the exception
of acetic acid. Bases, on the other hand, are
compounds that give off OH- in water.

47.  (The two definitions of a base are for our
purposes identical, as OH- combine with H+
to form water -- the Arrhenius and Bronsted-
Lowry definitions are for most purposes
identical).
 When you see the formula of a base, it's got
"OH" in it. The one exception to this is
ammonia, NH3.

48.  (NH3 combines with water to form NH4OH,
which is really the thing that's basic in
ammonia. So our definition is sort of true).
 Here are a couple of charts which show the
most common acids and bases. Some are
strong and some are weak, as indicated.

49. Formula Name Strong?
HCl hydrochloric acid yes
HBr hydrobromic acid yes
HI hydroiodic acid yes
HF hydrofluoric acid no
HNO3 nitric acid yes
H2SO4 sulfuric acid yes
H3PO4 phosphoric acid no
CH3COOH acetic acid no

50.  2. Bronsted-Lowry Model
 The basis for the model is proton transfer
 According to Bronsted-Lowry; acids are
compounds that give off H+ ions when you
stick them in water. This definition also says
that bases are compounds that can accept H+
ions when you stick them in water.

51.  Simply, acids have H+ in them and bases
have OH- in them.
 The conjugate base of an acid is whatever is
formed when the acid loses its H+ or the
base becomes the conjugate acid after it
accepts the proton because it can now donate
it back.

52.  The acid becomes the conjugate base after it
donates the proton because it can now accept
it back.
 As a general rule of thumb, the conjugate
bases of strong acids are weak. For example,
Cl- is the conjugate base of hydrochloric acid

53.  3. Lewis Model
 The basis for model is the electron pair
transfer
 The hydrogen requirement of Arrhenius and
Brønsted–Lowry was removed by the Lewis
definition of acid–base reactions, devised by
Gilbert N. Lewis in 1923, in the same year as
Brønsted–Lowry, but it was not elaborated by
him until 1938. Instead of defining acid–base
reactions in terms of protons or other bonded

54.  substances, the Lewis definition defines a
base (referred to as a Lewis base) to be a
compound that can donate an electron pair,
and an acid (a Lewis acid) to be a compound
that can receive this electron pair.
 In this system, an acid does not exchange
atoms with a base, but combines with it.

55.  Lewis definition can be applied to reactions
that do not fall under other definitions of
acid–base reactions. For example, a silver
cation behaves as an acid with respect to
ammonia, which behaves as a base, in the
following reaction:
 Ag+ + 2 :NH3 → [H3N:Ag:NH3]+
 The result of this reaction is the formation of
an ammonia–silver adduct.

56. Formula Name Strong?
NaOH sodium hydroxide yes
LiOH lithium hydroxide yes
KOH potassium hydroxide yes
Mg(OH)2 magnesium hydroxide no
Ca(OH)2 calcium hydroxide no
NH3 (NH4OH)
ammonia
(ammonium hydroxide)
no