UNIT 14 NOTES - Chapter 19 Acids, Bases, and Redox
Properties and Acids and Bases
Properties of Acids
1. sour to taste
2. have hydrogen in front of formula: HCl, H2SO4
3. pH range is from 0-6.9
4. dissolved in water (aqueous) –contains more hydronium(H3O+) ions than
hydroxide(OH-) ions
5. turns litmus paper red
6. feel “rough”
7. conduct electricity in solution
8. react with metals to produce hydrogen gas
9. corrosive
10. proton donors (H+)
11. neutralize bases
Properties of Bases
1. bitter to taste
2. the formula usually ends in OH (hydroxide): NaOH –an exception is NH3
(ammonia)
3. pH range is from 7.1-14
4. dissolved in water (aqueous) –contains more hydroxide ions than hydronium ions
5. turns litmus paper blue
6. feel “slippery”
7. conduct electricity in solution
8. a strong base can be just as harmful as a strong acid
9. neutralize acids
10. proton acceptors (H+)
11. react with oils and greases (used as cleaners)
(1)
Common Acids
Hydrochloric acid
Sulphuric acid
Stomach juice
Lemons
Vinegar
Apples
Oranges
Grapes
Sour milk
White bread
Fresh milk
pH
0.1
0.3
1-3
2.3
2.9
3.1
3.5
4
4.4
5.5
6.5
Common Bases
Human saliva
Distilled water
Blood plasma
Eggs
Seawater
Borax
Milk of magnesia
Ammonia water
Limewater
Caustic soda
pH
6-8
7
7.4
7.8
7.9
9.2
10.5
11.6
12.4
14
Acids can be identified by their reaction with some metals to produce hydrogen gas
For example: Aluminum reacts with aqueous solutions of hydrochloric acid to
produce hydrogen gas and aluminum chloride.
Theories
Arrhenius Theory (H or OH will appear in the formula)
-Acids produce ____H+ Ions________ in aqueous solution. EX:
-Bases produce _____OH- Ions________ in aqueous solution. EX:
Bronsted-Lowry Theory
-Acids are __________Proton Donors_____________________.
-Bases are _____________Proton Acceptors____________.
An H+ ion is transferred from an acid to a water molecule to produce the
__Hydronium Ion ____.
HX + H2O → H30+ + XWhen an acid donates a proton, what remains is called the Conjugate Base.
When a base receives a proton, the species that forms is called the Conjugate Acid.
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Conjugate Acid-Base Pairs.
A conjugate acid is the species produced when a base accepts a hydrogen ion from an
acid. The conjugate base is the species that results when an acid donates a hydrogen ion
to a base.
Example:
HC2H3O2
acid
↔
+ H2O
base
H3O+
conjugate
acid
+
C2H3O2−
conjugate
base
Identify the conjugate acid-base pairs in the following reactions.
NH4+(aq) + OH-(aq) ↔ NH3(aq) + H2O
HBr(aq) + H2O(l) ↔ H3O+ (aq) + Br-(aq)
CO32-(aq) + H2O(l) ↔ HCO3-(aq) + OH-(aq)
HSO4-(aq) + H2O(l) ↔ H3O+(aq) + SO42-(aq)
Amphoteric -
Strength of Acids and Bases
Strengths of Acids
 Strong acids are acids that ionize completely
 Because they ionize completely they are good conductors of electricity
 Rules to determine strong acids:
1) Binary acids: HCl, HBr, HI are strong
2)
Ternary acids: General rule: if the # of oxygen atoms exceeds the
number of hydrogen atoms by two or more, the acid is strong: HClO4,
H2SO4, HNO3

HCl  H+ + Cl This is a strong acid and the reaction goes to completion and is shown with only
one arrow. That means that if you have 1M HCl you yield 1M H+
and 1M Cl-.

Weak acids partially ionize/dissociate in water
HF  H+ + FThis is a weak acid and the reaction does not go to completion and is
shown with a double arrow.
(3)
Strengths of Bases
 Strong bases dissociate entirely into metal ions and hydroxide ions
 Rules to determine strong bases:
In general, Groups I and II metallic hydroxides are considered strong. All
others are weak.

NaOH  Na+ + OH This is a strong base and the reaction goes to completion and is shown with
only one arrow. That means that if you have 1M NaOH you yield 1M Na+
and 1 M OH-.

Weak bases partially ionize/dissociate in water
NH3 + H2O  NH4+ + OHThis is a weak base and the reaction does not go to completion and is
shown with a double arrow.
*Strong acids and strong bases ionize completely in water making them good conductors of
electricity. Weak acids and bases do not ionize completely in water and therefore do not
conduct electricity well. Weak acids and bases dissociate less than 5%. Acids and bases that
dissociate at greater than 5% are considered strong.
What is pH?
Autoionization of Water
A.
Water self-ionizes by the following rxn:
2 H2O <----> H3O+ + OH-
B.
Kw=[H3O+][OH-] where Kw=ion product constant for water
Kw=1.008x10-14 @ 25°C
C.
Pure Water Equilibrium
H2O  H+ + OHRemember, for pure water:
if: pH = 7, [H+] = 1 x 10-7M
and: [H+] = [OH-]
then: [OH-] = 1 x 10-7M
D.
pOH and pH are related by the statement:
pH + pOH = 14
Example Problem:
At 298 K, the H+ ion concentration of an aqueous solution is 1.0 x 10-5 M. What is the
OH- ion concentration in the solution? Is the solution acidic, basic or neutral?
(4)
pH
A.
pH= negative logarithm of the hydronium ion (hydrogen ion) concentration in
solution
B.
pH is a logarithmic scale (based on powers of ten)
1. numbers are more manageable
2. scale from 0 to 14
C.
pH = - log [H+] or pH = -log[H3O+] (same thing)
D.
Practice finding the pH of solutions with the following [H+]:
[H+] = 1 x 10-4 M
pH =
[H+] = 3.5 x 10-3 M
pH=
[H+] = 5.5 x 10-14 M
pH =
Using pH to find [H+]
A. To find the [H+] you will need to use the anti-log function or the 10x button on
your calculator.
B. Formula will be:
[H+] = inverse log-pH
C. STEPS
1. Push the second log (antilog key 10x)
2. enter the -pH value (change the sign of the pH)
3. close the parentheses
D. Practice with these pH and pOH values.
pH = 6.7
[H+] =
pOH = 2.5
[H+] =
(5)
pOH = 9.9
[H+] =
Putting it all together
A. By knowing any of the [H+], pH, or pOH values all other values can be
calculated.
B. Calculate the unknown values.
C. Determine whether the substance is acidic or basic.
[H+] = 3.46 x 10-4M
pH =
pOH =
pH = 8.90
pOH =
[H+] =
[H+] = 8.32 x 10-5 M
pOH =
pH=
pOH = 3.71
pH =
[H+] =
Measuring pH
Use pH paper (like litmus paper) or a pH meter
Naming Acids
Acids are a group of compounds that are given special treatment in naming. You will see
that acids are defined in several different ways when you look at the chemistry of acids in
more detail in Chapter 18. For now it is sufficient to know that acids are compounds that
give off hydrogen ions when dissolved in water.
The formulas of acids are of a general form HX, where X is a monatomic or polyatomic
anion. We have previously named the compound HCl, hydrogen chloride. When the
compound HCl is dissolved in water, it is named as an acid. Other compounds, such as
HNO3, exist only in water solution. They are always named as acids.
Consider the acid HX as dissolving in water. The acid can be named using three rules
that focus on the ending of the anion of the acid (See Table).
When the anion (X) ends in –ide, the acid name begins with the prefix hydro-. The stem
of the anion has the suffix –ic and it is followed by the word acid. Thus HCl (x =
chloride), dissolved in water, is named hydrochloric acid. H2S (x = sulfide) is
hydrosulfuric acid.
Table - Naming Acids
Anion Ending
-ide
-ite
-ate
Example
ClChloride
SO32sulfite
NO31nitrate
Acid Name
Hydro-(stem)-ic
acid
(stem)-ous acid
Example
Hydrochloric acid
(stem)-ic acid
Nitric acid
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Sulfurous acid
When the anion ends in –ite, the acid name is the stem of the anion with the suffix –
ous, followed by the word acid. H2SO3 (x = sulfite) is sulfurous acid.
If the anion ends in –ate, the acid name is the stem of the anion with the suffix –ic,
followed by the word acid. Thus HNO3 (x = nitrate) is nitric acid.
Example
Name these compounds as acids. a. HClO
Solution
a. hypochlorous acid (rule 2)
b. HCN
c. H3PO4
b. hydrocynaic acid (rule 1)
c. phosphoric acid (rule)
Formulas of acids are most easily written by using the preceding three rules in a reverse
fashion. For example, what is the formula of chloric acid? As rule 3 shows, chloric acid
(-ic ending) must be a combination of hydrogen ion (H+) and chlorate ion (ClO3-). The
formula of chloric acid is HClO3. Hydrobromic acid (hydro- prefix and –ic suffix),
according to rule 1, must be a combination of hydrogen ion and bromide ion (Br-). The
formula of hydrobromic acid is HBr. Hydrogen ion and phosphite ion (PO33-) must be
the components of phosphorus acid (rule 2). The formula of phosphorous acid is H3PO3.
Problems
1. Name these compounds as acids.
a. HF
b. HC2H3O2
c. H2SO4
d. HNO2
2. Write formulas for the following acids.
a. chromic acid
b. hydroiodic acid
c. chlorous acid
d. perchloric acid
Reactions
Neutralization Reactions (acid-base reactions)
A.
General Rxn. : acid + base → salt + water
Example: HCl
+ KOH →
KCl + H2O
B.
water produced from the union of a H+ from acid and OH- from base
C.
Salt: compound formed from the positive ion of an aqueous base (other than
hydrogen) and the negative ion of an aqueous acid (other than
hydroxide) {NOT ALWAYS NaCl}
D.
Equal amounts of H+ and OH- will neutralize completely
(7)
Precipitation Reactions
A.
Some reactions (double replacement) that occur in aqueous solutions produce
precipitates. Be sure to use the solubility chart to check the solubility of the
products.
Example: Ba(NO3)2(aq)
B.
+ Na2CO3(aq) → BaCO3(s) + 2NaNO3(aq)
A double replacement reaction cannot produce two aqueous products. If this
happens it is a “No Reaction”.
Example:
2NaCl(aq) + Ca(NO3)2(aq) →
CaCl2(aq) + 2NaNO3(aq)
(NO REACTION)
Oxidation-Reduction Reactions (Redox reactions)
A.
A reaction in which electrons are transferred from one atom to another.
B.
Oxidation is defined as the loss of electrons and Reduction is defined as the
gain of electrons. Remember: LEO says GER
C.
Assign oxidation numbers to elements in compounds.
D.
Single elements are assigned an oxidation of “0”.
Example: 2KBr + Cl2 →
2KCl
+ Br2
2K+1 + 2Br-1 + 2Cl0 → 2K+1 + 2Cl-1 + 2Br0
(This is a net ionic equation.)
2Br-1 → Br20 (bromine loses electrons and is oxidized)
Cl20 →
Cl-1 (chlorine gains electrons and is reduced)
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