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ACIDS AND BASES
Chapters 19 and 9.4
Acids and Bases
• Properties of ACIDS
Taste Sour/Tart
Stings and burns the skin
Reacts with bases
Turns blue litmus paper red
Reacts with metals to form H2
gas
Neutralizes Bases
Donates H+
Conduct electricity.
• Can be strong or weak electrolytes in aqueous solution
• Properties of BASES
Taste Bitter
Feels slippery on skin
Reacts with acids
Turns red litmus blue
Doesn’t react with metals
Neutralizes Acids
Accepts H+
Two important classification of compounds - Acids and Bases
Chapters 19 and 9.4
Acids Affect Indicators
Blue litmus paper turns red in
contact with an acid.
Chapters 19 and 9.4
Acids React with Active Metals
Acids react with active metals to
form salts and hydrogen gas:
HCl(aq) + Mg(s) → MgCl2(aq) + H2(g)
Chapters 19 and 9.4
Sulfuric Acid – H2SO4
Highest volume production
of any chemical in the U.S.
Used in the production of
paper
Used in production of
fertilizers
Used in petroleum refining
Chapters 19 and 9.4
Nitric Acid – HNO3
• Used in the production of
fertilizers
• Used in the production of
explosives
• Nitric acid is a volatile acid –
its reactive components
evaporate easily
• Stains proteins (including
skin!)
Chapters 19 and 9.4
Hydrochloric Acid - HCl
• Used in the “pickling” of
steel
• Used to purify magnesium
from sea water
• Part of gastric juice, it aids
in the digestion of proteins
• Sold commercially as
“Muriatic acid”
Chapters 19 and 9.4
Phosphoric Acid – H3PO4
o A flavoring agent in sodas
o Used in the manufacture
of detergents
o Used in the manufacture
of fertilizers
o Not a common laboratory
reagent
Chapters 19 and 9.4
Acetic Acid – HC2H3O2
Used in the manufacture of
plastics
Used in making
pharmaceuticals
Acetic acid is the acid present in
household vinegar
Chapters 19 and 9.4
Bases Affect Indicators
Red litmus paper
turns blue in contact
with a base.Phenolphthalein
turns purple in a
base.
Chapters 19 and 9.4
Examples of Bases
Sodium hydroxide (lye), NaOH
Potassium hydroxide, KOH
Magnesium hydroxide, Mg(OH)2
Calcium hydroxide (lime),
Ca(OH)2
Chapters 19 and 9.4
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
Acid-Base Theories
Arrhenius
(1883)
Brønsted–Lowry (1923)
Lewis (1923-38)
HCl NaOH
NH3
BF3
H2O
H2SO4
AlI3
CO32-
• An acid is a substance that dissociates in water to
produce hydrogen ions (H+)
• Example: Hydrochloric Acid (HCl)
• A base is a substance that dissociates in water to
produce hydroxide ions (OH-)
• Example: Sodium Hydroxide (NaOH)
Arrhenius Definition of A & B
Chapters 19 and 9.4
What happens when placed in
water???
Chapters 19 and 9.4
Svante Arrhenius (1859-1927)
Chapters 19 and 9.4
Brønsted-Lowry Definition
• An acid is any substance that can donate H+
ions.
• A base is any substance that can accept H+
ions.
• Expansion on Arrhenius definition of A&B
Defines A&B as not necessarily in water solution
Does not contain OH-; covers bases such as ammonia NH
3
Chapters 19 and 9.4
Johannes Bronsted Thomas Lowry
(1879-1947) (1874-1936)
Brønsted-Lowry Definition
• H+
is really only just a proton (no electrons or
neutrons), so definition is often in terms of protons
Brønsted-Lowry Base is a proton acceptor
Brønsted-Lowry Acid is a proton donor• Monoprotic Acids can only donate 1 H
+: HCl
• Diprotic Acids can donate 2 H+: H
2SO
4
• Triprotic Acids can donate 3 H+: H
3PO
4
• Polyprotic acids = diprotic and triprotic acids
Chapters 19 and 9.4
Chapters 19 and 9.4
Chapters 19 and 9.4
• Acid-base reactions with water proceed in both
directions:
Example: NH3
(g) + H2O (l) NH
4
+(aq) + OH
-(aq)
• Acid loses H+
= conjugate base
Example: H2O (acid) loses its H
+, turning it into OH
-
(conjugate base)
• Base gains H+
= conjugate acid
Example: NH3
(base) gains an H+, turning it into NH
4
+
(conjugate acid)
Conjugate Acid-Base Pairs
Chapters 19 and 9.4
Acids and bases come in pairs
• A “conjugate base” is the remainder of the
original acid, after it donates it’s hydrogen
ion
• A “conjugate acid” is the particle formed
when the original base gains a hydrogen ion
• Indicators are weak acids or bases that have
a different color from their original acid and
base
Chapters 19 and 9.4
The Hydronium Ion
• Water can pick up a H+
ion to form a hydronium ion:
H+
+ H2O H
3O
+H
3O
+= hydronium ion
• With acids, water is a Brønsted-Lowry base (accepts
protons
Example: HCl (g) + H2O(l) H
3O
+(aq) + Cl
-(aq)
• With bases, water is a Brønsted-Lowry acid (donates
protons)
Example: NH3
(g) + H2O (l) NH
4
+(aq) + OH
-(aq)
• Compound that can act as either a proton donor or
acceptor = amphoteric
Lewis Acids and Bases
• Gilbert Lewis focused on the donation or
acceptance of a pair of electrons during a
reaction
• Lewis Acid - electron pair acceptor
• Lewis Base - electron pair donor
• Most general of all 3 definitions; acids don’t even need hydrogen!
Gilbert Lewis (1875-1946)
Chapters 19 and 9.4
Lewis Acids and Bases
• Lewis acid = accepts a pair of electrons during a
reaction
• Lewis base = donates a pair of electrons during a
reaction
• Covers acids and bases not covered by Brønsted-
Lowry definition
Type Acid Base
Arrhenius H+
producer OH-producer
Brønsted-Lowry H+
(proton) donor H+
acceptor
Lewis Electron-pair
acceptor
Electron-pair donor
Table 19.4, Pg. 592, Text
Chapters 19 and 9.4
The pH Scale
• pH is based on the concentration of the
hydronium ion in a solution
• Concentrations range from 100
M (strong) to 10-14
M (weak) of [H3O
+]
• pH ranges from 0 to 14
If [H30
+] concentration = 10
-2 M, pH = 2
If [H30
+] concentration = 10
-10 M, pH = 10
Water is 10-7
M, pH = 7
• pH of 0-6 = Acidic
• pH of 7 = Neutral
• pH of 8-14 = Basic
pH
…..defined as the negative base-10 logarithm of the
hydronium ion concentration.
pH = −log [H3O+]
Problem: Calculate pH when [H3O+] = 2.3 x 10-3 M
Problem: Calculate [H3O+] when pH = 2.3 ?
10x
log
antilog10^ -2.3
10x
logbase 10 log log - 2.3 * 10^- 3
For pure H2O: [1.0 10−7 ] = 7.0
= 2.64
= 5.0 * 10-3
pH and pOH Calculations
H+ OH-
pH pOH
[OH-] = 1 x 10-14
[H+]
[H+] = 1 x 10-14
[OH-]
pOH = 14 - pH
pH = 14 - pOH
pO
H =
-lo
g[O
H- ]
pH
= -
log[H
+]
[OH
- ] =
10
-pO
H
[H+]
= 1
0-p
H
Chapters 19 and 9.4
pH
These are
the pH
values for
several
common
substances.
pH + pOH = 14
pH and pOH equilibrium
in pure Water
−log [H3O+] + −log [OH−] = −log Kw
• In pure water,
[H3O+] [OH−] = Kw
• Because in pure water [H3O+] = [OH−],
Kw = [1.0 10−7 ] [1.0 10−7 ] = 1.0 10−14
pH + pOH = pKw
7 + 7 = 14
[1.0 10−7 ] [1.0 10−7 ] = 1.0 10−14
pH and pOH equilibrium
in Water to which
Acids & Bases are Added
Kw = [1.0 10−8 ] [1.0 10−6 ] = 1.0 10−14
pH + pOH = pKw
8 + 6 = 14
H2O(l) + H2O(l) H3O+(aq) + OH−(aq)
OH-Add base
[H3O+] [OH−]H2O
Kw = [1.0 10−7 ] [1.0 10−7 ] = 1.0 10−14
Chapters 19 and 9.4
Measuring pH
• Why measure pH?
Solutions we use –• swimming pools
• soil conditions for plants
• medical diagnosis
• soaps and shampoos, etc.
• Sometimes we can use indicators, other times we might need a pH meter
Chapters 19 and 9.4
Acid-Base Strength
• An acid or a base is considered strong if they
completely dissociate into ions (H+
and OH-) in
water
• Strong Acids
HCl and H2SO
4
• Strong Bases
Hydroxides, e.g. NaOH
• Conjugate acid-base pairs have an inverse relationship
(works for both acids and bases)
The stronger the acid, the weaker the conjugate base
The weaker the acid, the stronger the conjugate base
Strong Acids and Bases to
know
Chapters 19 and 9.4
Measuring pH with wide-range paper
1. Moisten indicator
strip with a few drops
of solution, by using a
stirring rod.
2.Compare the color
to the chart on the vial
– read the pH value.
Chapters 19 and 9.4
Acid-Base Indicators
• Although useful, there are limitations to indicators:
usually given for a certain temperature (25 oC), thus may change at different temperatures
what if the solution already has color, like paint?
the ability of the human eye to distinguish colors is limited
Chapters 19 and 9.4
Some of the
many pH
Indicators
and their
ranges
Chapters 19 and 9.4
Red Cabbage Juice as an indicator
• Red cabbage juice
mixed with
baking soda (left)
and with vinegar
(right).
On the top, a drop
of unmixed juice.
Chapters 19 and 9.4
Acid-Base Indicators
• A pH meter may give more definitive results
some are large, others portable
works by measuring the voltage between two electrodes; typically accurate to within 0.01 pH unit of the true pH
needs to be calibrated
Chapters 19 and 9.4
Acids Neutralize Bases
HCl + NaOH → NaCl + H2O
-Neutralization reactions
ALWAYS produce a salt and
water.
-Of course, it takes the right
proportion of acid and base
to produce a neutral salt
Acid Base Salt
Strong Strong Neutral
Strong Weak Acidic
Weak Strong Basic
Weak Weak Neutral,
basic, or
acidic
Acid-Base Properties of Salts
• Neutralization reaction =
reaction of an acid and a
base
• Acid + base react – form a
salt and water
• Type of salt depends on
reactants
Chapters 19 and 9.4
Chapters 19 and 9.4
Chapters 19 and 9.4
Buffers
• Buffers are solutions in which the pH
remains relatively constant, even when
small amounts of acid or base are added
made from a pair of chemicals: a weak
acid and one of it’s salts; or a weak
base and one of it’s salts
Chapters 19 and 9.4
Chapters 19 and 9.4
Acid rain
Chapters 19 and 9.4
Causes of emissions
Chapters 19 and 9.4
pH readings nationwide
Chapters 19 and 9.4
Acid rain effects limestone and Marble
Chapters 19 and 9.4
Effects of Acid Rain on Marble(calcium carbonate)
George Washington:
BEFORE
George Washington:
AFTER
Chapters 19 and 9.4
Naming Acids and Bases
• Naming Acids: Three Rules
1. Name of anion ends in –ide
Acid name begins with hydro-
Stem of anion has suffix –ic
2. Name of anion ends in –ite
Stem of anion has suffix –ous
3. Name of anion ends in –ate
Stem of anion has suffix –ic
All three end with the word “acid”
Naming Bases
Named just like ionic compounds – cation + anion
Chapters 19 and 9.4
Acid – Base Humor