Acids and Bases
Definitions
Arrhenius Definitions
• Acid – substance that produces hydrogen
ions (H+) when dissolved in water
• Base – substance that produces hydroxide
ions (OH-) in water
Properties
ACIDS
BASES
Produce H+ in water
Produce OH- in water
Corrosive
Corrosive
Sour if dilute enough to
drink/ingest
Irritating when dilute
Bitter if dilute enough to
drink/ingest
Slippery
pH < 7
pH > 7
Properties
• Acidic properties are due to the presence
of H+ ions.
– The more H+ the more acidic….the more
corrosive, sour…
• Basic properties are due to the presence
of OH- ions
– The more OH- ions the more basic the
solution…..the more corrosive, bitter…
Properties
• Acids and bases neutralize each other.
– the reaction of an acid with a base produces
water and a salt (an ionic compound)
• The resulting solution is neutral (pH =7)
– no longer an acid or a base.
HCl + NaOH  NaCl
+
H2O
Indicators
• Acid-base indicators – color depends upon the
acidity or alkalinity* of the solution
– Litmus paper
– Phenolphthalein
* alkaline is another word for basic
Indicators
• Indicators in nature
– Some hydrangea
– Red cabbage juice
Modern Definition of Acid and Base
Bronsted – Lowry Definitions:
• Acid – proton (H+) donor
• Base – proton (H+) acceptor
• Bronsted-Lowry definitions expand the
number of substances that are classified
as acid or base.
Acids in Water
• Bronsted and Lowry knew that each H+
released by an acid attaches to a water to
form the hydronium ion, H3O+
HCl + H2O  H3O+ + Cl-
Label the acid and base in this reaction.
Strong vs. Weak Acids
• Acids can be classified as either strong or
weak.
– Strong acids release all of their H+ in water,
weak acids release very few of their H+
Strong vs. Weak Acids
•
Strong acids – ionize 100% in water
– Every acid molecule releases its H+
– There are only 4 strong acids, 3 listed:
•
HCl – hydrochloric acid
•
HNO3 – nitric acid
•
One H of H2SO4 – sulfuric acid
Strong acids in water
• HCl – already shown
H 2O 
HNO3 +
H2SO4
+
H 2O 
Weak Acids
• Weak acids – ionize slightly in water
– Very few of the acid molecules release their
H+, most stay in their molecular form (unionized)
– Most acids are weak acids
Weak Acids
• Examples of weak acids
– Acetic acid (vinegar), lactic acid, citric acid,
ascorbic acid, acetylsalicylic acid….
• All acids except the strong acids presented are
weak acids.
• All carboxylic acids are weak acids – these are the
type of acid most commonly found in living
organisms
Weak acids in water
• The symbol written between reactants and
products is different for weak acids.
HCN
+ H2O
New Terms
• Conjugate acid – proton (H+) donor in the
reverse reaction
– The reverse reaction is the reaction that
reforms the reactants.
• Conjugate base - proton (H+) acceptor in
the reverse reaction
Weak acids in water
• Write the reaction that shows the weak
acid HC2H3O2 dissolving in water.
• Label the acid, base, conjugate acid and
conjugate base in this reaction.
Strong and Weak Bases
• Strong bases ionize fully in water.
– Have OH- in their chemical formula
– NaOH is a strong base
• Weak bases ionize slightly in water.
– Most nitrogen compounds are weak bases.
– NH3 is a weak base
Weak base in water
NH3
+
H 2O
– Complete the reaction.
– Label the acid, base, conjugate acid and
conjugate base.
New Term
• Amphoteric – substance that can behave
as an acid and a base
– Water is amphoteric.
pH Scale
• The acidity or alkalinity of dilute acid or
base solutions is commonly expressed as
a pH.
• Definition of pH:
pH = -log [H+]
[ ] means concentration
pH Scale
• Acids: pH <7
[H+] > [OH-]
– The more H+ the more acidic the solution and
the lower the pH.
• Neutral: pH = 7
[H+] = [OH-]
• Bases: pH > 7
[OH-] > [H+]
– The more OH- the more basic the solution
and higher the pH.
pH Scale
• The pH scale is based on the following
finding.
– In all aqueous solutions:
[H+] x [OH-] = 1.0 x 10-14
pH Scale
• In a neutral solution [H+] = [OH-] = 1 x 10-7
• pH = -log [H+]
• pH = -log 1 x 10-7 = _____
pH Scale
Consider an acidic solution with a pH of 2.
• [H+] = ___________
• [OH-] = ___________
pH Scale
Consider an acidic solution with a pH of 3.
• [H+] = ___________
• [OH-] = ___________
pH Scale - Acids
pH
[H+]
[OH-]
2
1 x 10-2
1 x 10-12
3
1 x 10-3
1 x 10-11
4
pH Scale -Bases
pH
9
10
11
[H+]
[OH-]
pH Scale
1 2 3 4 5 6 7 8 9 10 11 12 13 14
Sample questions!
• What is the pH of a solution that is 1000 x
more acidic than pH 6?
• What is the pH of a solution that is 10x
less basic than pH 10?
Buffered Solutions
• Buffers – resist a change in pH even when
acid or base is added
• Made from weak acids or weak bases.
– Strong acids/bases cannot be buffered
– Most solutions in the body are buffered.
Change of Topics!
• Collision theory of a chemical reaction
explains why reactions occur faster:
– When the concentration of the reactant is
increased
– At higher temperatures
– In the presence of catalysts
Collision Theory
• Collision theory of a chemical reaction:
– for a reaction to occur the reactants must
collide in proper orientation and with sufficient
force to meet the activation energy (Ea).
– Any collision not meeting both of these criteria
does not result in a chemical reaction.
• The rest of the lecture will be presented on the
board.