Introduction to Organic and
Biochemistry
(CHE 124)
Reading Assignment
General, Organic, and Biological Chemistry: An Integrated Approach
3rd. Ed. Ramond
Chapter 7
Acids, Bases, and Equilibrium
Gasses, Solutions, Colloids, and Suspensions
Work Problems
7. 8, 12, 24, 26, 29, 30, 32, 36, 40, 44, 52, 56, 60, 66
Acid / Base
• Acid
– Sour taste (never taste lab chemicals!)
– Dissolves metals
– Turns litmus pink
• Base
– Bitter taste (never taste lab chemicals!)
– Feel slippery (soapy)
– Turns litmus blue
• See common acids and bases Table 7.1 p.
224.
Acid / Definitions
• Arrhenius definition
– Acid – compound that produces H+ (protons) in aqueous
solution.
• HCl
→
H+ + Cl-
– Base – compound that produces OH- in aqueous
solution.
• NaOH → Na+ + OH-
• Bronsted-Lowery definition
– Acid – releases H+ (protons).
– Base - H+ (proton) acceptor
• HCN + H2O ⇌ CN- + H3O+
• Acid
Base Base Acid
arrows mean reversible
• Reversible - means products can be converted into products
and products can be converted into reactants.
Hydrogen and Related Species
• Proton (hydron)
H+
• Hydronium ion (interchangable with proton) H30+
• Hydrogen atom
• Hydro (hydrogen) group
H·
H
• Hydride
H: or
• Hydrogen gas or molecule H:H or
H¯
H2
• Hydroxide
OH¯
• Hydroxyl group
HO
OH
or
Acid and Conjugate Base
• Acid and conjugate base differ by presence
or absence of a proton.
Conjugate
Acid
Base
HCN + H2O ⇌ CN- + H3O+
Conjugate
Base
Acid
Amphoteric – compound that can act as an acid or
a base.
Equilibrium
• Consider the reversible reaction of
decomposition of dinitrogen tetroxide to
form nitrogen dioxide.
– See Fig. 7.2 p. 226
N2O4 (g) ⇌ 2 NO2(g)
colorless
brown
Eventually the color stops changing (getting browner).
This is equilibrium – the rate of the forward and
reverse reaction are equal. The concentration of
each species remains constant.
Note the double arrow.
Equilibrium Constant
• If the concentration of the reactant and
product of an equilibrium equation are
determined then the following equation is true.
• Keq = [NO2]2 = 4.6 x 10-3
[N2O2]
Keq = Products
Reactants
[ ] = molarity
Writing Equilibrium Equation Keq
• To write an equilibrium constant (Keq) equation.
– Before you start, BALANCE the EQUATION!
– ONLY SPECIES WHOSE CONCENTRATION CAN
CHANGE ARE INCLUDED.
– Do NOT include solvents or solids in the equation.
aA + bB ⇌ cC + dD
A,and B are reactants
C and D are products
a,b,c,and d are coefficients
Keq
=
[C]c [D]d
[A]a [B]b
What does Keq Tell us?
• Keq > 1 (larger number)
– [reactant] < [product]
– Favors products
• Keq < 1 (small number)
– [reactant] > [product]
– Favors reactants
• Keq = 1
– [reactant] = [product]
Ka
• Ka = acidity constant
– Special name of Keq for acid base reactions.
• pKa = -log Ka.
Le Chatelier’s Principle
• Le Chatelier’s Principle states that when a
reversible reaction is pushed out of
equilibrium, the reaction responds to
reestablish equilibrium.
– Vary [reactant] or [product] by adding (or
removing) reactants or products.
Example of Le Chatelier’s Principle
carbonic anhydrase
H2O(l) + CO2 (g) ⇌ H2CO3(aq)
water
carbon dioxide
carbonic acid
– Describe where / why this reaction occurs?
– Describe what happens if you increase [CO2]
• Reaction proceeds in the forward direction (to the right)
– Describe what happens if you decrease [CO2]
• In which direction does the reaction proceed.
– Describe what happens if you increase [H2CO3]
Catalysts
• Catalyst increase the rate of the reaction by
lowering the activation energy.
– Catalyst
• Do not alter the equilibrium
• Do not alter the Keq.
Water is Amphoteric
• Amphoteric a compound that can act as an
acid or a base.
HCl + H2O
Acid
⇌ Cl- + H3O+
Base
Base
Acid
NH3 + H2O ⇌ HN4+ + OHBase
Acid
Acid
Base
Water Can Ionize
H2O (l) + H2O (l)
↔
H3O+ (aq) +
hydronium ion
Acid
Base
Acid
OH(aq)
hydroxide ion
Base
Kw = [H3O+][OH-] = 1 X 10-14
Kw is water equilibrium constant.
pH
pH = -log [H3O+]
• Measure of [H3O+]
• scale is continuum from 0 - 14
– 7 is neutral;
• Neutral - neither acidic or basic
– 0 - 6.99 is acidic
– 7.01 - 14 is basic (alkaline)
• one pH unit change represents 10 fold
change in [H+]
• See Fig. 7.6 p 233
• See Table 7.3 p. 233
pH of Strong Acids
• Strong Acid – dissociates 100% in water.
HCl → H+ + ClHCl
HBr
HI
HClO4
HNO3
H2SO4
hydrochloric acid (muriatic acid)
hydrobromic acid
hydriodic acid
perchloric acid
nitric acid
sulfuric acid
• [H+] is equal to the [H+] of the acid
• Weak Acid – dissociates less than 100% in
water.
– All other acids
pH of Strong Bases
• Strong Base – dissociate 100% in water.
NaOH → Na+ + OHLiOH
NaOH
KOH
Ca(OH)2
Sr(OH)2
Ba(OH)2
Lithium hydroxide
Sodium hydroxide
Potassium hydroxide
Calcium hydroxide
strontium hydroxide
barium hydroxide
• pOH is dependent on the concentration of the strong
base
• Weak Base – dissociates less than 100% in
water.
Neutralization
• Neutralization reaction of an acid and base to
form water and a salt.
HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)
Acid
Base
Salt
Water
– If equal amounts of acid and base are added, the
pH will equal 7.
• Titration
– A technique used to determine the concentration
of an acid or base solutions.
• Uses Buret and and an indicator
• See p. 238 Fig. 7.9
pH effects the Concentration of the
Acid and Conjugate Base
• A few points to understand:
– When pH = pKa
• [acid] = [conjugate base]
– When pH < pKa
• [acid] > [conjugate base]
– When pH > pKa
• [acid] < [conjugate base]
• This alters the charge on many biological molecules
by changing them form the acid form to the base
form (carboxylate ion)
– Use the fatty acid example.
Buffer
• Buffer - substance that resists changes in pH thus
stabilizing its relative pH.
Buffers are often a solution containing a
weak acid and its conjugate base
– See example on next slide.
–
– Buffers work within 1 pH unit either side of
the pKa of the weak acid.
Buffer
• Carbonic acid is a weak acid. It dissociates in aqueous solution
to form hydronium ion and bicarbonate
H2CO3
↔
H3O+
carbonic acid hydronium ion
+
HCO3bicarbonate
Buffering Blood
• pH of blood = 7.35 – 7.45
• Blood carries many acids which can alter it’s
pH.
– Fatty acids, lactic acid, phosphoric acid, carbonic
acid.
• Body uses two approaches to control pH (p.245
Fig 7.12)
– Use of Buffers (see next slide)
• Carbonic Acid / Bicarbonate buffer system
– Reduce [H3O+] (see following slides)
• Action of lungs
• Filtering by kidneys
Use of Buffers
• Carbonic acid is a weak acid that buffers blood. It dissociates
in aqueous solution to form hydronium ion (acid) and
bicarbonate.
H2CO3
↔
H 3O +
carbonic acid hydronium ion
+
OH↔
H2O
+
HCO3bicarbonate
+
H+
Acidosis
• Acidosis - low blood pH
• Leads to light headedness, coma, death.
– Respiratory Acidosis
• Characteristics
– Low blood pH; high blood PCO2; normal or high (if compensating)
blood HCO3-
• Causes
– Diseases / conditions that limit carbon dioxide exchange by lungs
such as ppneumonia, emphysema, cystic fibrosis, shallow breathing
or holding your breath.
– Metabolic Acidosis
• Characteristics
– low blood pH; normal or low (if lungs are compensating) blood PCO2;
low blood HCO3-
• Causes
– Presence of ketone bodies (acetone, acetoacetic acid, beta
hydroxybuyteric acid) due to starvation or poorly controlled
diabetes.
• See Table 7.7 p. 244 and Figure 7.12 p. 245
Reducing [H3O+]
• Lungs remove excess acid through increase in
respiration rate
– As the blood becomes more acidic, the respiratory center
of the brain signals for faster breathing.
– With faster breathing, CO2 is exhaled at a faster rate thus
reducing the partial pressure of carbon dioxide (PCO2).
This reduces the [carbonic acid] thus reducing the
[hydronium] producing an increase in pH.
– This happens when you exercise.
• Lungs remove excess base by reducing rate of
respiration.
– Breathing becomes slower and more shallow. PCO2
increase leads to increase [carbonic acid] and thus
[hydronium] and a drop in pH.
Reducing [H3O+] Cont’
• Kidneys remove excess acid by releasing
bicarbonate into the blood.
– The increase in [bicarbonate] shifts the equilibrium
toward carbonic acid. This reduces the
[hydronium].
Correcting Acidosis
CO2 + H20
↑ respiration rate
(breathing becomes
more rapid)
causes ↓ pCO2
(lungs remove
carbon dioxide from
blood and release
it into atmosphere)
shifts equation.
Think about exercise
↔ H2CO3 ↔ H3O+ +
HCO3-
kidneys
release
H3O+ in
urine.
kidneys
generate /
release
bicarbonate
shifts
equation
shifts
equation
Alkalosis
• Alkalosis - high blood pH.
• Leads to headaches, nervousness, cramps, and convulsions and
death.
– Respiratory alkalosis
• Characteristics
– high blood pH; low blood PCO2; normal or lower (if kidneys are
compensating) blood HCO3-
• Causes
– Occurs when CO2 is exhaled from the body more quickly than it is
produced by cells.
– Hyperventilation brought on by anxiety, CNS damage, aspirin
poisoning, fever, etc
– Metabolic alkalosis
• Characteristics
– high blood pH; normal or high (if lungs are compensating) blood
PCO2; high blood HCO3-
• Causes
– Excessive use of antacids and constipation.
• See Table 7.7 p. 244 and Figure 7.12 p 245
Correcting Alkalosis
CO2 + H20 ↔ H2CO3
↓ respiration rate
(breathing slows)
causes ↑ pCO2
shifts equation.
↔ H3O+ +
kidneys
generate
and
release
acid
into
blood
shifts
equation
HCO3kidneys
remove
HCO3from
blood
and
release it
into urine.
shifts
equation