1
INTRODUCTION
An acid is a substance that can yield(donate) a
hydrogen ion(H+) or hydronium ion when
dissolved in water. Eg: HCl, CH3COOH.
A base is a substance that can yield hydroxyl
ions(OH-). Eg: NaOH.
pH: Negative logarithm of H+ concentration.
[ pH = -log H+ (mol/L) ]
pK: Negative logarithm of dissociation
(ionization) constant.
Introduction of Acids & Bases
They are everywhere..
In your food
In your house
EVEN IN YOU!!!!!
Objectives
•
•
•
•
Define acids and bases.
Know the uses of acids and bases.
Understand what is pH and pH scale.
Know the use of Henderson-Hasselbalch
equation.
• Know buffer system of the body.
What is an acid?
An acid is a substance that can yield(donate) a
hydrogen ion(H+) or hydronium ion (H3O+) when
dissolved in water. Example: HCl, CH3COOH.
It is sour in taste and turns blue litmus paper red
The sting of Ant contain
formic acid( HCOOH)
Uses of Acids
• Acetic Acid = Vinegar
• Citric Acid = lemons, limes, &
oranges. It is in many sour
candies such as lemonhead &
sour patch.
• Ascorbic acid = Vitamin C
which your body needs to
function.
• Sulfuric acid is used in the
production of fertilizers, steel,
paints, and plastics.
• Car batteries
What is a base?
A base is a substance that can yield hydroxyl ions(OH-).
Another word for base is alkali. Example: NaOH, KOH,
Mg(OH)2.
It is bitter in taste and turns red litmus paper to blue
Uses of Bases
• Bases give soaps, ammonia,
and many other cleaning
products some of their useful
properties.
• The OH- ions interact strongly
with certain substances, such
as dirt and grease.
• Chalk and oven cleaner are
examples of familiar products
that contain bases.
• Your blood is a basic solution.
pH
0
pH is generally defined as the negative logarithm of the
hydrogen ion activities (concentration) expressed over
14 orders of magnitude.
pH = -log10 [H+]
1
2
3 4 5
6
7 8 9 10 11 12 13 14
Acidic
Basic
Neutral
[H+]>[OH-]
[H+] = [OH-]
[OH-]>[H+]
pH
• pH is a measure of how acidic or
basic a solution is.
• The pH scale ranges from 0 to 14.
• Acidic solutions have pH values
below 7
• A solution with a pH of 0 is very
acidic.
• A solution with a pH of 7 is
neutral.
• Pure water has a pH of 7.
• Basic solutions have pH values
above 7.
pH Scale
Base: proton acceptor
Devised by Sorenson (1902)
 [H+] can range from 1M and
1 X 10-14M
 using a log scale simplifies
notation
 pH = -log [H+]
Neutral pH = 7.0
Acid: proton donor
pH p
Scale
A pH below the reference range (<7.34 ) is referred
to as acidosis, whereas a pH above the reference
range (>7.44) is referred to as alkalosis.
Dissociation constant - pK
• pK indicates the strength of acids and bases
and their ability to dissociate in water.
• Strong acids have pK values < 3, whereas
strong bases have pK > 9.
• pH = pK when the protonated and unpronated
(dissociated) forms are in equal concentration.
Henderson Hesselbach equation
• The Henderson Hesselbach equation expresses
acid-base relationships in a mathematical
formula:
cA
pH = pK + log------cHA
• WhereA- = proton acceptor(eg. HCO3-),
•
HA = proton donor or weak acid (eg.H2CO3),
•
pK = pH at which there is an equal
concentration of protonated and
unprotonated species.
SOURCES OF HYDROGEN IONS
Hydrogen ions[H+] are produced in the body as a result of
metabolism, particularly from the oxidation of sulphur
containing amino acids of proteins ingested as food.
The total amount of H+ produced each day in this way is of
the order of 60mmol.If all this was to be diluted in the
extracellular fluid, [H+] would be 100, 000 times more than
normal! This just does not happen.
because of effective buffer systems present in the body
and the normal ventilatory and excretory function of lungs
and kidneys.
Regulatory Mechanisms That Maintain pH
Buffer systems of the blood. React very rapidly
(less than a second)
Kidney buffer system. Concentration of
bicarbonate controlled by the kidneys. Reacts
slowly (minutes to hours)
Can reabsorb bicarbonate
Can synthesize bicarbonate
Respiratory buffer system. Concentration of
carbonic acid controlled by the lungs. Reacts
rapidly (seconds to minutes)
Buffers
• Buffers are solutions in which the pH remains
relatively constant, even when small amounts of
acid or base are added.
– Contain a weak acid (CH3COOH) and its
salt(CH3COONa); H2CO3 & NaHCO3
How buffers work
• Equilibrium between acid and base.
• Example: Acetate buffer
CH3COOH  CH3COO- + H+
• If more H+ is added to this solution, it simply
shifts the equilibrium to the left, absorbing H+,
so the [H+] remains unchanged.
• If H+ is removed (e.g. by adding OH-) then the
equilibrium shifts to the right, releasing H+ to
keep the pH constant.
Buffer systems of the body
Blood Buffer System
Kidney Buffer System
Respiratory Buffer System
Blood Buffer Systems
• Plasma Buffer Systems-bicarbonate buffer system
-phosphate buffer system
-protein buffer system
• RBC buffer systems-haemoglobin buffer system
-bicarbonate buffer system
Buffer Systems of body (contd)
Kidney buffer systems• Phosphate buffer systems- most important
• Ammonia buffer system
• Bicarbonate buffer system
Respiratory buffer system• Bicarbonate buffer system