Body fluids, Acid, Base, Buffers, and Body Buffers

1.  Water ( Chemical formula : (H2O)
 A transparent fluid that forms world’s streams ,
lakes , oceans, rain.
 As a chemical compound, contains 2 Hydrogen
atoms and 1 Oxygen.
 Capillary Actor , Universal Solvent are the nick
names.

2.  A single water molecule is made up of 2
hydrogen atoms bonded with 1 oxygen atom
 The bond that forms water is a covalent bond

3.  Polar Molecule
 Capillary Action based molecule
 Surface Tension
 Universal Solvent
 Cohesion
 Adhesion
 Specific Heat

4.  A molecule that having electrically charged ends
 Water contains 2 Hydrogen +ve and 1 “O” –ve
ions
 The positive hydrogen ends of 1 molecule
attracted to the negative end of the oxygen of
another molecule – Polar Molecule

5.  Cohesion – the attractive force between water
molecules
 Water is Cohesion and its attracted by itself.
 3. Adhesion
 Adhesion –water molecules are attracted to
other substances
 Being water is Polar molecular, It having capacity
to attract with other molecules and surface

6.  Other name (or) nick name of Water
 Having specialty to dissolve lot of things
 SOLVENT = the thing doing the dissolving
 SOLUTE = the thing that dissolves away
 Why is water so good at dissolving things?
Because water is a polar molecule and is shaped
like a wedge, it is able to break up substances
into smaller pieces (dissolve)

7.  An action that having ability of a liquid to flow
in narrow space without the assistance of
external forces. Ex: Gravity
 Capillarity (or) Capillary motion are the Other
name.
 Cohesion + Adhesion

8.  The tightness across the surface of water that is
caused by the polar molecules pulling on one
another.
 Makes the surface act like a solid
SPECIFIC HEAT CAPACITY
Specific Heat = the amount of energy needed to
increase the temperature of something 1
degree C.
 Water has a really HIGH specific heat
 That means it takes a lot of energy for water to
increase its temperature.
 This is because of the STRONG ATTRACTION
between water molecules

9.  Acids produce H+ ions in aqueous solutions water
 HCl H+(aq) + Cl- (aq)
 • Taste sour • Corrode metals • React with bases
to form salts and water • Have a pH
Example of Acids
 HCl - Hydrochloric Acid Stomach acid
 C6H8O7 - Citric Acid Found in Oranges, Lemons
and other citrus fruit
 HNO3 - Nitric Acid used to cauterise warts and
in explosives, fertilisers and dyes

10.  Produce OHions in water water
 NaOH OH-(aq) + Na+ (aq)
 Taste bitter, chalky • Are electrolytes • Feel soapy,
slippery • React with acids to form salts and water •
Have a pH >7 • Turn Red Litmus paper Blue
Examples of Bases
 NaOH - Sodium Hydroxide Used to manufacture soaps
and detergents and as a drain cleaner
 KOH - Potassium Hydroxide Used to manufacture
liquid soaps and fertilizers and in alkaline batteries
 NH3 – Ammonia House hold cleaner
 Mg(OH)2 - Magnesium Hydroxide Used in Ant-acid
treatment such as Mylanta

11.  The strength of an acid (or base) is determined by the amount
of IONIZATION. (how easily it forms ions in water).
Strong = near complete ionisation
Weak = little ionisation
Weak Acids
Citric Acid
Vinegar (acetic acid)
Weak Bases
Baking soda (sodium hydrogen carbonate) • Sodium Carbonate
Strong Acids
HNO3
HCl
Strong Bases
NaOH
CaOH

12.  The pH scale is logarithmic and inversely
indicates the concentration of hydrogen ions in
the solution (a lower pH indicates a higher
concentration of hydrogen ions).
 The pH scale is a way of expressing the strength
of acids and bases.
 pH <7 –Acids
 pH = 7 - neutral
 pH >7-Base

13.  There are several ways to test pH
• Blue litmus paper (red = acid)
• Red litmus paper (blue = basic)
• pH paper (multi-colored)
• pH meter (7 is neutral, 7 base)
• Universal indicator (multi-colored)
• Indicators like phenolphthalein
• Natural indicators like red cabbage, radishes

14.  The pH is a measure of the concentration of
hydrogen ions in an aqueous solution. pKa
(acid dissociation constant) and pH are
related, but pKa is more specific in that it
helps you predict what a molecule will do at
a specific pH. Essentially, pKa tells you what
the pH needs to be in order for a chemical
species to donate or accept a proton.
 The relationship between pH and pKa is
described by the Henderson-Hasselbalch
equation.

15.  The pKa is the pH value at which a chemical
species will accept or donate a proton.
 The lower the pKa, the stronger the acid and the
greater the ability to donate a proton in aqueous
solution.
 The Henderson-Hasselbalch equation relates pKa
and pH. However, it is only an approximation and
should not be used for concentrated solutions or
for extremely low pH acids or high pH bases.

16.  If you know either pH or pKa, you can solve for the
other value using an approximation called the
Henderson-Hasselbalch equation:
 pH = pKa + log ([conjugate base]/[weak acid])
pH = pka+log ([A-]/[HA])
 pH is the sum of the pKa value and the log of the
concentration of the conjugate base divided by the
concentration of the weak acid.
 At half the equivalence point:
 pH = pKa
 It's worth noting sometimes this equation is written
for the Ka value rather than pKa, so you should know
the relationship:
 pKa = -logKa

17.  Only use strong acids or strong bases if the
pKa values fall between 5 and 9.
 Example pKa and pH Problem
 Find [H+] for a solution of 0.225 M NaNO2 and
1.0 M HNO2. The Ka value (from a table) of
HNO2 is 5.6 x 10-4.
 pKa = −log Ka = −log(7.4×10−4) = 3.14
 pH = pka + log ([A-]/[HA])
 pH = pKa + log([NO2
-]/[HNO2])
 pH = 3.14 + log(1/0.225)
 pH = 3.14 + 0.648 = 3.788
 [H+] = 10−pH = 10−3.788 = 1.6×10−4

18.  Buffers are compounds or mixtures of compounds
that by their presence in the solution resist
changes in the pH upon the addition of small
quantities of acid or alkali.”

19.  Sometimes it is necessary that a solution of a definite
pH be prepared and stored.
 The preservation of such a solution is even more
difficult than its preparation.
 If solution comes in contact with air, it will absorb
CO2 and becomes acidic.
 On the other hand, if solution is stored in a glass
bottle, alkaline impurities from the glass may alter
its pH.
 Due to these reasons, pharmaceutical solutions are
buffered as the buffer solutions are capable of
maintaining pH at some fairly constant value when
even small amounts of acid or base are added.

20.  Generally buffers are of two types;
 Acidic buffers
 Basic buffers
Acidic Buffers:
 An acidic buffer is a combination of weak acid
and its salt with a strong base. i.e. Weak acid &
salt with strong base (conjugate base).
 EXAMPLES: CH3COOH / CH3COONa
 H2CO3 / NaHCO3
 H3PO4 / NaH2PO4
 HCOOH / HCOONa

21.  A basic buffer is a combination of weak base and
its salt with a strong acid. i.e. Weak base & salt
with strong acid (conjugate acid).
 EXAMPLES:
 NH4OH / NH4Cl
 NH3 / NH4Cl
 NH3 / (NH4 )2CO3

22.  Besides the two general types of buffers (i.e.
acidic & basic), a third appears to exist.
This is buffer system composed of two salts:
 Monobasic potassium phosphate (KH2PO4 )
 Dibasic potassium phosphate (K2HPO4 ).

23.  The resistance of a buffer solution to a change in
pH is known as buffer action.
 Mechanism of Action of acidic buffers:
 Consider a buffer system of CH3COOH (Weak
electrolyte) and CH3COONa (Strong electrolyte).
 There will be a large concentration of Na+ ions,
CH3COONa – ions, and undissociated CH3COOH
molecules. When an acid is added
 If a strong acid (Hcl) is added in CH3COOH /
CH3COONa buffer, the changes that will occur
may be represented as:

24.  The hydrogen ions yielded by the Hcl are quickly
removed as unionized acetic acid, and the
hydrogen ion concentration is therefore only
slightly affected (because acetic acid produced
is very weak as compared to Hcl added)
 When a base is added If a strong base (NaoH) is
added in CH3COOH / CH3COONa buffer, the
changes that will occur may be represented as

25.  The hydroxyl ions yielded by the NaoH are
therefore removed as water. The supply of
hydrogen ions needed for this purpose being
constantly provided by the dissociation of acetic
acid.

26.  Consider a buffer system of NH4OH (Weak
electrolyte) and NH4Cl (Strong electrolyte).
 There will be a large concentration of NH4+ ions,
Cl – ions, and undissociated NH4OH molecules.
When an acid is added
 If a strong acid (HCl) is added in NH4OH / NH4Cl
buffer, the changes that will occur may be
represented as

27.  The hydrogen ions yielded by the Hcl are
therefore removed as water. The supply of OH
ions needed for this is constantly provided by the
ammonium hydroxide.
 When a base is added If a strong base (NaOH) is
added in NH4OH / NH4Cl buffer, the changes
that will occur may be represented as:

28.  The hydroxyl ions yielded by NaOH are therefore
quickly removed as unionized
ammoniumhydroxide and pH of solution is
slightly affected

29.  In KH2 PO4 / K2HPO4 buffer system, H2 PO4 -
serves as weak acid and HPO4 -2 serves as
conjugate base. When hydronium ions are
added, then
 HPO4 -2 + H3O+ → H2 PO4 - + H2O When
hydroxyl ions are added to this buffer, the
following reaction takes place;
 H2 PO4 - + OH- → HPO4 -2 + H2O

30. The body's chemical buffer system consists
of three individual buffers out of which
the carbonic acid bicarbonate buffer is
the most important.

31.  Cellular respiration produces carbon dioxide
as a waste product. This is immediately
converted to bicarbonate ion in the blood.
 On reaching the lungs it is again converted
to and released as carbon dioxide.
While in the blood , it neutralises acids
released due to other metabolic processes.
 In the stomach and deudenum it also
neutralises gastric acids and stabilises the
intra cellular pH of epithelial cells by the
secretions of bicarbonate ions into the
gastric mucosa.

32.  Phosphate buffer system operates in the internal
fluids of all cells. It consists of dihydrogen
phosphate ions as the hydrogen ion donor ( acid )
and hydrogen phosphate ion as the ion acceptor
( base ) .
 If additional hydroxide ions enter the cellular
fluid, they are neutralised by the dihydrogen
phosphate ion.
 If extra hydrogen ions enter the cellular fluid
then they are neutralised by the hydrogen
phosphate ion.

33.  Protein buffer system helps to maintain acidity
in and around the cells.
 Haemoglobin makes an excellent buffer by
binding to small amounts of acids in the blood,
before they can alter the pH of the blood.
 Other proteins containing amino acid histidine
are also good at buffering.
 The main purpose of all these buffers is to
maintain proper pH within the body system so
that all biochemical process can take place.