Acids, Bases and Buffers
In Earlier,
Acid may be defined as any substance which has a sour taste and its aqueous
solution turns blue litmus to red colour.
A base may
be defined as any substance which has a bitter taste and its aqueous solution
turns red litmus to blue colour.
But these
could not be explained the behavior of all acids or bases.
In
chemistry, acids and bases have
been defined differently by three sets of theories.
Theories of
acid and base
- Arrhenius
theory
Base is a
substance which dissociates to give hydroxide ions in water
Eg. Sodium hydroxide
in water produces hydroxide ion
Limitations
1. The
definition of acid or bases are only in terms of aqueous solutions and not in
terms of substance.
2. This
theory is not able to explain acidic or basic in non aqueous solvents. For Ex
Ammonium nitrate in liquid ammonia acts as an acid, but it does not give H+
ions.
3. This
theory is not able to explain for the basic substances which does not contain
OH- ions, For Ex, Ammonia is a basic.
4. It
cannot be explain acidic character of salts like AlCl3 in aqueous solutions.
5. The
neutralization of acid and base in absence of solvent could not be explained.
2. Bronsted
and lowry theory (Proton concept)
According to this theory, Acids are substances which is capable to donate
protons H+ to any other substances. They are also called as
proton donors.
For Ex
Bases are substances
which can accept protons H+ from any other substances. They are also called as
proton acceptors.
For Ex
Conjugate acid-base pair
Let as consider a reaction
Acid1 Base2 Acid1 Base2
In this reaction HCl donates proton to water,
therefore it is an acid. On the otherhand, water accept proton from HCl, and
therefore it is a base.
Acid Base
Here HCl donates
a proton and ammonia accepts this proton and forms ammonium chloride. So HCl is
an acid and ammonia is a base. The acid-base pairs, the members of which can be
formed each other mutually by gain or loss of protons are called conjugate
acid-base pairs.
Water is having
dual character because it can accept or donate protons.
Limitations
1. This
theory is not able to explain about the acids which do not contain proton.
2. In
large no. of acid-base reactions, proton transfer is not taking place.
3. Lewis
theory (Electron concept)
According to
this theory, acids are chemical substances which accept a lone pair of
electrons and are called electron acceptors. Bases are substances which donate
a lone pair of electrons in solution and are called electron donors. So the
process of neutralization is simply the formation of a co ordinate bond between
acid and base.
In the above
reaction boron trifluoride accepts the lone pair of electron donated by
ammonia. So BF3 is an acid and NH3is a base. The behave
as acid only when a base is available to accept proton or donate electron and
similarly bases
Limitations
1. According
to this theory, the strength of Lewis acids and bases is depend upon the type
of reaction, it is not possible to arrange them in any order of their relative
strength.
2. As
Lewis acid - base reaction involves electrons, they are expected to be fast
reactions, but many reactions are slow.
Relative
strength of Acids and Bases
The relative strength of an acid or base is based on
the efficiency of donating or accepting protons. These properties of a
substances will be affected by the environment, an acid can donate proton easily
in proton accepting medium.
With
respect to strength there are two classes, strong and weak. Strong acids and
bases are dissociate completely in aqueous media, But weak acids and bases do
not dissociate completely. The relative strength of an acid or base is
determined by dissociation constant K.
The dissociation constant K can
be calculated by using the following formula
K = [Concentration of products] or [Concentration of ionized] .
[Concentration
of reactants] [Concentration
of unionized]
For example
HCl is a strong acid, which
dissociates completely in water
HCl
+ H2O ----> H+ + Cl-
The dissociation constant Ka = [H+] [Cl-]
[HA]
The concentration of water is not
consider, because of large quantity was used.
Similarly, ammonia is a base
The dissociation constant or
strength of acid or base K is usually expressed in log. pKa is the negative
log of the equilibrium constant
pKa = - log Ka the
negative log of the equilibrium constant for acids
pKb = - log Kb the
negative log of the equilibrium constant for bases
pKw =
- log Kw the negative log
of the equilibrium constant for water
· Strong acids have a large Ka, indicating
that there are more products than reactants, larger Ka,
stronger acid, larger pKa, weaker acid,
· larger Kb,
stronger base, larger pKb, weaker base
Examples,
pKa value for
weak acids
The pH scale
We use the pH scale to describe how many hydrogen
ions are dissolved in a solution. The pH
comes from the negative log of the hydrogen ion concentration.
pH = - log [H+]
The term, p, means we have taken the negative log of
something.
Similarly
pOH = - log [OH-] the negative log of the hydroxide ion
concentration
Let us consider
the ionization of water
H2O ----------à
H+ + OH-
K =
[H+][OH-]
[H2O]
But
the concentration of water is constant, because, in large quantity, only one molecule
is dissociated.
K
[H2O] = [H+][OH-]
Kw = [H+][OH-] where Kw is ionic product of water and its
value is 1 x 10-14
Take the log of both sides. Remember when you have stuff multiplied
together, when you take the log, you add them.
so
log Kw = log [H+] + log [OH-]
Now
multiply by -1 through the entire equation to get:
-log Kw = -log [H+] - log [OH-] or pKw
= pH + pOH
-log
10-14 = -log [H+]
- log [OH-]
14
= pH + pOH
So for water pH = 7, ie neutral pH.
The pH scale is from 0 to 14. The pH value is less
than 7 for acid and more than 7 for base. How strong an acid is labeled as pH. The lower the pH, the stronger the acid.
Buffers
A buffer solution is any solution that maintains an
approximately constant pH in small additions of acid and base. The buffer
solution consists of mixture of weak acids or bases with its salts.
Types of buffer solutions
1)
Acidic buffer: It consists of a weak acid and its conjugate base or salt.
It buffers on the acidic side of neutral.
Example,
Acetic acid and sodium acetate
2)
Basic buffer: consists of a weak base and its conjugate acid or salt. It
buffers on the basic side of neutral.
Example,
Ammonium hydroxide and Ammonium chloride
Properties of
buffer solutions
Ø The pH of buffer
solution remains constant.
Ø The pH of buffer
solution does not change on dilution
Ø The pH does not
change after addition of small quantity of acid or base
Ø The pH of buffer
solution does not change on keeping for long time.
Buffer actions
The resistance to change
in pH possessed by buffer is called buffer action
Acid
buffer action:
Let us consider
Acidic buffer which is prepared by mixing acetic acid and sodium acetate in water.
Acetic acid is a weak acid and ionized to acetate ion and hydrogen
ion to some extent only.
Sodium acetate is a strong
electrolyte, so fully dissociated to acetate ion and sodium ion
Therefore the buffer mixture
contains less amount of hydrogen ion and more amount of acetate
ion.
If a little amount of strong
acid is added to the buffer mixture the hydrogen ion from strong
acid, will be attracted by acetate ion, and feebly ignitable acetic
acid is formed, the acidity does not change much.
It is possible to calculate the pH of acidic buffer by
Henderson-Hasselbalch equation
Let
us consider the dissociation on an acid
Ka = [H+]
[A-]
[HA]
when
expressed in logarithmic form we get:
log Ka = log [H+] +
log [A-]
[HA]
multiplying
both sides by -1:
-log Ka =
-log [H+] - log [A-]
[HA]
Consider pKa = - log Ka. pH = -log [H+]
pKa = pH - log [A-]
[HA]
multiplying
by -1 through the entire equation to get rid of the - signs we get:
-pKa = -pH + log [A-]
[HA]
Rearranging
the above equation
pH = pKa + log [A-]
[HA]
pH = pKa + log [Conjugate base]
[Acid]
Or
Base buffer
action:
Similarly
Let us consider basic buffer which is prepared by mixing ammonia and ammonium chloride in water.
Base
Buffer Example:
NH3 (aq) +
H2O
(l) ------à NH4+ (aq) +
OH-(aq)
When strong acid
is added, NH3 accepts protons from incoming acid to make NH4+. When strong base
is added, NH4+ donates a proton to form NH3 and H2O. pH remains the same.
It
is possible to calculate the pH of acidic buffer by Henderson-Hasselbalch
equation
Let us consider
the dissociation on an acid
Kb = [B+]
[OH-]
[BOH]
[OH-] = kb [BOH]
[B+]
when
expressed in logarithmic form we get:
log [OH-] = log Kb + log [BOH]
[B+]
multiplying
both sides by -1:
-log [OH-] = -log Kb + log [B+]
[BOH]
pOH = pKb + log [B+]
[BOH]
Hence
pH + pOH = 14
So
pH = 14 – pOH
pH = 14 – ( pKb - log [B+]
[BOH]
pH =14- ( pKb + log [Conjugate acid]
[Base]
Buffer capacity(b)
It is the effectiveness of buffer in quantity basis. It is defined as the
amount of strong acid or strong base that must be added to the buffer to
produce unit change of pH.
Hence b =d[B] / dpH.
Since, the addition of base
increase the pH and addition of acid decrease the pH. The value of b
depends on the nature of buffer and the pH, which is determined by the relative
concentrations of the acid and its conjugate base. The buffer functions are
best around pKa value.
pH = pKa, when [HA] = [A-] and
there are equal amount of acid and its conjugate base.
For Example
In a buffer, the more A- and
HA molecules available, after the addition of a strong acid or base will have
on the pH of a system. Consider the addition of a strong acid such as HCl.
Initially, the HCl donates its proton to the weak base (A-) through the reaction
A- + HCl → HA + Cl-
This changes the pH by lowering the ratio
[A-]/[HA], but as long as there is still a lot of A- present,
the change in pH will be small. But if we keep adding HCl, the weak base A- will
eventually run out. Once the A- is gone, any additional HCl will
donate its proton to water HCl + H2O → H3O+ + Cl-
This will dramatically increase the
concentration [H+] and so the pH drops.
Example
In
Vivo biologic buffer systems
•Blood
Primary buffers
– Plasma
NaHCO3--H2CO3,
NaHPO4--NaH2PO4, protein
Secondary
buffers :Erythrocytes;
hemoglobin-oxyhemoglobin,
2HPO4-KH2PO4
•Lacriminalfluid
-pH: 7.4
(range 7 –8 or slightly higher)
•Urine
pH: 6.0 (range 4.5 –7.8)
-below
normal…hydrogen ions are excreted by the kidney.
-above pH
7.4…hydrogen ions are retained by action of the kidney.
Buffers in Pharmaceutical systems
In pharmaceutical
formulations, lot of buffers are used because
v Some
compounds are soluble in particular pH, For ex, inorganic salts of Ferric ion,
phosphates, borates are soluble in acidic media.
v
The pH of
blood is maintained around 7.4 by carbonic acid and sodium carbonate buffer
system and hemoglobin buffer system.
v
It is used in
the assay of enzyme activity because the enzyme activity is maximum at a
particular pH.
v
Used in the estimation
of metallic salts by complexometric titration with EDTA, as the metal
complex is more stable at a particular pH.
v
Certain
pharmaceutical preparations are stabilized by suitable buffers.
Eg. Neutral adrenaline eye drops stabilized by
borate buffer
v Penicillin preparations are buffered with sodium citrate, calcium carbonate and aluminum
hydroxide
v The
colour of the natural dyes are maintained only on particular pH, For ex, Red
colour of cherry and rasbery syrups has been maintained at acidic pH.
v Some
liquid preparations was stable only on particular pH range, For ex, Ascarbic
acid and penicillin are stable in acidic pH. Adrenaline injection is most stable
in pH 2.5 – 5.
v Optimum
pH conditions for activity of medicinal compounds are maintained. For ex, the
germicidal activity of sodium hypochlorite is increased at lowering pH.
v Buffers
of known pH are used as a standard in analytical lab.
Official buffers
The buffer solution recommended by pharmacopoeia are called official
buffers.
Examples,
1.
Hydrochloric
acid buffer
– 0.2M KCl and 0.2M HCl (pH = 5.8 to
8.0)
2.
Phthalate buffer
-
0.2M Pot. Hydrogen phthalate and 0.2M NaOH (pH = 4.2 to 5.8)
3.
Phosphate buffer
- 0.2M
pot. dihydrogen phosphate and 0.2M NaOH (pH = 5.8 to 8.0)
4.
Alkaline borate
buffer- 0.2M
Boric acid, 0.2M KCl and 0.2M NaOH (pH = 8 to 10)
5.
Acetate buffer -
Ammonium acetate or Sodium acetate and acetic acid (pH
= 2.8 to 5.0)
Isotonic solution
v It is a solution which has colligative properties similar to those of
body fluids. Isotonic is biological compatibility.
v An isotonic solution refers to two solutions having the same osmotic
pressure across a semi permeable membrane. This state allows for the free
movement of water across the membrane without changing the concentration of
solutes on either side.
v They have equal concentrations of solutes and water.
v Example- Sodium chloride solution 0.9% w/v which has a Freezing point
depression (△tf)= - 0.52℃
Hypertonic solution
v In biology, a hypertonic solution is
one with a higher concentration of solutes outside the cell than inside the
cell.
v
When a cell is
immersed into a hypertonic solution, the tendency is for water to
flow out of the cell in order to balance the concentration of the solutes.
v
It has a Freezing point depression -0.52℃ < tf < 0 ℃
v
It produce swelling and bursting of the cells
(ex. Hemolysis)
Hypotonic solution
v A hypotonic solution is any solution that has a
lower osmotic pressure than another solution.
v In the biological fields, this generally refers to a solution that
has less solute and more water than another solution
v It has a Freezing point depression
tf<-0.52℃
v It produce Shrinkage of the living cells
Some Isotonic Buffer Solutions
Preparation of Isotonic Solutions
Isotonic solution is
prepared by the drug solution which have equal concentration of 0.9%w/v of sodium
chloride.
Freezing Point Depression Method
Amount of NaCl needed in
100mL of solution = [0.9 x (0.52-tf’)] / 0.52
Ex. 100ml of a drug
solution , How much NaCl=? Fp=-0.18o C
Amount of NaCl needed =
[0.9 x (0.52- 0.18)]/0.52=0.59g
Sodium Chloride Equivalent Method
This method is based on
the calculation of Sodium Chloride Equivalent Values (E) and prepare the isotonic solution equivalent to 0.9% w/v
concentration.
E value can be calculated by the formula E = 17
(Liso / MW)
For Ex , 1g drug = ? g sodium chloride
Ex. E value? MW=340g/mole Liso =3.4
E
value=17(3.4/340)=0.17g
Measurement of tonicity
v Hemolytic
method - Red blood cells are applied
in the tonic solution. It liberates oxyhemoglobinin direct proportion to the
number of cells hemolyzed.
v Determine colligative properties like Freezing
point lowering Tf, Osmotic pressure etc. Based on the values when compare with
the value of isotonic solution, we can determine the tonicity.
For Isotonic solution Tf=
0.52 ºC (Freezing point lowering)
Calculating Tonicity Using Liso values•
Tonicity can be
calculated by the Van’tHoff expression
Method of adjusting tonicity and pH
Class I…add Sod.
Chloride to lower the freezing point of soln. to -0.52°
1. Cryoscopic method
2. Sodium chloride equivalent method
Class II…add Water
to form an isotonic soln.
1. White-Vincent method
2. Sprowlsmethod
Cryoscopic method
Amount of NaCl needed in
100mL of solution = [0.9 x (0.52-tf’)] / 0.52
Sodium Chloride Equivalent Method
This method is based on
the calculation of Sodium Chloride Equivalent Values (E) and prepare the isotonic solution equivalent to 0.9% w/v
concentration.
E value can be calculated by the formula E = 17
(Liso / MW)
White-Vincent Method
It is based on the
following formula
Wi= weight in grams of
the ith solute in the formulation
Ei= sodium chloride
equivalent of the ith solute in the formula
v = volume of sodium
chloride solution (0.9%) that contains 1g of NaCl (this volume is 111.1 mL)
Example
Sprowls’ Method
The Sprowls’ value is the
volume of an isotonic solution that can be prepared by the addition of enough
water to 0.3 g drug
“0.3g”represents the
amount of drug in a fluid ounce of 1% solution(1 fluid ounce≒0.0294L)
No comments:
Post a Comment