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Imran Nur Manik, profile picture
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Buffer capacity MANIK

The document provides an extensive overview of buffer solutions, detailing their composition, mechanisms, and pH stabilization processes. It introduces the Henderson-Hasselbalch equation, explains buffer capacity, and discusses various buffer systems in biological contexts, especially the bicarbonate buffer in blood. Additionally, it covers osmotic pressure and tonicity, describing the implications of isotonic, hypertonic, and hypotonic solutions in pharmacology.

Health & Medicine◦
An unbuffered solution or a buffered solution acid added base added acid added base added Md. Imran Nur Manik
A buffer solution is a solution which resists changes in pH when a small amount of acid or base is added. The resistive action is the result of equilibrium between the weak acid (HA) and its conjugate base (A-). HA(aq) + H2O(l) → H3O+ (aq) + A- (aq) Typically a mixture of a weak acid and a salt of its conjugate base or weak base and a salt of its conjugate acid.Md. Imran Nur Manik
Two types :  ACIDIC BUFFERS – Solution of a mixture of a weak acid and a salt of this weak acid with a strong base. e.g. CH3COOH + CH3COONa ( weak acid ) (Salt )  BASIC BUFFERS – Solution of a mixture of a and a e.g + ( Weak base) ( Salt) Md.ImranNurManik
HOW BUFFERS WORK Equilibrium between acid and base. Example: ACETATE BUFFER  CH3COOH⇌CH3COO ⎯ +H+  CH3COONa⇌CH3COO⎯+Na+  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 Md.ImranNurManik
Equilibrium between acid and base. Example: Basic Buffer NH4OH⇌NH4 ++ OH⎯ NH4Cl ⇌NH4 ++ Cl⎯  If more OH⎯ is added to this solution, it simply shifts the equilibrium to the left, absorbing OH⎯ , so the [OH⎯] remains unchanged.  If OH⎯ is removed (e.g. by adding H+) then the equilibrium shifts to the right, releasing OH⎯ to keep the pH constant HOW BUFFERS WORK (CONT..) Md.ImranNurManik
HANDERSON HASSELBALCH EQUATION  Lawrence Joseph Henderson wrote an equation, in 1908, describing the use of carbonic acid as a buffer solution.  Karl Albert Hasselbalch later re-expressed that formula in logarithmic terms, resulting in the Henderson–Hasselbalch equation. Md.ImranNurManik
Ka = [H+] [A-] [HA] Take the -log on both sides The Henderson-Hasselbalch Equation derivation -log Ka = -log [H+] -log [A-] [HA] pH = pKa + log [A-] [HA] = pKa + log [Salt] [Acid] HA H+ + A- pKa = pH-log [A-] [HA] Apply p(x) = -log(x) And finally solve for pH… Md.ImranNurManik
Problem: Find the pH of a buffer solution containing 0.20 mole per litre CH3COONa and 0.15 mole per litre CH3COOH. Ka for acetic acid is 1.8 10–5. Problem: Calculate the concentration of acetic acid to be added to a 0.1 M solution of sodium acetate to give a buffer of pH 5 (pKa of acetic acid=4.66). Problem: The Ka of propionic acid is 1.34×10–5. What is the pH of a solution containing 0.5M propionic acid, C2H5COOH, and 0.5 sodium propionate, C2H5COONa. What happens to the pH of this solution when volume is doubled by the addition of water? Problem: A buffer solution contains 0.015 mole of ammonium hydroxide and 0.025 mole of ammonium chloride. Calculate the pH value of the solution. Dissociation constant of NH4OH at the room temperature is 1.80×10–5 Problem: Estimate the pH at 25°C containing 0.10 M sodium acetate and 0.03 M acetic acid pKa for CH3COOH = 4.57.
Buffer capacity is a measure of the efficiency of a buffer, in resisting changes in pH. The buffer capacity is defined as It is also known as buffer efficiency, buffer index, and buffer value. Conventionally, the buffer capacity (β) is expressed as A buffer solution can resist a small amount of change of pH on adding acid or alkali to the solution. Buffer capacities ranging from 0.01-0.1 are usually adequate for most pharmaceutical solutions.
In 1922, Van Slyke first introduced an approximate equation to determine the buffer capacity by the following equation: In which β = Buffer capacity, delta Δ=a finite change, and ΔB =the small increment in gram equivalents (gEq)/litre of strong base added to the buffer solution to produce a pH change of ΔpH. According to equation, the buffer capacity of a solution has a value of 1 when the addition of 1 g Eq of strong base (or acid) to 1 litre of the buffer solution results in a change of 1 pH unit. The higher the buffer capacity the less the buffer solution changes its pH. A more exact equation for buffer capacity: The buffer capacity calculated from above equation is only approximate. It gives the average buffer capacity over the increment of base added. Koppel and Spiro and Van Slyke developed a more exact equation, Where, C = the total buffer concentration (i.e. the sum of the molar concentrations of acid and salt).
Body fluids contain buffering agents and buffer systems that maintain pH at or near pH=7.4. Important endogenous (natural) buffer systems include carbonic acid/sodium bicarbonate and sodium phosphate in the plasma and haemoglobin, and potassium phosphate in the cells. An in vivo value of pH < 6.9 or pH > 7.8 can be life threatening. Pharmaceutical solutions generally have a low buffer capacity in order to prevent overwhelming the body’s own buffer systems and significantly changing the pH of the body fluids. Buffer concentrations of between 0.05 and 0.5 M and buffer capacities between 0.01 to 0.1 are usually sufficient for pharmaceutical solutions.
Components pH range HCl, Sodium citrate 1 – 5 Citric acid, Sodium citrate 2.5 - 5.6 Acetic acid, Sodium acetate 3.7 - 5.6 K2HPO4, KH2PO4 5.8 - 8 Na2HPO4, NaH2PO4 6 - 7.5 Borax, Sodium hydroxide 9.2 – 11
H2CO3⇌ H++HCO3 – The pH of blood is controlled by a bicarbonate (H2CO3/HCO3 –) buffer system. When the pH gets too high (high OH concentration), the OH– reacts with carbonic acid (H2CO3) to form HCO3 –) and H2O. When the pH gets too low (high H+ concentration), the H+ reacts with HCO3 – to form H2CO3. Because H2CO3 is a weak acid, the H+ stays associated with the H2CO3. Since pH is an important factor in many physiological processes, a change in the blood pH is a potentially life threatening condition requiring immediate regulation. The pH of blood: The HCO3 -/H2CO3 buffer system is present in blood in greatest concentration and is very important in maintaining the pH of blood within normal limit. The concentration of HCO3 - and H2CO3 in blood are 0.02M and 0.00125M respectively and hence the [HCO3 -]/[H2CO3] ratio is 20/1. In blood, the pKa value for first ionization stage at body temperature is 6.1. pH = pKa + Log [salt]/[Acid] = 6.1+ Log20/1= 6.1+1.2=7.4 The phosphoric and protein buffers of plasma are of relatively little important as compared with bicarbonate buffer in regulating pH.
Buffer is very important for biological system. Some of the pictures are as follows 1. Buffer maintains constant [H+] in the body required for optimum cellular activity. 2. The pH of blood (around 7.4) is controlled by a bicarbonate (H2CO3/HCO3 –) buffer system. 3. The phosphate buffer system (HPO4 2-/H2PO4 -) plays a role in plasma and erythrocytes. H2PO4 -+H2O⇌H3O++HPO4 2- M/A: Any acid reacts with monohydrogen phosphate to form dihydrogen phosphate HPO4 2- + H3O+ H2PO4 - + H2O monohydrogen phosphate dihydrogen phosphate The base is neutralized by dihydrogen phosphate H2PO4 - + OH- HPO4 2- + H3O+ dihydrogen phosphate monohydrogen phosphate 4. Proteins as a buffer: Proteins contain –COO- groups, which, like acetate ions (CH3COO-), can act as proton acceptors. Proteins also contain –NH3 + groups, which, like ammonium ions (NH4 +), can donate protons. M/A: If acid comes into blood, hydronium ions can be neutralized by the –COO- groups -COO- + H3O+- COOH + H2O If base is added, it can be neutralized by the –NH3 + groups -NH3 + + OH-- NH2 + H2O
PROTEIN BUFFER SYSTEM  Proteins are very large, complex molecules in comparison to the size and complexities of acids or bases  Proteins are surrounded by a multitude of negative charges on the outside and numerous positive charges in the crevices of the molecule - - - - - - - - - - - - ---------- --- - - - - - - - - + +++ + + + + ++ + + + ++ + + + + + + + + + + Md. Imran Nur Manik
PROTEIN BUFFER SYSTEM  H+ ions are attracted to and held from chemical interaction by the negative charges - - - - - - - - - - - - ---------- --- - - - - - - - - + +++ + + + + ++ + + + ++ + + + + + + + + + + H+ H+ H+ H+ H+ H+ H+ H+ H+ H+ H+ H+ H+ H+ H+H+H+H+H+H+H+ Md. Imran Nur Manik
PROTEIN BUFFER SYSTEM  OH- ions which are the basis of alkalosis are attracted by the positive charges in the crevices of the protein - - - - - - - - - - - - ---------- --- - - - - - - - - + +++ + + + + ++ + + + ++ + + + + + + + + + + OH- OH- OH- OH- OH- OH- OH- OH- OH-OH- OH- OH- Md. Imran Nur Manik
PROTEIN BUFFER SYSTEM - - - - - - - - - - - - ---------- --- - - - - - - - - + +++ + + + + ++ + + + ++ + + + + + + + + + + OH- OH- OH- OH- OH- OH- OH- OH- OH-OH- OH- OH- H+ H+ H+ H+ H+ H+ H+ H+ H+ H+ H+ H+ H+ H+ H+H+H+H+H+H+H+ Md. Imran Nur Manik
To prepare a pharmaceutical buffer solution having definite pH and capacity, the pharmacist should maintain the following steps: 1. Select a weak acid having a pKa approximately equal to the pH at which the buffer is to be used. This will ensure maximum buffer capacity. 2. From the buffer equation, calculate the ratio of salt and weak acid required to obtain the desired pH. (4-10) 3. Consider the individual concentrations of buffer salt and acid needed to obtain a suitable buffer capacity. (A concentration of 0.05 to 0.5M is usually sufficient; and a buffer capacity of 0.01 to 0.1 is generally adequate.) 4. Other factors of some importance in the choice of a pharmaceutical buffer include- • Availability of chemicals • Sterility of the final solution • Stability of the drug and buffer on aging • Cost of materials • Freedom from toxicity e.g. a borate buffer, because of its toxic effects cannot be used to stabilize a solution to be administered orally or parenterally. 5. Finally, one should determine the pH and buffer capacity of the completed buffer solution using a reliable pH meter.
Osmotic Pressure The flow of the solvent through a semipermeable membrane from pure solvent to solution or from a dilute solution to concentrated solution is termed osmosis (Greek Osmos means “to push”.) Osmotic pressure may be defined as the external pressure applied to the solution in order to stop the osmosis of the solvent into the solution separated by a semipermeable membrane. A membrane which is permeable to solvent and not to solute is called semipermeable membrane. Animal and vegetable membranes are not completely semipermeable. Cupric ferrocyanide, Cu2Fe(CN)6, membrane deposited in the walls of porous pot is perfectly a semipermeable membrane. Md. Imran Nur Manik
The flow of the solvent through a semipermeable membrane from pure solvent to solution or from a dilute solution to concentrated solution is termed osmosis. Osmotic pressure (π) may be defined as the external pressure applied to the solution in order to stop the osmosis of the solvent into the solution separated by a semipermeable membrane. Isotonic Solutions Solutions having the same osmotic pressure are said to be isotonic. In terms of physiological fluids, the solutions having osmotic pressure equal to the osmotic pressure of intracellular fluid is called isotonic solution. (πsoln = πcell) In pharmacy and medical science, isotonic solution is that solution; which have equal tonicity with body fluid i.e. blood, serum, plasma or lacrimal fluid. 0.9% NaCl solution is also regarded as isotonic solution.
Hypertonic Solutions As compared to the blood plasma if a solution has higher osmotic pressure is said to be hypertonic solution. (πsoln >πcell). Physiological solutions having a greater osmotic pressure than that of body fluid or 0.9% NaCl solution is referred to as hypertonic solution. Hypotonic Solutions As compared to the blood plasma if a solution has lower osmotic pressure is said to be hypotonic solution. (πsoln <πcell). Physiological solutions with an osmotic pressure lower than that of body fluid or 0.9% NaCl solution is referred to as hypertonic solution. Paratonic Solutions The solution that is not isotonic that means both the hypertonic and hypotonic solutions are called paratonic solution.
Effect of tonicity on body( Injection on blood) The solutions which are not isotonic with plasma may be harmful to use. On injecting the hypotonic solutions into blood stream, it may enter the blood cells in an attempt to produce equilibrium, the cells swells rapidly until they burst leading to hemolysis. As this damage is irreversible may lead to serious danger to RBC. When hypertonic solution is injected into the blood stream, the water comes out of the membrane of RBC in order to reach equilibrium. The cells shrink leading to crenulation which is only a temporary damage. When the osmotic pressure of two solutions becomes equal the damaged cells will come to its original position. Hence hypertonic solutions may therefore be administered without permanent damage to the blood cells. They should be injected slowly to ensure rapid dilution into the blood stream and to minimize the crenulation of blood cells.
Calculation for the Preparation of Isotonic Solutions For the preparation of isonotic solutions, the quantities of substances to be added may be calculated by the following methods: 1. Based on the freezing point data (Freezing point depression). 2. Based on molecular concentration. 3. Based on Sodium Chloride equivalents. 4. Graphical method based on vapor pressure and freezing point determinations. Method: Based on freezing point depression. The freezing point is a colligative property often used in the calculation of the isotonic solution as it can be measured easily and accurately. The temperature at which blood plasma and tears (Lacrhrymal secretions) freeze is –0.52ºC which is the same value of a 0.9% solution of NaCl. All solutions which freeze at –0.52ºC will be isotonic with blood plasma and lachrymal fluid. The freezing points are usually expressed in terms of 1% solutions.
The quantity of needed for making the solutions isotonic with blood plasma may be calculated from the general formula given below: Percentage W/V of needed= Where, a= freezing point depression of unadjusted solution. b= freezing point depression of 1% W/V of the . Problem: Find out the concentration of required to render or make a 1% solution of cocaine hydrochloride isotonic with blood plasma. The freezing point of 1% W/V solution of cocaine hydrochloride is –0.090ºC and that of NaCl is –0.576ºC. Problem: Find out the concentration of required to render or make a 1.5% solution of proocaine hydrochloride isotonic with blood plasma. The freezing point of 1% W/V solution of proocaine hydrochloride is –0.122ºC and that of is –0.576ºC.
Method: Based on Sodium Chloride Equivalents This method has gained popularity. NaCl method is defined as the weight of solution chloride which will produce the same osmotic effect as 1 g of the drug to prepare an isotonic solution. Formulla: Percentage of NaCl for adjustment to isotonicity=0.9–Percent strength of drug solutionNaCl eqivalent of the the drug. Problem: Calculate the percentage of KNO3 required to make a 0.5% isotonic solution of AgNO3.The NaCl equivalent of KNO3 is 0.56 and NaCl equivalent of AgNO3 is 0.33. Hints: 0.9-(0.50.33)=0.735÷0.56=1.313% Md. Imran Nur Manik
Md. Imran Nur Manik