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Chapter: Pharmaceutical Drugs and Dosage: Pharmaceutical solutions

The required concentration of API in an aqueous solution is determined by the drug’s dose and reasonable amount of solution that can be administered.


The required concentration of API in an aqueous solution is determined by the drug’s dose and reasonable amount of solution that can be admin-istered. In addition, factors such as drug’s solubility and taste play a role in determining drug concentration. For example, the taste of bitter or unpleasant drugs tends to be concentration dependent. In addition, taste-masking strategies, such as drug adsorption to ion-exchange resin, limit the maximum drug concentration in solution depending on the maximum amount of drug that can be adsorbed on the resin and resin concentration in solution.

Solubilization of the API, that is, increasing the soluble concentration of the API in the vehicle, is frequently required to prepare aqueous solutions. The most commonly used approaches for solubilizing API are the use of one or more of pH control, surfactant(s), and/or cosolvent(s). Drugs that are poorly soluble in water may be dissolved in a mixture of water and a water-miscible solvent such as alcohol, glycerol, polyethylene glycol, or propylene glycol. The proper selection of a solvent depends on the physicochemical characteristics of the solute and the solvent.

Temperature is an important factor in determining the solubility of a drug and in preparing its solution. Sometimes the manufacturing process requires the use of elevated temperature to prepare a solution. After manu-facturing, the solution can be brought back to room temperature without drug precipitation or crystallization.

pH and buffer capacity

The pH of the vehicle is an important determinant of solubility of an ion-izable drug. Most drugs are weak acids (e.g., having a carboxylate group) or weak bases (e.g., having an amine group). Weak acids are ionized at basic pH. Weak bases are ionized at acidic pH. Ionized forms of the drugs are more soluble than unionized forms. Thus, pH affects the solubility of the drug. Depending on the slope of the pH-solubility profile of a drug, a slight increase or decrease in pH can cause some drugs to precipitate from a solution. Therefore, an adjustment of pH can aid in solubilizing ionizable drugs, and use of buffer to prevent pH shift on storage can minimize the risk of precipitation or crystallization.

Buffers are binary mixtures of compounds in solution that resist changes in solution pH upon the addition of small quantities of acid or base. These binary mixtures could be (a) a combination of a weak acid and its conjugate base (i.e., its salt) or (b) a combination of a weak base and its conjugate acid (i.e., its salt). A weak acid buffer is exemplified by the combination of acetic acid and sodium hydroxide, which forms the salt sodium acetate. A weak base buffer is exemplified by histamine and hydrochloric acid, which forms protonated histamine chloride salt. Buffer solutions are generally not prepared from weak bases and their salts because bases are usually highly volatile and unstable.

pH and buffering capacity

The most important characteristics of a buffer solution are its pH, which can be calculated using the Henderson–Hasselbach equation, and its buffer capacity, which is defined as the magnitude of the resistance of a buffer to pH changes. The stable pH of the solution generated by a buffer depends on the concentration of the two species and the pKa of the weak acid or the weak base. It is determined by the Henderson– Hasselbalch equation.

The extent to which a buffer resists change in solution pH is known as the buffering capacity. Buffering capacity of a buffer is related to the concentra-tion of the acid and the base, that is higher the concentration, greater the buffering capacity. Buffering capacity is generally expressed as the concen-tration of the buffer. Thus, a 2 M acetate buffer has 10× more buffering capacity than a 0.2 M acetate buffer. The ratio of the acetate salt to the acetic acid may be the same in both buffers.

If strong acid, such as 0.1N HCl, is added to a 0.02 M solution contain-ing equal amounts of acetic acid and sodium acetate, the pH is changed only 0.09 pH units because the base acetate (Ac) ties up the hydrogen ions according to the reaction:

Ac + H 3O+ HAc + H2O

If strong base, such as 0.1 N NaOH, is added to the buffer mixture, acetic acid neutralizes the hydroxyl ions as follows:

HAc + OH H 2O + Ac

Example of How to Make a Buffer


An acetate buffer is created by the addition of sodium acetate to acetic acid. Alternatively, sodium hydroxide can be added to a solution of acetic acid. In the presence of the strong base, sodium hydroxide, an equimolar amount of acetic acid, converts to the sodium acetate salt or the acetate ion in situ.

When sodium acetate is added to acetic acid, the dissociation con-stant, Ka, for the weak acid is expressed by the equation:

The dissociation constant is a known constant for each acid. The pKa of acetic acid is 4.75. This means that an equal concentration of acetic acid and sodium acetate in solution will result in a solution pH of 4.75.

The pH of the final solution is obtained by rearranging the equilib-rium expression for acetic acid:

The aforementioned equation can be expressed in logarithmic form, with the sign reversed as follows:

log [H3O+] = − logKa – log [Acid] + log [Salt]

This is the Henderson–Hasselbalch equation for a weak acid:

pH = pKa + log [Salt]/ [Acid]

The term, pKa, is the negative logarithm of Ka, which is called the dis-sociation constant.


The buffer equation for solutions of weak base and their salts can be derived in a manner similar to that for the weak acid buffers. Accordingly,

[OH ] = pKa [Salt]/[Base]

Using the relationship [OH] = Kw/[H3O+], we can obtain the follow-ing buffer equation:

pH = pKw pKb + log [Salt]/[Base]

Surfactants and cosolvents

Surfactants are commonly used in the dosage form to impart an amphiphi-lic character to the aqueous vehicle and/or associate with the hydrophobic drug to increase its solubility. When low concentrations of surfactants are added to the aqueous solution, they associate with the hydrophobic parts of a solute and increase the solubility of the solute in a concentration-dependent manner. At a certain concentration, known as the critical micelle concentra-tion (CMC), there are enough surfactant molecules in solution that several surfactant molecules self-associate, with hydrophobic parts of the molecule buried inside and the hydrophilic part on the outside, facing the aqueous environment, to make structures known as micelles. Typical micelles contain 6–12 molecules of the surfactant. Micelles are subvisible soluble colloidal structures with a hydrophobic core. This allows the partition and retention of hydrophobic drug in the core of the micelle, thus dramatically increasing total drug solubility. The slope of concentration dependence of solubilization of a solute by a surfactant is significantly higher above the CMC than below.

Cosolvents increase drug solubility by altering the dielectric constant and hydrogen bonding capability of the vehicle and by providing a hydrophobic microenvironment. Commonly used cosolvents include ethanol, polyethyl-ene glycol, and propylene glycol. In addition, cyclic polysaccharides, such as cyclodextrins, that have a hydrophobic cavity and a hydrophilic exterior are often used for drug solubilization.

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