Chemical Stability Studies
Title : Chemical Stability Studies
Course Co ordinator : Dr. Smt. Jayanti Vijaya Ratana.
Let us try to understand what is known as chemical stability. I said earlier that the rate of decomposition depends on the frequency and intensity of the collisions of the molecules. Let us take off in this class from that point.
When you cut an apple and expose it to the air very quickly it gets deep brown colour – it is getting oxidized. Milk is being turned into curd by a reaction known as fermentation. Every one knows the smell of a poorly stored sample of oil – which is called a rancid oil. What has happened to the oil? It has been hydrolyzed! Lipolytic enzymes in the presence of moisture hydrolyze oils and fats producing glycerol and free fatty acids and thus develop rancidity in fats and oils. So to store oil properly we have to exclude air, light, microorganisms and moisture. In all these reactions; hydrolysis, or oxidation or fermentation a common villain is – moisture. All reactions are speeded up in the presence of moisture.
Why? Moisture gives a fertile ground for the growth of microorganisms. Moisture takes part as a reactant in many chemical reactions and plays the role of a solvent vector in many reactions. It has better thermal conductivity than solids and allows better heat transfer – hence molecules have more kinetic energy and you observe more decomposition. So, we can put a ditto below foods and write about drugs. All that is true for foods is also true for drugs.
But I want to take you a little deeper first and make you see how a chemical reaction is happening – whether it is oxidation or hydrolysis. Let us try to remember the second law of thermodynamics – a spontaneous reaction always proceeds towards greater entropy (more disorder) and towards loss of free energy. A chemical reaction also tends to occur if it increases the disorder of the system.
Remember that we are talking about several types of reactions, such as, (1) the splitting up of a molecule into two units, (2) the transformation of an isomer from one form to another, (3) the reaction between two different molecules, and so on. For all these reactions some basic steps are necessary. I will explain a unimolecular reaction first, with the example of the transformation of a Cis – dichloroethylene to trans – dichloroethylene.
Fig.1: Energy profile for the conversion of Cis – to trans – dichloroethylene
The energy of a molecule going through this conversion process is represented in figure-1. A molecule of the cis compound which is stable and is having low energy, picks up energy as it changes to the transition state configuration and then releases energy as it rolls downhill to the trans form. The abscissa of the graph, called the reaction coordinate is a measure of the progress of the reaction. For this reaction to take place, a molecule of cis – dichlorothylene must gain an amount of energy shown in Figure – 1 as Ea, which is called the activation energy for the process. However, all the molecules which have sufficient activation energy do not transform into the product; some of them shed their excess energy and return to the starting state rather than proceeding to form the reaction product. We can write a rate equation as:
Where [A] = concentration of reactant A
When experiments were done on several chemical reactions, the exponential dependence of rates on temperature was established. The rate for energized molecules is a temperature independent factor and varies over wide limits (109 – 1016 sec-1) but many results cluster around the value of 1013 sec-1.
So for a unimolecular reaction
Rate = k [A]
K = ?e-(E/RT) …(1)
Where ? = frequency factor
= rate for energized molecules
2.303 logk = 2.303 log? – Ea/RT
In a bimolelcular reaction two molecules react with each other to form products. Molecules must collide to react. These collision rates are much faster and only a fraction of the molecules colliding result in reaction. The rates of these reactions also increase with temperature just as unimolecular reactions do and they also have to reach or attain an activation energy to result in a reaction.
Taking an example of the reaction of a hydroxide ion with methyl iodide,
HO- + CH3 I Ch3OH + I-
Fig.2: Energy profile for the reaction of hydroxide ion with methyl iodide
Just as we have written for the unimolecular reaction let us write a rate expression in terms of molecular collisions for the bimolecular reaction also.
A + B products
the rate is
The last factor in the rate expression is called the “orientation” factor. It tells us that all the energetic collisions do not result in a reaction – this is because they are not properly oriented to cause a reaction. This term corresponds to the ? factor in the unimolecular rate expression (eq.1).
In this example, collision of a hydroxide ion with the iodine atom of methyl iodide would not result in reaction. Only collisions involving direct contact of the oxygen atom with the carbon atom would be effective in causing reaction to occur. The collision rate is a number that depends on the concentrations of the molecules colliding.
You can see that the expression for k is the same as that for a unimolecular reaction. But the difference lies in the interpretation and units of ?. For bimolecular rate processes, ? has units of liter mole-1 sec-1 whereas for unimolecular reactions, ? is given in sec-1. On the basis of their experience scientists give us these guidelines:
(1)A typical reaction will proceed at an appreciable rate at room temperature if Ea is 10 kcal mole-1 or less.
(2)For such a reaction, the fraction of high energy molecules increases with temperature in such a way that a 10o rise in temperature will roughly double the reaction rate.
Now, I want to compare a unimolecular reaction to a person trying to hit a goal in a game of basketball. The man gathers strength, raises and throws the ball.
His attempts are many, every time he gathers sufficient energy but the hits are few and the ‘hit’ comes only when every move is properly coordinated, i.e. the right orientation is there
To understand a bimolecular reaction visualize two young girls playing a skipping game with a long rope. One girl swings the rope, starting from below her feet, coming from behind – above her head and in its forward sweep it includes the second girl and comes from beneath her feet by then one swing is over. The second girl jumps into the air as the rope is going below her feet. The girls try a number of times, but perfect smooth circle of the rope is completed only some times, i.e. when both the girls are perfectly oriented to each other and are in perfect coordination.
Degradation in solid dosage forms and solid drugs is usually not much because of one strong reason i.e., the usual range for activation energies for tablet formulation decomposition is about 10 to 20kcal/mole-1 except if diffusion or photolysis is rate determining. Then the rate is about 2 to 3 kcal/mole-1 which rarely occurs in tablet degradation. For reactions in which the heat activation energies range is more than 50 k cal/mole-1, the rate of degradation is not of any practical significance at the temperature of shelf life storage of solid formulations.
I explained all these points before going into degradation mechanisms because I want you to have an idea of how actually a chemical reaction is taking place. So a formulation suffers chemical instability when the drug degrades and the drug content as claimed on the label is not available in it. Other chemicals which are used in the formulation as excipients may also decompose and though their content is not theraupeutically necessary the decomposition is of high pharmaceutical significance. It may alter the appearance as well as the release rate of the drug.
Decomposition of the drugs is happening when the drug molecules are energized and are taking a particular orientation, but after taking the orientation or positioning, what exactly is happening? Are the molecules splitting into two? Are they loosing some electrons? Or they gaining some electrons? Are they racemizing?
This later step is what we are calling as the pathway or mechanism of degradation. Some well known mechanisms are hydrolysis, oxidation-reduction, racemization, decarboxylation, ring cleavage and photolysis. But the most frequently seen mechanisms are hydrolysis and oxidation reduction. Let us try to understand how these reactions are happening a little later.