States of Matter Related to Pharmaceutical Formulations PDF

States of Matter Related to Pharmaceutical Formulations  When molecules interact with each other, they do so by the forces of both Intermolecu attraction and repulsion. Forces of attraction are essential for molecules to lar Forces...

States of Matter Related to Pharmaceutical Formulations  When molecules interact with each other, they do so by the forces of both Intermolecu attraction and repulsion. Forces of attraction are essential for molecules to lar Forces come together. Cohesive Forces – similar molecules Adhesive Forces – different molecules  Inverse Relationship between Attractive force and Distance between molecules: where n ~ 6 or 7 for the interaction of two hydrogen atoms or nitrogen atoms and n ~ 3 or 4 for chloroform molecules.  This inverse relationship between the forces of attraction and distance is known as the Lennard- Jones potential as described by John Edward Lennard-Jones Attractive forces can be represented by means of a potential energy function. As the forces of attraction between the molecules increase, the potential energy becomes increasingly negative,  Shortrange repulsive forces operate when the molecules come close together and electron clouds interact. Repulsive forces (FR) are proportional to an exponential relationship with the reciprocal of the distance separating the molecules (r)  Putting the interactive force components together gives the net potential energy diagram in Figure 2- 3. The quantitative details of the curve vary from one type of molecule to another, but the general form is the same. As two molecules approach each other, the energy changes are gradual and negative due to attractive forces then the energy starts rising rapidly as the intermolecular distances become smaller and repulsive forces begin to dominate. The distance between the molecules at which the attractive and repulsive forces just balance each other is the collision diameter.  4 main types Van der waals attractive forces Intermolecul Dipole-dipole, dipole- induced dipole, induced ar Attractive dipole-induced dipole) Ion-dipole forces Forces Hydrogen bond  van der Keesom forces occur when polar molecules possessing Waals permanent dipoles, having both a partial positively charged end and a partial Forces negatively charged end, interact. These molecules align (keesom themselves so that the negative pole of one molecule interacts with the positive pole of the other forces) The energy of this attraction ranges from 1 to 7 kcal/mole.  A polar molecule can produce a Dipole- temporary electric dipole in nonpolar molecules that are easily induced polarizable The forces of attraction are weaker, dipole being about half those of dipole- dipole forces forces Dipole induced-dipole forces are generally lower in energy, but (Debye multiple interactions can have a stabilizing effect on states of matter forces)  Induced These are forces that originate dipole- from molecular internal vibrations in nonpolar induced molecules to produce attraction that arises because of dipole or synchronized fluctuating dipoles in neighboring atoms. dispersion These forces are also forces temporary. This force is responsible for the liquefaction (London of gases. The energy of this attractive force is 0.5 to 1 kcal/ forces). mole.  These forces are also temporary. This force is responsible for the liquefaction Summary of gases. The energy of this attractive force is 0.5 to 1 kcal/ mole.  Molecules that are polar are attracted to either positive or negative charges. Ion Dipole Pharmaceutical salts will have ion Forces dipole forces holding the drug molecule and the counterion together Ion-induced dipole The forces of attraction are induced by the close proximity of a charged ion to the nonpolar molecule These attractive forces account for the solubility of ionic crystalline substances in water, where a cation attracts the large negative oxygen of water and an anion attracts the hydrogen atoms. The Hydrogen Bond A specific type of electrostatic attraction between molecules and parts of molecules is the hydrogen bond Hydrogen bonding is the attraction of a hydrogen atom for a strongly electronegative atom such as oxygen, nitrogen, fluoride, and, to some extent, sulfur. Hydrogen can form an electrostatic bond to these atoms because of its small size, generating an intense electrostatic field.  The hydrogen bond is considered a uniquely strong type of dipole-dipole interaction. However, the hydrogen bond is also partly covalent in nature, exhibiting such “covalent” bonding properties as bonding strength, unidirectionality, and short interatomic distances Hydrophobic interaction Hydrophobic interactions are forces of attraction between nonpolar atoms and molecules in water. They cause the nonpolar species to be driven together and are critical for the structure and stabilization of many molecules including proteins and aggregates of amphiphiles Hydrophobic attractive forces are less due to interactions between nonpolar molecules than to the lack of ability of these molecules to bond with water molecules.  When these non–hydrogen bonding nonpolar molecules are introduced into water, the hydrogen bonding network between water molecules is disrupted. To minimize this disruption and reestablish a hydrogen bonding network, the water molecules reorient to form a structured water “cage” around the nonpolar molecules, thereby stabilizing the interaction between them. STATES OF MATTER In addition to their potential energy, molecules of gases, liquids, and solids have kinetic or thermal energy. Kinetic Energy is proportional to temperature Gaseous States Gases are described as molecules that have higher kinetic energy that produces rapid motion, are held together by weak intermolecular forces, have no regular shape, are capable of filling all available space, are compressible, and, for many gases, are invisible  Gases at normal temperatures and pressure cand be defined by the ideal gas laws such as Boyle’s and Charles’ Boyle’s law = Inverse proportionality between Volume and Pressure Charles’ law = Direct proportionality between Temperature and Volume Ideal Gas Formula Kinetic Molecular Theory The kinetic molecular theory is a fundamental concept in the field of physics and chemistry that describes the behavior of particles (atoms or molecules) in gases. It provides a theoretical framework for understanding the macroscopic properties of gases based on the motion and interactions of their constituent particles.  Particles in Constant Motion: Gas particles are in constant random motion, moving in straight lines until they collide with other particles or the container walls. These collisions are perfectly elastic, meaning no energy is lost during the collision. Negligible Volume: Gas particles are considered to have negligible volume compared to the total volume of the container they occupy. In other words, the volume occupied by the particles themselves is much smaller than the space between them. No Attractive or Repulsive Forces: Gas particles are assumed to have no significant attractive or repulsive forces between them. They only interact during collisions, and these interactions are purely elastic.  High Kinetic Energy: Gas particles possess kinetic energy due to their motion. The temperature of a gas is directly proportional to the average kinetic energy of its particles. Temperature Determines Kinetic Energy: The temperature of a gas is a measure of the average kinetic energy of its particles. As temperature increases, the average kinetic energy of the particles also increases. Equal Distribution of Energy: In a gas at a given temperature, all particle speeds follow a distribution known as the Maxwell-Boltzmann distribution. This means that while the average kinetic energy increases with temperature, particles of varying speeds are present in the gas. Pressure Arises from Collisions: The pressure exerted by a gas on its container walls results from the constant collisions of gas particles with those walls. More collisions lead to higher pressure. Blood Gases O2, CO2 Plasma gas concentrations are determined and expressed in terms of Henry’s law of gas solubility "At a constant temperature, the amount of a given gas that dissolves in a given type and volume of liquid is directly proportional to the partial pressure of that gas in equilibrium with that liquid” Solubility of a gas in a liquid is directly proportional to the partial pressure of the gas above the liquid:  Dalton’s law of partial pressures – the total pressure is equal to the sum of the partial pressures in the solution We measure blood gas as mmHg PO2 = Partial pressure of O2 PCO2 = Partial pressure of CO2 Normal partial pressure on average of O2 is 80mmHg, PCO2 is around 35-45mmHg  PCO2 refers to the pressure exerted by dissolved CO2 gas in the blood. PCO2 is influenced by respiratory function PCO2 is generally considered to reflect the inverse relationship between itself and lung ventilation If PCO2 is high = poor ventilation If PCO2 is low = hyper ventilation Liquid State A liquid occupies a definite volume and takes the shape of the container required to hold it. Liquids are denser than gases and possess less kinetic energy than do gases They are considered less compressible than gases and more compressible than solids. Liquids flow very readily, and the flow is influenced by friction - Viscosity. They can also be frozen (become solids), have boiling points (become gases), and have vapor pressure and surface tension Liquefaction of Gases When a gas is cooled, it loses kinetic energy as heat, causing the molecules' velocity to decrease. Applying pressure to the gas brings molecules within the range of van der Waals interaction forces, causing the gas to transition into a liquid state. Liquids, due to these forces, are denser and have a definite volume compared to gases. Transitions from gas to liquid and from liquid to solid depend on both temperature and pressure. Elevating the temperature past a critical point prevents gas liquefaction, leading to the concept of critical temperature.  The critical temperature marks the point beyond which a gas cannot be liquefied regardless of applied pressure. Critical pressure is the highest vapor pressure of the liquid at this state. Cooling below critical temperature progressively decreases pressure needed for liquefaction. All known gases can be liquefied by cooling and applying appropriate pressure. Supercritical fluids, involving extreme temperature and pressure, form an intermediate phase. Aerosols: Application of Gaseous Liquefaction Gases can be liquefied under high pressures in a closed chamber as long as the chamber is maintained below the critical temperature. When the pressure is reduced, the molecules expand and the liquid reverts to a gas In aerosols a drug is dissolved or suspended in a propellant, a material that is liquid under the pressure conditions existing inside the container but that forms a gas under normal atmospheric conditions  The container is so designed that, by depressing a valve, some of the drug–propellant mixture is expelled owing to the excess pressure inside the container. If the drug is nonvolatile, it forms a fine spray as it leaves the valve orifice; at the same time, the liquid propellant vaporizes off. Vapor Pressure Vapor pressure is the pressure exerted by the vapor (gaseous phase) of a substance in equilibrium with its liquid or solid phase in a closed container. It is a measure of how easily a substance can evaporate or transition from a liquid or solid state into a vapor at a specific temperature. Substances with higher vapor pressures at a given temperature are more volatile and tend to evaporate more readily. Clausius–Clapeyron Equation: Heat of Vaporization The Clausius–Clapeyron equation is a fundamental equation in thermodynamics that describes the relationship between the vapor pressure of a substance and its temperature during phase transitions, specifically between a liquid and its vapor. The equation mathematically relates the rate of change of vapor pressure to the heat of vaporization and the temperature. It can be used to estimate vapor pressures at different temperatures and to predict the conditions under which phase transitions occur.  Heat of vaporization - amount of heat energy required to transform a given quantity of a substance from a liquid to a vapor at a constant temperature and pressure. The heat of vaporization is a measure of how much energy is needed to overcome intermolecular forces that holds the liquid from transitioning to a vapor. It varies with the substance and is typically expressed in units of joules per mole (J/mol) or calories per gram (cal/g). Boiling Point The temperature at which the vapor pressure of the liquid equals the external or atmospheric pressure is known as the boiling point Surface Tension Surface tension is another physical property of liquids. The surface of a liquid is considered two- dimensional. Therefore, the dimensions of surface tension are force per unit length;  Solid State Solids are characterized as having a fixed shape and being nearly incompressible compared to gases and liquids. They have strong intermolecular forces and therefore very little kinetic energy Solids Solids are characterized by shape, particle size, and melting point; some solids are volatile enough to have a sublimation point. Other characteristics of solids, such as those having importance in the manufacture of solid dosage forms such as tablets, are surface energy, hardness, elastic properties, compaction, and porosity. Crystalline Solids In crystalline solids, the molecules or atoms are arranged in repetitious three-dimensional lattice units infinitely throughout the crystal. Crystals usually have definite melting points Varieties of Crystalline materials Homomeric – identical Heteromeric – different crystal forms Solvates Salt Crystals Cocrystals Polymorphism Polymorphs are chemical entities, including pharmaceutical agents, that may exist in more than one crystalline structure, Polymorphs have different physical properties, including different melting points, solubilities, and stability Ritonavir Case Study Ritonavir is a protease inhibitor used to treat AIDS. It was discovered in 1992 and was one of the first treatments on the market for AIDS patients. It was crystalline but was not bioavailable as the crystalline solid.  It was marketed in 1996 as Norvir®, an oral liquid and semisolid capsules. Both formulations contained ritonavir dissolved in ethanol/water–based solutions. Only one crystalline form was identified during development. Two-hundred-forty lots of Norvir® capsules were produced without incident. In 1998, Norvir® capsules failed dissolution. A new polymorph, Form II, was identified in the capsules and was found to be significantly less soluble than Form I  The ethanol/water solutions for the capsules were not saturated with respect to Form I but were 400% supersaturated with respect to Form II. Oral solutions could no longer be stored at the recommended 2° to 8°C without risk of crystallization. The presence of Form II made the original formulation unmanufacturable and limited supplies were available  The product had to be removed from the market. An oral solution was available in the interim, but due to the extremely poor taste of ritonavir, many patients were not willing to switch. A new formulation was developed using Form II. It is estimated that hundreds of millions of dollars were spent trying to recover Form I and $250 million lost in sale Solvates and Hydrates Solids can also incorporate other compounds into the lattice to produce new crystalline materials. When water is included in a lattice it is called a hydrate. When a solvent is incorporated into the lattice it is called a solvate. Most solvates are not chosen as drug substances due to the possible toxicity of common solvents, however, hydrates are commonly used as drug substances.  Most solvates are not chosen as drug substances due to the possible toxicity of common solvents, however, hydrates are commonly used as drug substances. This difference in solubility, if significant, can lead to a lower bioavailability for the hydrate compared to the anhydrous form. Salt Crystals The crystal lattice can also accommodate other molecules such as acids and bases to form salts. When there is a pKa difference between the molecules, two ionized species would be formed The ionized compounds will interact in the lattice to form a crystalline salt, it may be a weak acid or a weak base. This is called a counterion The counterion should be non-toxic Salts will have different properties from the free acid or the free base that was used to make the salt  Crystalline salts can also exhibit different forms such a polymorphs, hydrates, and solvate Cocrystal A cocrystal is simply defined as a homogeneous, multicomponent phase of fixed stoichiometry where the chemical entities are held together in a crystal lattice by intermolecular forces. Cocrystals are a good option to change properties when an ionizable group is not available. Amorphous Solids Solid material is referred to as amorphous, when there is no long-range order over many molecular units to produce a lattice or crystalline structure Maybe referred to as glasses or supercooled liquids because of the random order of arrangement Because of weaker interactions between molecules, amorphous materials are less stable  The amorphous form of a drug will usually be more soluble than crystalline materials since there is no crystalline lattice that needs to be broken before the material can dissolve A second form of amorphous material used in drug products is an amorphous dispersion. An amorphous drug is stabilized by a polymer or combinations of polymers and surfactants. These has increased solubility with stability closer to a crystal material Polymeric Solids Polymers are large molecules formed by covalent assembly of small molecules (monomers) to form long repeating chains A variety of polymers are used to make the amorphous dispersions discussed in the previous section. The polymers help stabilize the amorphous drug in the solid state and may help prevent crystallization upon dissolution Stability of Solids Chemical Stability Chemical stability can result from a number of different chemical reactions depending on the molecule, conditions, and other components of the dosage form Physical Stability Another type of stability that needs to be addressed is physical stability or the ability of the solid form to resist change upon standing/storage or under stress/ processing conditions