States and Properties of Matter and Physicochemical Properties of Drug Molecules

Chapter 2

States and Properties of Matter and Physicochemical Properties of Drug Molecules

States of Matter and Properties of Matter

States of Matter

Matter can be defined as anything that has mass and occupies space. Based on its composition and properties, matter can be classified as elements, pure compounds, pure substances and mixtures.

A substance is a form of matter that has a constant composition. Physicochemical properties of a substance are dependent on the organizational arrangement of its constituent atoms. For example, n-butane has the same chemical formula as iso-butane, C4H10. Physical properties namely, boiling point, melting point and relative density of both these compounds are given in Table 2.1. Vapor pressures of these compounds at a temperature and their chemical properties like reactivity differ due to different arrangement of the same atoms in each molecule. They have different structural formulas as n-butane: CH3-CH2-CH2-CH3 and iso-butane: CH3−CH−(CH3)−CH3 and thus the physicochemical property of substance varies with structural arrangement.

Solid, liquid and gas represent the three basic states of matter as shown however plasma and Bose Einstein condensate are considered as other states of matter. In pharmaceutical viewpoint, basic three states are significant while other two states have limited applications in pharmacy, but they have major applications in physics. The plasma state is not related to blood plasma, but it represents an ionized gas at very high temperatures. There is no sharp distinction between solid, liquid and gaseous states because they may exist in any state depending upon intensity of intermolecular forces and physical forces like temperature and pressure. For a molecule to exist in aggregate as compound there must be some intramolecular binding force. Knowledge of these forces is important to understand the properties of solids, liquids and gases as well as solutions, suspensions, emulsions and powders etc. The summaries properties of solids, liquids and gases that identify their microscopic behavior responsible for each property.

An element cannot be further divided by chemical means whereas compound is a form of substance in which two or more atoms are linked chemically. Molecular compounds can be broken down to pure elements by chemical means and are defined by its atomic number. Some elements have isotopes, radioactive 125I, for example, frequently used in thyroid cancer treatment is an isotope of the stable 127I. All isotopes have the same atomic number, but they have different mass number (i.e. different number of neutrons). Pharmacist frequently uses radioisotopes as a means to study in-vivo fate of biologically active macromolecules and synthetic drug compounds. Radioisotopes are also used in diagnostic applications. A combination of two or more substances is known as mixture, which may or may not retain original physicochemical properties of its constituent components. There are two types of mixture namely, homogeneous mixture and heterogeneous mixture.

Homogenous Mixture:

In homogenous mixture of solid and liquid the chemical and physical properties of individual components cannot be determined by any single instrumental method of analysis. Depending upon temperature substances exist in different states. Aspirin, for example, as shown in Fig. 2.3, indicate that below 135 °C it exists in solid crystalline form whereas above this temperature it exists in liquid form.

Dissolving aspirin crystals in water makes aqueous aspirin solution. Water destroys the intermolecular forces between the aspirin molecules that exist as crystalline arrangement during the process of solution formation. In the formation of a molecular dispersion there must be some mutual interaction between solute and solvent. Thus, the properties of the individual components of the mixture get changed. All the physical properties of aspirin are changed upon interaction with water. Similarly, the properties of water are also get changed by the presence of the aspirin. Another physical property called absorption of electromagnetic radiation is changed due to homogeneous mixing. Halothane, for example, shows different absorption of light in the visible and ultraviolet region as pure liquid and as a solution in organic solvents. The homogeneous mixtures of liquids and solids and mixtures of gases are always homogeneous.

Heterogeneous Mixture:

A heterogeneous mixture is one in which the individual components of the mixture retains their original physicochemical properties. The composition of a heterogeneous mixture may or may not be uniform throughout. The commonest example of heterogeneous liquid mixture is pharmaceutical suspension. Suspensions are liquids in which the insoluble drugs are present in the fine state and are somewhat uniformly dispersed in aqueous media. Kinetic forces exerted by the water molecules on the suspended drug molecules are primarily responsible for their suspension in solvent. The larger particles are more difficult to keep uniformly suspended in the water. Since the drug solubility is less, the physicochemical properties of drug and water in pharmaceutical suspensions remain practically intact.

By means of physical methods components of homogeneous and heterogeneous mixtures can be separated and recovered as pure substances. However, for homogenous mixtures great care need to be taken to recover pure components. For example, water present in simple syrup can be removed by boiling syrup and condensing generated vapours to get back pure water leaving behind the pure dry sugar powder. The sugar is recovered in a pure form, but not in its original, crystalline state. A tablet prepared by direct compression of a drug and other excipients such as lactose, polyvinyl pyrrolidone (PVP) and magnesium stearate is an example of a heterogeneous solid mixture. Lactose powder tries to remain as a separate entity from the magnesium stearate and the solid drug. For the excipients to exert its effect in the tablet, they must retain their distinct identity along with their physicochemical properties within the powder mixture. PVP is the disintegrant and its swelling property facilitates disintegration of tablet in dissolution media. Interaction of PVP with the drug or with any of the other excipients may change or even neutralize its disintegration property. Similarly, interaction of magnesium stearate, a lubricant, with other excipients may eliminate its lubricant properties. But most importantly, active drug-excipient interactions that are not expected could lead to product instability, ineffective therapy or sometimes toxicity. The carbonate salts, for example, is commonly used in effervescent tablets that may cause hydrolysis of an ester drug in the presence of moisture. Similarly, interaction of the drug with excipients may lead to complex formation, which may have reduced solubility that may affect drug performance

Changes in the State of Matter

In the solid-state particles are held near by intermolecular, interatomic or ionic forces therefore the particles of solid oscillate about fixed position. As the temperature of solid is increased, the particles acquire enough energy to breakdown the ordered arrangement of the lattice and pass in to the liquid form. On further application of energy by increasing temperature, liquid molecules pass in to the gaseous state. The transition between different states of matter and the processes involved in these transitions.

The examples of the substances that exist in different physical states are nitrogen (gas), water (liquid) and glucose (solid) under the normal temperature (22 °C) and pressure (1 atm) conditions. Ice water, liquid water and vapour water, is classic example of a substance that exist in three different states. Some solids with high vapour pressure like iodine and camphor can pass directly in to gaseous state without melting called as sublimation. A change in which gas state directly changes to solid state is called condensation. A substance may co-exist in two or three states simultaneously at temperature and pressure conditions. For example, ice in liquid water at temperature very close to freezing point or coexistence of ice, liquid and water vapour at triple point of water.

The changes in the physical states of a substance are reversible in nature. These are due to rearrangement of the molecules in a substance, while on other hand; chemical changes are due to change in specific orientation or arrangement of the atoms and groups of the substance. Chemical changes may be irreversible or completely or partially reversible. Chemical changes always result in a formation of a new compound having different properties. An example of an irreversible chemical change is decomposition of water causing the molecules to breakdown in to new substances hydrogen and oxygen. An example of reversible chemical change is esterification of salicylic acid with malonic anhydride to form aspirin.

Latent Heat

The amount of heat required to raise the temperature of one gram of the solid is called the heat capacity. The temperature of solid continuously increases until it reaches to its melting point. At melting point the temperature will hold steady for a while, even though heat is added to the solid. It will hold steady until the solid completely melts. The temperature rising stops because melting requires energy. All the energy added to a crystalline solid at its melting point goes into melting, and none of it goes into raising the temperature. Then again, the temperature of the solid will begin to increase. This heat is called the latent heat of melting. Once the solid get melted, the temperature begins to rise but at a slower rate. The molten solid (liquid) has a higher heat capacity than the solid crystalline state therefore it absorbs more heat with a smaller increase in temperature. Hence, when a crystalline solid melt it absorbs a certain amount of heat, the latent heat of melting, and it undergoes a change in its heat capacity. Any change like melting, freezing, boiling or condensation brought about by heat which has a change in heat capacity and a latent heat involved, is called a first order transition. But when an amorphous solid is heated to its Tg, the temperature increases. It increases at a rate determined by the solid’s heat capacity. There is no latent heat of glass transition. At Tg, the temperature does not stop rising. The temperature keeps upon increasing above Tg but at different rate than below Tg. The solid does undergo an increase in its heat capacity when it undergoes the glass transition due to change in heat capacity. Any change brought about by heat, which has a change in heat capacity, but a latent heat is not involved, is called a second order transition. In first order transition melting is observed with crystalline solid, and in second order transition the glass transition is observed with amorphous solid.

Vapour Pressure

Physical properties of liquids are controlled by strength and nature of intermolecular attractive forces. The most important properties are vapour pressure, viscosity, surface tension and light absorption and refraction. A liquid placed in a container partially evaporates to establish a pressure of vapour above the liquid. The established pressure depends on the nature of the liquid, and at equilibrium it becomes constant at any given temperature. This constant vapour pressure is the saturated vapour pressure of liquid at that temperature. Until the vapour pressure is maintained, no further evaporation observes. As shown in Fig. 2.6, at lower pressures a liquid evaporates into the vapour phase while at higher pressure the vapour tend to condensate till equilibrium establishes. During vaporization heat is absorbed by liquid. At any given temperature, the amount of heat required per gram of liquid is definite quantity called as heat of vaporization of liquid (∆Hv). It is difference in enthalpies of vapour (Hv) and liquid (Hl), respectively. Therefore, ∆Hv = Hv – Hl … (2.1) During evaporation ∆Hv is always positive while during condensation it becomes always negative. As per definition of change of enthalpy, ∆Hv is the difference in internal energy of vapour and liquid. ∆Hv = ∆Ev + P ∆Vv … (2.2) where, P is vapour pressure and ∆Hv is change in volume during vapour to liquid transition.

The temperature of a substance depends on the average kinetic energy of its molecules. Average kinetic energy is considered because there is an enormous range of kinetic energies for these molecules. Even at temperatures well below the boiling point of a liquid, some of the particles are moving fast enough to escape from the liquid. During this process the average kinetic energy of the liquid decreases. As a result, the liquid becomes cooler. It therefore absorbs energy from its surroundings until it returns to thermal equilibrium. But as soon as this happens, some of the water molecules once again have enough energy to escape from the liquid.

In an open container, this process continues until all the water evaporates. In a closed container, some of the molecules escape from the surface of the liquid to form a vapour. Eventually, the rate at which the liquid evaporates to form a gas becomes equal to the rate at which the vapour condenses to form the liquid. At this point, the system is said to be in equilibrium. As shown in Fig. 2.7, the space above the liquid is saturated with water vapour, and no more water evaporates. The pressure of the water vapour in a closed container at equilibrium is called the vapour pressure.

The relationship between vapour pressure and temperature is not linear. The vapour pressure of water increases more rapidly than the temperature of the system.
Measurement of Vapour Pressure: Vapour pressures of liquids are measured by static and dynamic methods.

Static Method: Vapour pressure of liquid is generally measured by the isoteniscopic method, which is precise, flexible and convenient over a range of temperatures. A simple apparatus is showN. It consists essentially an isoteniscopic bulb of 2 cm diameter.

A liquid under test is filled in bulb-up to half level mark, which is connected to mercury manometer and a pump. The air inside the bulb is removed by application of vacuum. Now there is no air present in the bulb. To maintain equilibrium, part of liquid evaporates. The system is maintained at constant temperature so that the equilibrium between liquid and vapour attains. The generated vapours exert pressure on mercury present in column. The difference in height of mercury in column is determined which is equal to vapour pressure of that liquid. By maintaining the system at any other temperature, it is possible to determine vapour pressure at that temperature. This method is used for liquids having vapour pressures on higher sides close to one atmosphere.

Dynamic Method:

This method is proposed by Walker and is useful especially in determinations of very low vapour pressure of liquid mixtures. Great care is required to obtain excellent results. An illustrative apparatus. An inert gas such as nitrogen is passed through the given liquid at constant temperature. The inert gas is saturated with the vapors of liquid under test and leaves the flask at exit of the tube. If P is total vapour pressure in the apparatus at saturation, n is the moles of gas passed through and nv is number of moles of vapour collected. The nv is given as nv = Wv Mv … (2.3)

where, Wv is loss in weight of liquid and Mv is molecular weight of liquid. The partial pressure of vapour, P’ is same as vapour pressure of liquid at saturation and can be given as where, m is loss in weight of liquid as vapour, V is volume of gas passed through, M is molecular weight of liquid and R is gas constant.

Boiling Point of Liquids: The boiling point is temperature at which vapour pressure of liquid equals the 760 mmHg pressure. However, increasing temperature can boil liquids at any temperature from its freezing point to critical temperature either or decreasing applied external pressure. Hence, boiling point of liquid is temperature at which vapour pressure of liquid is equal to pressure acting on its surface. Boiling is characterized by formation of bubbles within it and release from the surface. The change in boiling point with pressure is calculated if molar heat of vaporization (∆Hv) for the liquid is known. If T1 is the boiling point at pressure P1 and T2 is boiling point at pressure T2 then using equation (2.5) boiling point of liquid at given pressure is obtained. log P1 P2 = ∆Hv 2.303 R T2 − T1 T1T2 … (2.6).

However, when ∆Hv is not known its value is estimated from Trouton’s rule which states that ∆Hv Tb = Constant … (2.7) where, Tb is normal boiling point of liquid on absolute temperature scale. Boiling points of some liquids at one atmospheric pressure are given in Table 2.3.

Boiling Point and Vapour Pressure: Bubbles are formed on heating liquid, which rises to liquid surface and bursts. When liquid vapourizes, the molecule in vapour state remain together as tiny bubbles with the vapour pressure within it. This vapour pressure in bubbles within the liquids is different than that of atmospheric pressure. When a bubble rises to surface, it burst to have equal vapour pressure that of atmosphere. Therefore, boiling point of a liquid is a temperature at which vapour pressure of liquid is equal to atmospheric pressure. Reducing external pressure can reduce the boiling point and at low temperature it is equal to external pressure. Similarly, increasing external pressure increases boiling point of liquid and at high temperature it is equal to external pressure.

Sublimation

Sublimation is another form of phase transitions. Here solid turns directly into a gas. As a sublimating material changes from a solid to a gas, it never passes through the liquid state. As we know water exists in its three forms namely ice, water, and steam. Sublimation is just one of the ways water or another substance can change between its potential phases. Substances such as water and carbon dioxide (CO2) can be plotted on as pressure vs. temperature to understand their state of matter (solid, liquid, or gas) at a given temperature and pressure. At a typical atmospheric pressure, water is a solid at temperatures below 0° C, a liquid from 0 to 100° C, and a gas at higher temperatures. But atmospheric pressure, however, can change, particularly with altitude. Higher altitudes yield lower atmospheric pressures. Water doesn't always change phase at the same temperatures. For example, with lower pressures, liquid water changes to a gas at temperatures lower than 100° C. If the pressure is dropped low enough, water reaches what's known as a triple point. At pressure and temperature of triple point a substance can exist in solid, liquid, and gaseous forms.

Below this point, solid water sublimes, changing directly into a gas with a rise in temperature and never pass through the liquid phase. The CO2 has a triple point at a pressure higher than 1 atmospheric pressure, meaning that at Earth's standard atmospheric pressure, CO2 will sublime as it heats and is converted from solid to a gas.

Critical Point

A liquid need not always have to be heated to its boiling point before it changes to a gas. The kinetic energy of the molecules is proportional to the absolute temperature of the gas. Due to high kinetic energy gas molecules are in the state of constant motion. In liquids, only few molecules have lower or higher kinetic energy. It is illustrated. At low temperature, the number of molecules having high kinetic energy is less as shown by ABCD while at high temperature the number of molecules having higher kinetic energy increases as shown by FBCE. The molecules with high kinetic energy are important to escape from liquid state to vapour state. Upon cooling, kinetic energy gradually decreases. Since the temperature being decreased a stage is attained at which gas molecules loses their energy that they are unable to overcome forces of attraction between them. This situation brings the gas molecules near to have contact with each other achieving more condensed liquid state. This state also can be possible to achieve by increasing pressure of the gas but it has a limitation that pressure is effective only below specific temperature. This temperature is called as critical temperature. It is defined as the temperature above which gas cannot be liquefied, even if very high pressure is applied.

The critical temperature of water is 374 °C or 647 K and its critical pressure is 218 atm. If liquid such as water is sealed in evacuated tube, a specific amount of it evaporates to produce vapour at constant temperature. Like gas, water vapour exerts pressure and maintains equilibrium between liquid and vapour phases. Exerted vapour pressure is characteristic of every liquid and is constant at any given temperature.

The vapour pressure of water at 25 °C is 23.76 mmHg while at 10 °C it is 760 mmHg and therefore it is clear that vapour pressure increases continuously with temperature. As water is heated further, it evaporates to more amount resulting in increased vapour pressure.

When temperature reaches 374 °C the water meniscus becomes invisible. At critical temperature, physical properties of liquid and vapour become identical and no distinction can be made between the two. This point is also called as critical point. The temperature, saturated vapour pressure and molar volume corresponding to this point are designated as critical temperature (Tc), critical pressure (Pc) and critical volume (Vc) respectively. For water these critical constants are; Tc = 374 K, Pc = 219.5 atm and Vc = 58.7 mL/mole. The critical points for different gases are given in Table 2.4.

Eutectic Mixtures

A two-component system containing a solid and liquid in which the two components are completely miscible in the liquid states and are completely immiscible in the solid state. This is because the solid phase consists of pure component. This mixture is known as eutectic mixture. The temperature at which such system exists in liquid phase is known as eutectic temperature. Above this temperature, the components are liquid and below this temperature they are solids. Physically eutectic systems are solid dispersions. Some examples of this type are thymol – salol, thymol – camphor, menthol – camphor etc.

In the melting temperature of two substances A and B are plotted against mixture compositions. The curves separating the regions of A + Liquid and B + Liquid from regions of liquid AB are termed liquidus curves. The horizontal line separating the fields of A + Liquid and B + Liquid from A + B all solid, is termed the solidus. Upon addition of B to A or A to B, their melting points are reduced. The point, E, where the liquidus curves and solidus intersect, is termed the eutectic point. At the eutectic point in this two-component system, all three phases, that is Liquid, crystals of A and crystals of B, all exist in equilibrium. The eutectic point represents a composition (eutectic mixture composition) at which any mixture of A and B has the lowest melting point. Note that the eutectic is the only point on the diagram where this is true. At the eutectic point the maximum numbers of allowable phases are in equilibrium. When this point is reached, the temperature must remain constant until one of the phases disappears. A eutectic is an invariant point. Below eutectic temperature no liquid phase exists.

If we cool solution of A and B which is richer in A than the eutectic mixture, then the crystal of pure A will appear. As the solution is cooled further, more and more of A get crystallize out and the solution becomes richer in B. When the eutectic point is reached, the remaining solution crystallizes out forming a microcrystalline mixture of pure A and pure B. If salol – thymol combinations is to be dispensed as dry powder, it is necessary that the ambient temperature should be below its eutectic point of 13 o C. Above this temperature, it exists in liquefied form. At eutectic point their contribution with respect to composition is 34 % thymol and 66 % salol.

Gases

The gaseous state is the simplest state amongst the three states of matter. A microscopic representation of gaseous state. The molecules in gas are wide apart in empty space and are free to move in any direction in the container they are contained in. The gas molecules exert pressure on the walls of the container in all directions. Gases have indefinite expansion ability to fill the entire container. If movable piston is fitted into container containing gas, then on application of pressure by piston they get easily compressed. When two or more gases placed together they rapidly diffuse throughout each other and form a homogenous mixture. Upon heating gas in the container inside pressure increases and if container is fitted with piston under this condition its volume increases.

Chemical properties of gases vary significantly whereas Physical properties are simpler to understand. Gaseous state can be described by considering small scale action of individual molecules or by large action of the gas. By studying these properties, we can understand the behaviour of gases. The model called as kinetic molecular theory can easily describe the properties.

Kinetic Molecular Theory of Ideal Gases:

The statements made in this theory are only for what is called an ideal gas. They cannot all be rigorously applied to real gases but can be used to explain their observed behavior qualitatively. The kinetic molecular theory is based upon the following postulates:

All matter is composed of tiny discrete particles (molecules or atoms).
Ideal gases consist of small particles (molecules or atoms) that are far apart in comparison to their own size.
These particles are dimensionless points, which occupy zero volume.
These particles are in rapid, random and constant straight-line motion. Well-defined and established laws of motion can describe this motion.
There are no attractive forces between gas molecules or between molecules and the sides of the container with which they collide
Molecules collide with one another and the sides of the container.
Energy can be transferred in collisions among molecules.
Energy is conserved in these collisions, although one molecule may gain energy at the expense of the other.
Energy is distributed among the molecules in a fashion known as the MaxwellBoltzmann Distribution.
At any instant, the molecules in each sample of gas do not at all possess the same amount of energy. The average kinetic energy of all the molecules is proportional to the absolute temperature. Above mentioned postulates are meant for ideal gas only and are only approximately valid for real gases.

Characteristics of Gases:

The volume (V), pressure (P), temperature (T) and the number of moles (n) in the container are measurable characteristic properties of the gas.

Volume:

The volume of container is the volume of gas sample and is expressed in unit liter (L) or milliliter (mL).

Pressure:

Atmospheric pressure is measured using a barometer. If a tube, completely filled with mercury (Hg), is inverted into a dish of mercury, mercury will flow out of the tube until the pressure of the column of mercury equals the pressure of the atmosphere on the surface of the mercury in the dish. The height of the mercury in the tube is 760 mm for 1 atm of pressure. Column of mercury is used to measure pressure of a gas closed in a container. The height ‘h’ of mercury column of manometer, indicate how much higher the pressure of gas is in the container than outside.

Pressure of a gas is proportional to average force per unit area that gas molecules exert on the walls of the container. The greater the number of gas molecules in each container, the higher is the pressure as the greater average number of collisions occurring with the wall of the container. If the volume of the container is reduced, the average number of collisions will increase. Pressure is directly proportional to the kinetic energy of the gas molecules therefore higher the temperature the greater is the kinetic energy and greater the pressure of the gas.

Temperature:

Temperature of gas is measured in Kelvin temperature scale. The product of pressure and volume per mole is proportional to the average molecular kinetic energy. The average kinetic energy is proportional to the absolute temperature.

Number of Moles of Gas:

The concentration of gas in a container can be obtained as ration of mass ‘m’ of the gas sample to the molar mass, M. Moles of gas = Mass (m) Molar mass (M) of the gas … (2.8) Gases are classified into two type namely ideal gases and non-ideal (real) gases. An ideal gas is one that obeys certain laws while real gases are those, which obey these laws only at low pressures. In ideal gases, the volume occupied by its molecules is negligible compared to total volume at all temperatures and pressures and at these conditions the intermolecular attraction is extremely small. In case of real gases both these parameters are appreciable, and magnitude depends on nature, temperature and pressure of gas. Ideal gas is hypothetical gas and real gas contains molecules that have definite volume and intermolecular attraction between each other. When influence of these parameters is negligible gas is considered as ideal gas. This is practically observed at low pressures and high temperatures when the free space between gas molecules is large that very little or negligible attractive forces exist between the molecules.

Gas Laws:

Physical laws describing the behavior of gas under various conditions of pressures, volumes and temperature is known as gas laws. These laws are described below.

Boyle’s Law:

Robert Boyle, in 1662, formulated a generalization that the volume of any definite quantity of gas at constant temperature is inversely proportional to its pressure. Mathematically it is expressed as; V ∝ 1 P or V = k P (when temperature is held constant) … (2.9)

where, V is the volume and P is pressure of the gas whereas k is proportionality constant. This constant is dependent of temperature, weight of gas, its nature and the PV units. The equation (2.9) is the mathematical expression of Boyle’s law; at constant temperature, the volume occupied by a fixed weight of a gas is inversely proportional to the pressure exerted on it. Boyle’s law describes the behavior of an ideal gas and approximates the behaviour of a real gas. The approximation is very poor at high pressures and low temperatures. If in certain condition, as shown in Fig. 2.15, pressure and volume of gas are P1V1 and at any other condition they are P2V2, then at constant temperature, this can be expressed as, P1V1 = k = P2V2 … (2.10)

Charles’s Law:

Charles in 1787 investigated that gases such as hydrogen, carbon dioxide and oxygen expand to an equal amount upon heating from 0 °C to 80 °C at constant pressure. However, Gay-Lussac in 1802 showed that volume of all gases increases with each 1 °C increase in temperature and was approximately equal to 1/273.15 volume of gas at 0 °C. Consider the change in volume of one mole of an ideal gas with the change in temperature when the pressure is held constant.

On simplifying equation (2.11); V2 = V1      (273.15 + t) 273.15 … (2.12) If, 273.15 + t, is designated as T2 and 273.15 as T1 then equation (2.12) becomes; V2 V1 = T2 T1 … (2.13) Therefore, this law is stated as the volume of definite quantity of gas at constant pressure is directly proportional to absolute temperature. It is expressed as; V ∝ T … (2.14) V = kT (when pressure is held fixed) … (2.15) A plot of the volumes at various temperatures is given in Fig. 2.18.

The volume is a linear function of temperature (°C) with V = 0 at −273.15 °C. On defining temperature in absolute or Kelvin scale as, T (K) = (t °C) + 273.15 … (2.16) Then the plot of volume versus temperature (K) yields Fig. 2.19, in which the volume is directly proportional to the absolute temperature. Charles’ law describes the behaviour of an ideal gas and approximates the behaviour of a real gas. The approximation is very poor at high pressures and low temperatures.

Example 2.2: If 10.3 g sample of a gas occupies 10.3 L at 650 torr and 400 K, what volume will be gas occupies at the same pressure and 25 °C?

Solution: Since n and P are held constant, and V1 = 10.3 L, T1 = 400 K, T2 = 25 °C + 273 = 298 K, V2 =? On substituting given values; V1 T1 = V2 T2 = constant 10.3L 400 K = V2 298 K V = 7.67 L.

Avogadro’s Law: It states that at constant pressure and temperature the volume occupied by a gas is directly proportional to the number of moles of the gas. Mathematically it is expressed as; V = n × constant (when P and T are held fixed) … (2.17) If V1 and V2 are volumes and n1 and n2 are number of moles of gas at constant temperature and pressure, then; V1 n1 = V2 n2 … (2.18) The 1 mole of an ideal gas at 1 atm and 0 °C (Standard Temperature and Pressure, STP) occupies 22.414 L (or dm3 ).

Ideal Gas Law: The ideal gas law relates to the volume and pressure of a gas at a constant temperature. On combining Boyle’s law, Charles’s and Gay-Lussac law and Avogadro’s law we find that the volume of gas depends on pressure, temperature and number of moles of gas in the container.

Summary: Boyle’s Law: V ∝ 1 P (when n and T are held constant) Charles’s Law: V ∝ T (when n and P are held constant) Avogadro’s Law: V ∝ n (when P and T are held constant) Therefore, volume should be proportional to the product of these three terms as; V ∝ 1 P × T × n … (2.19) Replacing proportionality symbol (∝) with equal to symbol (=) and adding the proportionality constant, (R), we get; V = R × 1 P × T × n … (2.20) ∴ PV = n RT … (2.21) where, P represents the pressure of the gas, V stands for the volume of the gas, n represents the number of moles of the gas, R stands for the molar gas constant which is always 0.08205 L atm/K.mol and T represents the temperature of the gas. The equation (2.21) is known as ideal gas equation. As can be understood from the above equation, the pressure and the volume are inversely proportional. As the pressure increases the volume decreases, and as the volume increases the pressure decreases. But the volume and temperature are directly proportional. As the volume increases the temperature also increases. The ideal gas law can be very useful when one needs to find the approximate molecular weight of a gas. The n is replaced by g/M, which is grams of the gas divided by molecular weight.

Applications of the Ideal Gas Law: The ideal gas law PV = nRT has four parameters and a constant, R. This equation can be rearranged to give an expression for each of P, V, n or T. For example, P = nRT/V and P = (nR/V) T. These equations are Boyle’s law and Charles law, respectively. Similar expressions can be derived for V, n and T in terms of other variables. Thus, ideal gas law has many applications; however, it is important to use proper numerical value for the gas constant R as per the units we have for the parameters. Furthermore, n/V is number of moles per unit volume and this quantity has the same units as the concentration. The concentration is a function of pressure and temperature as given in equation below. C = P RT … (2.22)

At 1 atm pressure and room temperature of 298 K the concentration of an ideal gas is 0.041 mol/L. The Avogadro’s law can be further applied to correlate gas density ρ (weight per unit volume or nM/V) and molecular mass, M, of a gas. The following equation is easily derived from the ideal gas law: PM = nM V RT … (2.23) Thus, we have PM = ρ RT       ‡ ρ = nM V ρ = PM RT M = ρ RT P … (2.24)

Example 2.3: An air sample containing only nitrogen and oxygen gases has a density of 1.3390 g/L at STP. Find the weight and mole percentages of nitrogen and oxygen in the sample. Solution: From the density (ρ), we can evaluate an average molecular weight (also called molar mass). PM = ρ RT M = 22.4 × ρ       ‡ RT P = 22.4 L/mol = 22.4 L/mol × 1.3390 g/L = 30.01 g/mol.

Assume that we have 1 mol of gas, and x mol of which is nitrogen, then (1 − x) is the amount of oxygen. The average molar mass is the mole weighted average, and thus, 28.0 x + 32.0 (1 − x) = 30.01 − 4 x = − 1.99 x = 0.497 mol of N2, and 1.0 − 0.497 = 0.503 mol of O2 Now, to determine weight percentages we need to find the amounts of nitrogen and oxygen in 1.0 mol (30 g) of the mixture. Mass of 0.497 mol nitrogen = 0.497 × 28.0 = 13.916 g Mass of 0.503 mol oxygen = 0.503 × 32.0 = 16.096 g Percentage of nitrogen = 100 × (13.916/30) = 46.38 % Percentage of oxygen = 100 × (16.096/30) = 53.65 % = 100 − 46.38 = 53.62 %

Aerosols

Gases can be liquefied by increasing pressure, provided we work below the critical temperature. When the pressure is reduced, the molecules expand and the liquid reverts to a gas. This reversible change of state is the basic principle involved in the preparation of pharmaceutical aerosols. In such products, the drug is dissolved or suspended in a 'propellant', a material that is liquid under the pressure conditions existing inside the container and forms a gas under normal atmospheric conditions. Chlorofluorocarbons and hydro fluorocarbons have traditionally been utilized as propellants in these products because of their physicochemical properties. However, in the face of increasing environmental concerns (ozone depletion) their use is tightly regulated which has led to the increased use of other gases such as nitrogen and carbon dioxide.

Inhalers

The delivery of drugs by inhalation is a critical issue in obstructive airway diseases such as bronchial asthma and chronic obstructive pulmonary disease. The inhaled drugs are targeting the lungs directly and being a lower dose with a quick onset of action, and better therapeutic index. Now day’s inhalers are a major component of patient’s therapeutic management. Several effective molecules have been developed till date but their true effectiveness in real life can be affected and modulated substantially by the device used for inhalation. An increasing number of inhalation devices have been engineered, either for single or combined molecules. However, it was assumed since long ago that the ideal device should be:

Effective: such as, able to consent the inhalation of a sufficient fraction of drug with a particle size ≤ 6 µ, independently of the patient’s inspiratory flow.
Reproducible: such as, able to always consent the inhalation of the same drug amount, also in terms of its respirable fraction.
Precise: such as, able to consent to know at any moment the amount (or the number of doses) of the drug remaining in the device, and whether or not the inhalation was correctly performed: thus the need for providing dry powder inhalers (DPIs) of a “dose counter” and of a “double-dosing protection counter”, in order to avoid a further inhalation if the patient is unaware or not sure of having taken the previous one.
Stable: such as, able to protect the drug(s) contained from the effects of temperature and/or humidity changes.
Comfortable: such as, easy to use in different circumstances (particularly in critical conditions), and possibly containing several doses of the drug(s) for a long-term use.
Versatile: such as, it should consent the use of other drugs by inhalation.
Environmentally compatible, such as not containing chemical contaminants.
Affordable: such as, of acceptable cost, and possibly rechargeable
The DPIs’ family independently of wet nebulizers, pocket devices can be basically grouped in three major classes:

Metered Dose Inhalers (MDIs), still largely used for single and combined molecules, and which need a propellant for the dose delivery.

Dry Powder Inhalers (DPIs), which do not require any propellant, and are increasingly prescribed for single and combined molecules.

Soft Mist Inhalers (SMIs), at present consisting of in only one device for only one molecule (Respimat for Tiotropium bromide).

DPIs are available in wide variety of design and represent a substantial improvement in the inhalation therapy. They fit majority of the above mentioned requirements. Mainly, they eliminate the use of propellants; simplify the inhalation technique; reduce the patient’s co-operation and improve the patient’s compliance to treatment; favor a higher deposition of drugs within the lungs; reduce the variability of the inhaled dose; reduce the incidence of both local and systemic side effects, and finally ameliorate the consistency of the dose and then the outcomes substantially. The most advanced DPIs also fitted the most sophisticated patients’ requirements in terms of minimization of the number of actions needed for preparing the actuation.

In case of a low-resistance DPI, the only driving force for the distribution and the micro dispersion of the drug to inhale is the patient’s inhalation airflow rate which depends on the patient’s airflow limitation and disease severity. The role of the resistance-induced turbulence is obviously negligible in these cases. Therefore, the required regimen of turbulence is achieved only by increasing the inhalation airflow. It frequently represents the main critical limitation for airway obstructive patients. Under these conditions, the variability in the dose consistency is higher and the effective inhaled dose can be far from the original claim. This also is due to the higher oropharyngeal impact of the powdered drug. In correct sense, the “low resistance DPIs” should not be mandatory associated to the concept of “the most effective DPIs” because just in these cases patients are required for a higher inspiratory performance, which frequently cannot be achieved by patients affected by a disease-induced airflow limitation.

Relative Humidity

Relative humidity is the ratio of the partial pressure of water vapour to the equilibrium vapour pressure of water at a given temperature. Relative humidity depends on temperature and the pressure of the system of interest. It requires less water vapour to attain high relative humidity at low temperatures; more water vapour is required to attain high relative humidity in warm or hot air. The relative humidity (RH or φ) of an air–water mixture is defined as the ratio of the partial pressure of water vapour (PH2 O) in the mixture to the equilibrium vapour pressure of water (p * H2 O) over a flat surface of pure water at a given temperature:

Relative humidity is normally expressed as a percentage; a higher percentage means that the air–water mixture is more humid.

Vapour Concentration (Absolute Humidity)

The vapour concentration or absolute humidity of a mixture of water vapour and dry air is defined as the ratio of the mass of water vapour (Mw) to the volume (V) occupied by the mixture. Dv = Mw /V, expressed in grams/m3 or in grains/cu ft. The value of Dv can be derived from the equation PV = n RT. Relative humidity is the ratio of two pressures; %RH = P Ps × 100 … (2.26) where, P is the actual partial pressure of the water vapour present in the ambient and Ps the saturation pressure of water at the temperature of the ambient. Relative humidity sensors are usually calibrated at normal room temperature (well above freezing). Consequently, it generally accepted that this type of sensor indicates relative humidity with respect to water at all temperatures (including below freezing). As already noted ice produces a lower vapour pressure than the liquid water. Therefore, when ice is present, saturation occurs at a relative humidity of less than 100 %. For instance, a humidity reading of 75 % RH at a temperature of −30 °C corresponds to saturation above ice.

Method of Calibration:

A frequent method of calibrating a relative humidity instrument is to place the humidity sensor in a closed container. By putting a known solution of water and another substance inside the container, a known humidity is established at equilibrium. This humidity value is used to provide a reference against which the instrument can be adjusted or calibrated.

Temperature stability:

Obtaining equilibrium conditions is one of the most critical requirements of the method. This means that there should be no difference of temperature between the humidity sensor, the solution and the head space above the solution. Unstable temperature during calibration will not permit this. A temperature stability of 0.02°C/min or better is required during the calibration process for the method to be accurate.

Temperature of calibration:

The relative humidity values generated by the different solutions used for the purpose of calibration are affected by temperature. Therefore, a correction must be made for the temperature of calibration. However, no correction is required for the effect of temperature on the total pressure inside the calibration container. The temperature of calibration may also be restricted by the design of the instrument. For instance, an instrument that provides a compensation for the effect of temperature on the humidity sensor does so by assuming that the temperature of calibration is always the same. In that case, the manufacturer provides a recommendation as to the range of calibration temperature that result in the best overall accuracy for the instrument.

Significance of RH:

Climate control: Climate control refers to the control of temperature and relative humidity in buildings, vehicles and other enclosed spaces for the purpose of providing for human comfort, health and safety, and of meeting environmental requirements of machines, sensitive materials (for example, labile pharmaceuticals) and technical processes.

Human discomfort:

Humans are sensitive to high humidity because the human body uses evaporative cooling, enabled by perspiration, as the primary mechanism to get rid of waste heat. Perspiration evaporates from the skin more slowly under humid conditions than under arid. Because humans perceive a low rate of heat transfer from the body to be equivalent to a higher air temperature, the body experiences greater distress of waste heat burden at high humidity than at lower humidity, given equal temperatures. For example, if the air temperature is 24 °C (75 °F) and the relative humidity is zero percent, then the air temperature feels like 21 °C (69 °F). If the relative humidity is 100% at the same air temperature, then it feels like 27 °C (80 °F). In other words, if the air is 24 °C (75 °F) and contains saturated water vapour, then the human body cools itself at the same rate as it would if it were 27 °C (80 °F) and at 20% relative humidity (an unstated baseline used in the heat index). The heat index and the humidex are indices that reflect the combined effect of temperature and humidity on the cooling effect of the atmosphere on the human body. In cold climates, the outdoors temperature causes lower capacity for water vapour to flow about. Thus, although it may be snowing and at high humidity relative to its temperature outdoors, once that air comes into a building and heats up, its new relative humidity is very low, making the air very dry, which can cause discomfort and can lead to ill health, although, dry air is good for those suffering from some lung disorder.

Effect on skin:

Low humidity causes tissue lining nasal passages to dry, crack and become more susceptible to penetration of Rhinovirus cold viruses. Low humidity is a common cause of nosebleeds. The use of a humidifier in homes, especially bedrooms, can help with these symptoms. Indoor relative humilities should be kept above 30% to reduce the likelihood of the occupant's nasal passages drying out. Humans can be comfortable within a wide range of humilities depending on the temperature from 30% to 70% but ideally between 50% and 60%. Very low humidity can create discomfort, respiratory problems, and aggravate allergies in some individuals. In the winter, it is advisable to maintain relative humidity at 30 percent or above. Extremely low (below 20%) relative humilities may also cause eye irritation.

Buildings:

For climate control in buildings using HVAC systems, the key is to maintain the RH at a comfortable range low enough to be comfortable but high enough to avoid problems associated with very dry air. When the temperature is high and the relative humidity is low, evaporation of water is rapid; soil dries, wet clothes hung on a line or rack dry quickly, and perspiration readily evaporates from the skin. Wooden furniture can shrink, causing the paint that covers these surfaces to fracture. When the temperature is low and the relative humidity is high, evaporation of water is slow. When relative humidity approaches 100%, condensation can occur on surfaces, leading to problems such as mold growth, corrosion, decay, and other moisture-related deterioration. Condensation can pose a safety risk as it can promote the growth of mold and wood rot as well as possibly freezing emergency exits shut. Certain production and technical processes and treatments in factories, laboratories, hospitals, and other facilities require specific relative humidity levels to be maintained using humidifiers, dehumidifiers and associated control systems.

Water vapour is independent of air:

The notion of air "holding" water vapour or being "saturated" by it is often mentioned in connection with the concept of relative humidity. This, however, is misleading because the amount of water vapour that enters a given space at a given temperature is independent of the amount of air that is present. Indeed, a vacuum has the same equilibrium capacity to hold water vapour as the same volume filled with air; both are given by the equilibrium vapour pressure of water at the given temperature.

Pressure dependence:

The relative humidity of an air–water system is dependent not only on the temperature but also on the absolute pressure of the system of interest. This dependence is demonstrated by considering the air–water system shown. The system is closed (i.e., no matter enters or leaves the system).

If the system at State A is isobarically heated (constant pressure) the RH of the system decreases because the equilibrium vapour pressure of water increases with increasing temperature. This is shown in State B. If the system at State A is isothermally compressed (constant temperature) the relative humidity of the system increases because the partial pressure of water in the system increases with the volume reduction. This is shown in State C. At above 202.64 kPa the RH would exceed 100% and water may begin to condense. If the pressure of State A is changed by simply adding more dry air, without changing the volume, the relative humidity would not change. Therefore, a change in relative humidity can be explained by a change in system temperature, a change in the volume of the system, or change in both of these system properties.

Enhancement factor:

The enhancement factor (fw) is defined as the ratio of the saturated vapour pressure of water in moist air (e ' w ) to the saturated vapour pressure of pure water (e* w ). fw = e ' w / e* w … (2.27) The enhancement factor is equal to unity for ideal gas systems. However, in real systems the interaction effects between gas molecules result in a small increase of the equilibrium vapour pressure of water in air relative to equilibrium vapour pressure of pure water vapour. Therefore, the enhancement factor is normally slightly greater than unity for real systems. The enhancement factor is commonly used to correct the equilibrium vapour pressure of water vapour when empirical relationships, such as those developed by Wexler, Goff, and Gratch, are used to estimate the properties of psychrometric systems. Buck has reported that, at sea level, the vapour pressure of water in saturated moist air amounts to an increase of approximately 0.5% over the equilibrium vapour pressure of pure water. The term relative humidity is reserved for systems of water vapour in air. The term relative saturation is used to describe the analogous property for systems consisting of a condensable phase other than water in a non-condensable phase other than air.

Measurement:

A device used to measure humidity is called a hygrometer; one used to regulate it is called a humidistat, or sometimes hygrostat. The humidity of an air–water vapour mixture is determined through the use of psychrometric charts if both the dry bulb temperature (T) and the wet bulb temperature (Tw) of the mixture are known. These quantities are readily estimated by using a sling psychrometer. There are several empirical formulas that can be used to estimate the equilibrium vapour pressure of water vapour as a function of temperature. The Antoine equation is among the least complex of these, having only three parameters (A, B, and C). Other formulas, such as the Goff-Gratch equation and the MagnusTetens approximation, are more complicated but yield better accuracy. The formula presented by Buck is commonly encountered in literature: e* w = (1.0007 + 3.46 x 10−6P) × (6.1121)e(17.502T/240.97+T) … (2.28) where T is the dry bulb temperature expressed in °C, P is the absolute pressure expressed in millibars, and e* w is the equilibrium vapour pressure expressed in millibars. Buck has reported that the maximum relative error is less than 0.20% between −20 °C and +50 °C when this particular form of the generalized formula is used to estimate the equilibrium vapour pressure of water.

Liquid Complexes

Liquid complexes are binary mixtures that have coexistence between two phases: solid–liquid (suspensions or solutions of macromolecules such as polymers), solid–gas (granular), liquid–gas (foams) or liquid–liquid (emulsions). They exhibit unusual mechanical responses to applied stress or strain due to the geometrical constraints that the phase coexistence imposes. The mechanical response includes transitions between solid-like and fluid-like behavior as well as fluctuations. Their mechanical properties can be attributed to characteristics such as high disorder, caging, and clustering on multiple length scales.

Complex systems are distinguished by their behaviour as determined by competing processes of self-organization (ordering) and self disorganization (disordering) creating a hierarchical adaptive structure. A notion of complexity is also used in amorphous materials exhibiting slow and non-exponential relaxation, in particular in glass-forming liquids and glasses. However in liquid complexes, complexity is not yet a quantifiable but rather a qualitative characteristic. Numerous experimental and theoretical studies and, more recently, computer simulations revealed important macro-and mesoscopic details associated with materials complexity such as dramatic slowing-down of structure changes on cooling, wide spectrum of relaxation times and stretched-exponential (KWW) relaxation kinetics and dynamic heterogeneity on microscopic length-scales. These features and the sometime observed power law correlations are often used as practical but rather qualitative criteria of complexity in materials. In the Literature, the assumed physical cause of materials complexity is the dynamic competition between aggregation of particles into preferred structures, and factors preventing crystallization. Understanding the origins of complexity and the dynamics of structure in complex materials is most important but hardest problems in condensed matter.

Not every liquid becomes complex on cooling. Three-dimensional (3D) liquids with simple two-particle interactions (molten metal’s and salts, liquefied noble gases, Morse particles) aggressively crystallize on cooling before they show any significant signs of complexity. Classical 3D complex liquids have complicated and competing interactions and special supercooling regimes are necessary to avoid crystallization on supercooling. Two-dimensional (2D) liquids with simple interactions have a continuous or almost continuous crossover from simple liquid state to crystal. At crossover temperatures, Fig. 2.21, particles in these equilibrium liquids aggregate to form a dynamic mosaic of crystalline-ordered regions (crystallites) and less-ordered clusters. At the high-temperature end of the mosaic states, crystallites are small and separated island of order in a disordered (amorphous) matrix. Crystallites fraction of the system increases at lower temperatures where crystallinity percolates. At even lower temperatures, crystallites merge into a multiconnected crystalline matrix with expected algebraic decay of orientation order (hexatic liquid) or long range order. The mosaic is a feature observed at temperatures where the correlation length for orientations is finite and the 2D liquid is in normal (not hexatic) state.

Liquid Crystals

The three distinct states of matter as solid, liquid, and gas have been discussed so far. However, there is a state of matter, which does not meet the necessary requirements of any of these three categories. For example, a substance like cholesterol or mayonnaise is somewhere between a liquid and a solid. This is not quite liquid or quite solid, but is a phase of matter whose order is intermediate between that of a liquid and crystal. It is often called a mesomorphic state which is state of matter in which the degree of molecular order is intermediate between the perfect three dimensional, long-range positional and orientational order found in solid crystals and the absence of long-range order found in isotropic liquids, gases, and amorphous solids. It is also called as meso intermediate. Physically, they are observed to flow like liquids showing some properties of crystalline solids. Hence this state is considered to be the next (fourth) state of matter known as liquid crystal (LC) state. The LC state is also known as mesophase and can be defined as the condensed matter that exhibit intermediate thermodynamic phase between the crystalline solid and simple liquid state. LC’s can be considered to be crystals, which have lost some or all of their positional order while maintaining full orientational order. They are free to move, but like to line up in about the same direction. The degree of mobility of the molecules in the LC’s is less than that of a liquid.

The liquid crystals are of thermotropic and lyotropic types. The lyotropic liquid crystals are induced by the presence of solvent. Thermotropic liquid crystals are induced by a change in temperature and are essentially free of solvent. Liquid crystals are liquids featuring a certain level of orientational order. Specifically, molecules in LCs tend to point to a certain direction, while they still have translational (positional) freedom. Although they are best known for their application in displays, liquid crystals are also an essential part of all life forms. Lyotropic liquid crystals are essential organic substances, DNA, lipids of cellular membranes and proteins are some examples of well-known liquid crystals. In liquid crystals drug delivery crystalline solids exhibit short as well as long-range order with regard to both position and orientation of the molecules. Whereas liquids are amorphous in general but may show short-range order with regard to position and/or orientation. Liquid crystals show at least orientational long-range order and may show short-range order, whereas positional long-range order disappears.

The LC state is widespread in nature such as lipoidal forms found in nerves, brain tissue and blood vessels. LC’s may also be associated with arthrosclerosis and formation of gallstones. They are believed to have structures similar to those of cell membranes. In general, most molecules that form a liquid crystalline state are organic, elongated, rectilinear, rigid, and found to have strong dipoles and easily polarizable groups. The existence of liquid crystalline state may be because of heating of solids or from the action of certain solvents on solids. Cholesterol acetate, a liquid, which exhibit optical properties, is first of its kind known LC’s. Since, this state of matter possesses orientational or weak positional order; they display some physical properties of crystals but flow like liquids. When transition between the phases is temperature dependent, they are called thermotropic and when transitions are dependent of different components these LC’s are called lyotropic.

Thermotropies are mostly used in technical applications, while lytropics are important for biological systems such as membranes. Liquid crystals due to anisotropic intermolecular forces usually consist of steric rod or disc like organic molecules aligning themselves with long-range order. There are three types of liquid crystals.

Types of Liquid Crystals:

Nematic Crystal:

In the simple liquid crystalline state the molecules possess only orientational but no positional order are called nematic crystal phase. In the nematic phase the molecules can rotate about one axis (i.e. uniaxial) and are mobile in three directions. They are polarizable thread or rod like organic molecules on the order of 25 ° A in lengths and 5 ° A in height. The order of nematic crystal is a function of temperature.

The name nematic has been given with respect to thread-like textures as observed under polarizing microscope. A unit vector called nematic director can describe the direction of considered alignment. Because of their tendency to organize themselves in a parallel fashion they demonstrate interesting and useful optical properties. As nematics are characterized by orientational order of the constituent molecules, the molecular orientation and hence the material’s optical properties, can be controlled with applied electric fields.

Cholesteric crystal:

LC’s when made of chiral (asymmetric) molecules that differ from their mirror image a cholesteric liquid crystal e.g., cholesterol acetate, is obtained. Cholesteric can be similar to nematic but differ in the considered orientation that it forms a helical structure with the helical axis perpendicular to the director.

Physical Properties:

Physical properties of liquid crystals are anisotropic due to orientational order. These properties are the heat of diffusion, the magnetic susceptibility, the dielectric permittivity or the optical birefringence. Liquid crystals are sensitive to electrical fields, a property that has been used in display systems. Liquid crystals are mobile and found to show flow properties of liquids like rotational viscosity acting on dynamic director deformations, respectively.

Pharmaceuticals and Cosmetic Applications of LCs:

Liquid Crystal Emulsion:

A large part of cosmetic products are made in the form of emulsions, a form that allows the simultaneous use of lipophilic and hydrophilic ingredients in the required dosages. A product in the form of an emulsion also has the advantage of having the most convenient appearance and texture that also facilitates its application. They can be formulated to be liquid, milk type emulsions of variable consistency, creams, or even super liquid sprayable emulsions. It is well known fact that an emulsion is the best carrier for active ingredients and functional substances. The theory of stabilizing an emulsion through the formation of a network of liquid crystals is different than the HLB theory. The gasification of the water phase obtainable with hydrosol vatable polymers or with emulsifiers that are able to form a reticular organised structure in liquid crystal form, eliminates the need to use waxy components in large quantities and consistency factors that are no longer in harmony with the modern conception of light and easy to spread emulsions. LCs (mesophases) provides the following advantages to emulsion

Stability: Emulsion stability of the multilayers around the oil droplets act as a barrier to coalescence. If oil droplets coalesce emulsion breaks. This barrier for coalescence acts as increased stability property of the emulsion.
Prolonged hydration: Lamellar liquid crystalline and gel network contain water layer, which shows that 50% of the water of oil in water (o/w) emulsion can be bound to such structures. Such water is less prone to evaporation when applied to the skin and permits a long lasting moisturization / hydrating effect, necessary for drug entry.
Controlled Drug delivery: Liquid crystals prevent the fast release of the drug dissolved in the oil phase of an emulsion. This is attributed to the lamellar liquid crystalline multilayer, which reduces the interfacial transport of a drug dissolved within the oil droplets. Microscopic observations under polarized light show the exceptional thickness of liquid crystalline lamellar layer around the oil droplets.

Function and Properties of LCs Emulsion System

LCs, when present at the oil/water interface, the liquid crystals help to give the system rigidity and, by limiting the fluctuation of the components at the interface give great stability to the emulsion. Furthermore, the liquid crystal system enhances the moisturizing ability of the emulsion. The quantity of inter-lamellar water can be extremely high and become immediately available when the cream is applied to the skin. For these reasons these emulsions have a shinny surface, a fresh and original feel and they leave a light and pleasant sensation on the skin. In recent years, the moisturizing effect of creams and lotions has become increasingly more important and cosmetic chemists are constantly searching for better methods of retaining water in the superior layers of the skin. The evaporation of the bonding water in emulsions containing anisotropic lamellar phases is slower and permits a hydro retentive action that prolongs the moisturizing effect. The associations that are formed because of the excess water are particularly interesting; in these cases the ability of the crystalline phase to swell is strictly linked to the stability and the behaviour of the emulsion because, in a liquid crystal system, the quantity of inter-lamellar water and of hydrophile elements can amount to 70% of the total external phase.

Controlled Release of Bioactive Materials:

The release of the active substance from liquid crystalline delivery systems is often controlled by diffusion, and some systems using the photo induced or thermal phase transition of the liquid crystals as the release for bioactive materials.

Drug Loading:

According to the nature of the drug, it can be added in both the aqueous as well as oil phase. Loading totally depends on solubility of active constituents and their partition between existing phases. For example, cefazolin, cefuroxime, clomethiazole, clindamycin phosphate, 4-phenylbutylamine, prilocaine, oestriol, isosorbide mononitrate, insulin, indomethacin, clotrimazole, gramicidin, nitroglycerin, lidocaine hydrochloride etc.

Other Applications:

Lyotropic LC’s include organic substances that are essential for life. Examples of lyotropic liquid crystals include DNA, proteins, cholesterol etc. LC pharmaceuticals are a unique class of lyotropic LC’s that represent novel drug candidates for the treatment of a wide range of diseases. LCs are useful in cosmetic and pharmaceutical compositions as well as methods comprising delivery systems for the controlled release and enhanced penetration of biologically active materials (for example, vitamin A) to the skin. The delivery systems comprise cholesteric liquid crystals wherein the active material is retained within the lamellar molecular structure (i.e., between the molecular sheets) of the cholesteric LC. Another example of LC is new investigational antitumor drug called Telecine™, a compound that also has antiviral and antibacterial applications. LCs are also used in solubilization of water insoluble substances. LC’s have its applications in most areas due to its remarkable features of anisotropic optical properties. As a result of strong Bragg’s reflection of light cholesteric LC’s have vivid iridescent colours. In some of the LC’s the pitch of spiral and reflected colour changes with temperature therefore can be used to measure temperature of the skin and other surfaces. This can be useful in detecting elevated temperatures under the skin as in certain disease states.

Glassy State

The rapid cooling of a liquid below its melting point (Tm) leads to an amorphous state with structural characteristics of a liquid but with a much greater viscosity.

The enthalpy and volume changes immediately below Tm exhibit no discontinuity with those observed above Tm, so the amorphous state is considered to be equilibrium super cooled state. The amorphous state is also called as rubbery state because of the macroscopic properties of amorphous solids in this region. The 3-D long-range order that normally exists in a crystalline material does not exist in the amorphous state and position of molecule relative to another molecule is more random as in the liquid state. Therefore they are considered as super cooled liquids. Typically, an amorphous solid exhibit short-range order over a few molecular dimensions and has physical properties quite different than the crystalline solids.

Amorphous state can also be characterized by rate and extent of molecular motions. The molecular motions in the supercooled liquids are usually less than 100 s and viscosity is between 10−3 to 1012 Pa.s and both properties are strongly temperature dependent. Further cooling of supercooled liquid reduces molecular mobility of a liquid to a point where material is unable to attain equilibrium in time scale as it loses its thermal energy leading to change in temperature dependence of the enthalpy and volume. The temperature at which this occurs is called as glass transition temperature (Tg). Below Tg the material is kinetically frozen into thermodynamically unstable glassy state with respect to equilibrium liquid and crystalline state. At Tg physical properties like hardness, volume, and percent elongation-to-break and young’s modulus undergo change.

Melting is observed in case of crystalline solids, while the glass transition happens only to amorphous solids. A given sample may often have both amorphous and crystalline domains within it, so the same sample can show a Tm and a Tg. But the chains that melt are not the chains that undergo the glass transition. When crystalline solids are heated at a constant rate, the temperature increases at a steadily.

Solids

The state, in which a substance has no tendency to flow under stress, resists forces that tend to deform it, and remain in definite size and shape is called as solid state. In solid state the molecules are closely bound to one another. A solid hold its shape. The volume of solid is fixed by the shape of solid. There are two types of solids namely; crystalline solids and amorphous solids. They differ from one another by the way their particles are arranged and their melting points.

Physicochemical Properties of Drug Molecules: Determinations and Application

The molecular structure of the compound uniquely defines all its physical, chemical and biological properties. It is generally recognized that physicochemical properties play an important role in product development including studies on biological performance of drugs. A study of the physical properties of drug molecules is a prerequisite for product preformulation, formulation development and optimizing storage and usage conditions. It often leads to a better understanding of the relationship between molecular structure and drug action. The most important physical properties related to product formulation and biological performance is summarized below:

Classification:

Physical properties of substances may be classified in to three types;

Additive Properties:

Additive properties are derived from sum of the properties of individual properties of atoms or functional groups present within the molecule. The examples of this type are mass or molecular weight, volume etc. Consider the case of acetic acid (CH3COOH). Obtaining molecular weight of acetic acid involves addition of molecular weights of individual atoms that makes it. Acetic acid contains; C = 2, H = 4 and O = 2. So the molecular weight is calculated as; Molecular weight of acetic acid = C × 2 + H × 4 + O × 2 = 12 × 2 + 1 × 4 + 16 × 2 = 60 g/mol.

Constitutive Properties:

These properties are depending on the structural arrangement of atoms and functional groups as well as bond structure that exists within the molecules. The examples of this type are optical activity, surface tension, viscosity etc. Consider the case of lactic acid. It exists in two forms namely d-lactic acid and l-lactic acid. The specific rotation of d-lactic acid is +3.8° while l-lactic acid shows it as -3.8°.

Combined Additive-Constitutive Properties:

Many physical properties are constitutive and yet to have some measure of additivity is called as additive-constitutive properties. The example of this type is molar refraction.

The standard contributions of atoms, groups and structural unit’s to the molar refractions are listed. Molar refractions of ethyl methyl ketone and 2, 3-butanol is obtained as follows.

The molar refraction of both these compounds is obtained as sum by substituting values of their contributions for groups and bond structures as given in Table 2.10. According to definition of additive property, molar refractions of both these molecules are sum of atoms and groups that makes these molecules. Therefore molar refraction is an additive property. Although the number and types of atoms in both these molecules are same their arrangements in molecules is different and they show different molar refraction values as ethyl methyl ketone has 20 while 2, 3 - butanol has 18.60. Therefore, molar refraction is additive as well as constitutive property.

Colligative Properties:

Colligative properties are defined as the properties which depend upon the total number of non-volatile solute particles present in the solution. Dilute solutions which contain negligibly small amount of non-volatile solute exhibit colligative properties. The examples of these properties are lowering of vapour pressure, freezing point depression, boiling point elevation and osmotic pressure. These properties are used to determine molecular weights of compounds. In this chapter various physical properties discussed are refractivity, optical activity, dielectric constant and induced polarization, dipole moment and dissociation constant, magnetic properties, molecular and Raman spectra, nuclear magnetic resonance and x-ray diffraction.

Refractive Index

In 1621, a Dutch physicist named Willebrord Snell derived the relationship between the different angles of light as it passes from one transparent medium to another. When light passes from one transparent medium to another, it bends according to Snell’s law which states: Νi × Sin (Ai) − Nr × Sin (Ar) … (2.32) where, Ni is the refractive index of the medium the light is leaving, Ai is the incident angle between the light ray and the normal to the medium to medium interface, Nr is the refractive index of the medium the light is entering; Ar is the refractive angle between the light ray and the normal to the medium to medium interface. In other words refractive index of substance is the ratio of velocity of light in vacuum or air to that in the substance When monochromatic light passes through a less dense medium such as air or vacuum and enters a denser medium, the advancing waves at interface are modified and brought closer together, This leads to decrease in speed and shortening of wavelength. When light passes the denser medium, a part of wave slows down more quickly as it passes through interface and makes it bend towards the interface. This phenomenon is called as refraction. If light passes from denser medium to less denser medium then it is refracted away from the interface. This effect observed between mediums is expressed as refractive index (n). The refractive index is a constant for a given pair of materials under specified conditions. It can be defined as ratio of speed of light in material 1 to the speed of light in material 2. This is usually written 1n2 and is the refractive index of material 2 relative to material 1. The incident light is in material 1 and the refracted light is in material 2. When the incident light is in a vacuum this value is called the absolute refractive index of material. Refractive indices of most substance are more than air because the velocity of light in air is greater than in the substance for example, absolute refractive index of water is 1.330, soda lime glass 1.510. By definition the refractive index of a vacuum is 1. In practice, air makes little difference to the refraction of light with an absolute refractive index of 1.0008. The refractive indices of some liquids are given.

Bending light:

The bending of light rays is due to the refraction. As light passes from one transparent medium to another, it changes speed, and bends. How much this happens depends on the refractive index of the mediums and the angle between the light ray and the line perpendicular (normal) to the surface separating the two mediums, Fig. 2.43. Each medium has a different refractive index. The angle between the light ray and the normal as it leaves a medium is called the angle of incidence. The angle between the light ray and the normal as it enters a medium is called the angle of refraction.

Critical angle:

If the angle of incidence is increased there is an increase in angle of refraction. The maximum angle of incidence that can be achieved is 90o . sin i = sin 90 sin r = 1 sin r … (2.34) The r in equation (2.34) is called as critical angle. Refraction simulator is used to know how light bends toward the normal when the light enters a medium of greater refractive index, and away from the normal when entering a medium of lesser refractive index. When the light is moved to an angle close to 90° or − 90° in the medium with a higher refractive index we approach the critical angle and the refracted light approaches 90° or − 90°. At critical angle the angle of refractions becomes 90° or − 90° and the light is no longer transmitted across the medium1/medium2 interface. For angle with greater in absolute value than the critical angle, all the light is reflected. This is called total reflection.

Specific Refraction:

Initially, it was not possible to draw any conclusion regarding the nature of the substance from the refractive index. In 1880, scientist Lorentz showed the property specific refraction which was found to be more useful in characterization of substance independent of temperature. The specific refraction is mathematically expressed as RS = (n2 − 1) (n2 + 2) × 1 ρ mL/g … (2.35) where, RS is specific refraction in mL/g, n is refractive index of substance and ρ is density of substance at the temperature at which refractive index is determined.

Molar Refraction:

Molar refraction is defined as molecular weight times the specific refraction of substance. It is more useful property than specific refraction as it is characteristic of the substance and useful in structural studies like finding nature of bonding in molecules and in determination of dipole moment. It is expressed as RM = (n2 − 1) (n2 + 2) × M ρ mL/mol … (2.36) where, RM is molar refraction in mL/mol and M is molecular weight of the substance under study. The experimentally determined values of molar refractions are compared with the theoretical values giving their contribution of atoms, groups and structural units to the molar refraction.

Measurement of Refractive Index:

Refractive index is determined by using instrument called refractometer. Abbes refractometer, immersion refractometer and Pulfrich refractometer are used for this purpose. Abbes refractometer is commonly used at laboratory scale because of its advantages over other refractometers. It is most convenient, reliable and simple instrument with small sample size requirement suitable for range of substances. Ordinary light source, easy maintenance and economy and easy determinations are some of the other advantages of this instrument. The components of Abbes refractometers include light reflection mirror, dispersion compensator, telescope, and index arm and prism box. The schematic of abbes refractometer is shown in Abbes refractometer may be calibrated with anyone of the liquid specified in Table 2.12 at temperatures below 25 °C using D-line of sodium.

Applications:

Since refractive index is a fundamental physical property of a substance it is often used to analyze and identify a particular substance, confirm its purity, or measure its concentration. Refractive index values are useful in determination of molecular weights and structures of organic compounds from their molar refraction values. Refractive index is used to measure refraction characteristics of solids, liquids, and gases. Most commonly it is used to measure the concentration of a solute in an aqueous solution. For a solution of sugar, the refractive index can be used to determine the sugar content. Similarly alcohol content in bioproduction is also determined from the refractometry. Dielectric constant and molar polarizability values can be obtained from the refractive index. Refractive index of a material is the most important property of any optical system that uses refraction for example, lenses and prisms.

Optical Rotation

Ordinary light consists of vibrations, which are evenly distributed in all directions in a plane perpendicular to the direction of propagation, called as unpolarized light. When the vibrations of light are restricted to only one plane, the light is said to be polarized light. Some substances rotate the plane of polarized light are called as optically active substances. This property of optically active substance is measured as angle of rotation. The property in which rotation of plane polarized light is observed is known as optical activity. Optically active substances include organic molecules with a central carbon atom to which four different groups are attached, making the molecule very asymmetric (chiral carbon).

The use of one name for glucose, dextrose, refers to the fact that it causes linearly polarized light to rotate to the right or dexter side. Those optically active substances that rotate plane polarized light to the left or counterclockwise are known as leave rotatory (l) or (−) substance. Laevulose, more commonly known as fructose causes the plane of polarization to the left. Fructose is even more strongly leavo rotatory than the glucose. Other examples of optically active substances are lactic acid, tartaric acid, 2-methyl -1-butanol etc. Optical rotation occurs because of optically active substances have different refractive indices for left and right polarized light. Another way to make this statement is that left and right polarized light travel through an optically active substance at different velocities. Optical activity is considered to be due to the interaction of plane polarized radiations with electrons in molecules which shows electronic polarization. This interaction rotates the direction of vibration of radiation by altering electric field.

Optically active substances can be categorized in to two types:

Those which are optically active only in the crystal state due to their characteristic crystal structure and becomes optically inactive in the fused or dissolved state, for example, sodium chlorate, quartz crystal etc and,

Those which shows optical activity in all states viz. Crystalline (solid), fused (liquid) and gaseous state, due to their structural configurations.

Specific Rotation:

When a polarized light passes through an optically active substance, all the molecules in the path of light rotates plane of polarization by some constant amount which is a characteristic of that substance. The total rotation in the emergent light beam is proportional to the path length (l) and density (ρ) of the substance. This relation is mathematically expressed as θ ∝ ρ … (2.37) ∴ θ = [α] ρ … (2.38) where, α, is proportionality constant called as specific rotation. The amount of optical rotation depends on the number optically active species through which the light passes and thus depends on both the sample path length and analyte concentration. Specific rotation provides a normalize quantity to correct for this dependence, and is defined as; [α] T λ = θ l × ρ … (2.39) where, θ, is measured optical angle of rotation in deg cm2 /g or degrees, l is sample path length in decimeters (dm) and ρ is density if the substance is pure liquid, λ is the wavelength of light used for observation, usually 589 nm, the D line of a sodium lamp unless otherwise specified and T is the temperature in °C.

Solution: Substituting the given values in equation [α] T λ = θ l × ρ [α] T λ = 104 l × 0.805 = 129 .

Molar Rotation: Molar rotation is characteristic property of optically active substances. It is obtained from multiplication of specific rotation and molecular weight of the compound as µ = M[α] × 100 … (2.41) where, µ is molar rotation, M is molecular weight and [α] is specific rotation.

Enantiomeric Purity: The molecules that are non-superimposable mirror images are called enantiomers. In case of optically active substances if only one enantiomer is present then the substance is considered to be optically pure, while if it consists of mixtures of two enantiomers (a racemic mixture), it will not rotate plane of polarized light and is optically inactive. A mixture that contains one enantiomer in excess displays a net plane of polarization which is characteristic of the enantiomer that is in excess. The optical purity or the enantiomeric excess (%ee) of a sample can be determined as follows: Optical purity = % enantiomeric excess = % enantiomer1 − % enantiomer2 = 10 × [a] of mixture [a] of pure sample %e = 100 ([R] − [S]) ([R] + [S]) … (2.42) where, [R] and [S] are the concentrations of R and S isomers, respectively.

Measurement of Optical Activity:

Measurement of orientation of plane polarized light is called polarimetry, and the instrument used is called a polarimeter. The simplest polarimeter, Fig. 2.46, consists of monochromatic light source, a polarizer, a sample cell, a second polarizer which is called the analyzer and a light detector. Polariser and analyzer are made up of Nicol prisms. When analyzer is oriented 90° to the polarizer no light reaches to the detector. The polarizer is placed near to the light source while analyzer is placed between sample cell and the detector. The sample cell of suitable size and capacity with outward projection at the centre, to trap the air bubble is usually used. When an optically active substance is placed in the sample cell and beam of light is passed through, it rotates the polarization of the light reaching the analyzer so that there is a component that reaches the detector. The angle that the analyzer must be rotated from the original position is the optical rotation.

Dielectric Constant

A polar molecule can sustain a separation of electric charge either through the induction by an external electric field or by a permanent charge separation within a molecule. The separation of charge can be best understood from the concept called dielectric constant. Consider the example of parallel plate condenser.

The parallel plates are separated by some medium across a distance r and connected to voltage supply source. The electricity will flow across the plates from left to right through the battery until potential difference of the plates equals that of the battery which is supplying the initial potential difference. The capacitance, C, is equal to the amount of electric charge, q, stored on the plates, divided by V, the potential difference, between the plates. C = q V … (2.43)

The capacitance of condenser depends on the type of thickness of the condenser separating the plates. The Co is used as capacitance reference medium on which to compare other mediums. The Co is the capacitance between the plates when a vacuum fills the space between the plates. The ratio of capacitance of test material (Cx) divided by the capacitance of reference material is termed as dielectric constant.

If the polar molecules are placed between plates of charged capacitor, the molecules can undergo an induced polarization. This occurs because of the separation of the electric charge within the molecules as it is placed in the electric field between the plates. This polarization is usually temporary and is independent on the ease with which the molecules can be polarized. This temporary induced polarization is proportional to field strength of capacitor and induced polarizability, αp, which is characteristic property of the particular molecules. The ease with which a molecule is polarized by any external force (electric field, light or any other molecule) is known as polarizability. The dipole moment and polarizability of some solvents...

Dipole Moment

Dipole is a pair of separated opposite electric charges. Electric dipole is an assemblage of atoms or subatomic particles having equal electric charges of opposite sign separated by a finite distance. Dipoles are characterized by their dipole moment, a vector quantity with a magnitude equal to the product of charge or magnetic strength of one of the poles and the distance separating the two poles.

where, µ is dipole moment, q is charge on atom and r is distance of separation of charge. The direction of the dipole moment corresponds for electric dipoles, to the direction from the negative to the positive charge. The direction of an electric field is defined as the direction of the force on a positive charge, electric field lines away from a positive charge and toward a negative charge.

Molecular Dipoles:

Many molecules have dipole moments due to non-uniform distributions of positive and negative charges on its various atoms. In the case of HCl, the bonding electron pair is not shared equally rather is attracted towards the more electronegative chlorine atom due to its higher electro-negativity which pulls the electrons towards it. It leads to development of positive charge to H atom and negative charge to chlorine atom. A molecule having positive and negative charges at either terminal is referred as electric dipoles or just dipole. Dipole moments are often stated in Debyes; The SI unit is the coulomb meter.

Molecular dipoles are of three types.

Permanent dipoles: These occur when two atoms in a molecule have substantially different electro-negativity with one atom attracting electrons more than another becoming more electronegative, while another atom becomes more electropositive.

Instantaneous dipoles: These occur due to chance when electrons happen to be more concentrated in one place than another in a molecule, creating a temporary dipole. 3. Induced dipole: These occur when one molecule with a permanent dipole repels another molecule’s electrons, inducing a dipole moment in that molecule. In a diatomic molecule, the dipole moment is a measure of the polar nature of the bond; i.e. the extent to which the average electron charges is displaced towards one atom. In a polyatomic molecule, the dipole moment is the vector sum of the dipole moments of the individual bonds. In a symmetrical molecule, such as tetrafluoromethane (CF4), there is no overall dipole momenta though the individual C-F bonds are polar.

Molecular dipole moments: In most molecules even though the total charge is zero, the nature of chemical bond is such that the positive and negative charges do not overlap. These molecules are said to be polar because they possess a permanent dipole moment. The example of this type is water molecule. The molecules with mirror symmetry like oxygen, nitrogen carbon dioxide and carbon tetrachloride have no permanent dipole moments. Even if there is no permanent dipole moment, it is possible to induce a dipole moment by the application of an external electric field and is called as polarization. The magnitude of the dipole moment induced in the molecules is a measure of the polarizability of that molecular species.

Permanent dipole moment: The permanent dipole moment differs from induced polarization in the sense that it is a permanent separation and it happens only to be in the polar but not in the non-polar molecules. These charges that separate balance out each other and therefore have a net charge of zero. The water is example of permanent dipole moment. The permanent dipole moment is defined as the vector sum of the individual charge moments within the molecules.

Applications: The structure of the molecule can be confirmed from the dipole moment values, for example, chlorobenzene, benzene, carbon dioxide etc. The cis and trans isomers can be differentiated form dipole moment values, for example, cis and trans dichloroethylene. Dipole moments can be used to determine percent ionic characteristic of bond of the molecule, e.g. H-Cl a covalent bond, ionic characteristic is 17%. Permanent dipole moments can be correlated with the biological activities to obtain information about the physical parameters of molecules. The more soluble the molecule the easier it passes the lipoidal membrane of insects and attacks the insect’s nervous system. Therefore, the lower is the dipole moment the greater is the insecticidal action. For example, p, m and o isomers of DDT show different insecticidal activities due to their differences in permanent dipole moment as p- isomer shows µ=1.1 and has predominant toxicity, o-isomer shows µ = 1.5 with intermediate toxicity while m-isomer shows µ = 1.9 with least toxicity. The variations in activities of different isomers are due to the greater solubilities in non-polar solvents.

Dissociation Constant

Dissociation is the process by which a chemical compound breaks-up into simpler constituents as a result of either added energy (dissociation by heat), or the effect of a solvent on a dissolved polar compound (electrolytic dissociation). It may occur in the gaseous, solid, or liquid state, or in solution. An example of dissociation is the reversible reaction of hydrogen iodide at high temperatures 2HI(g) H2(g) + I2(g) The term dissociation is also applied to ionization reactions of acids and bases in water. For example, HCN + H2O H2O + + CN− This is often regarded as a straight forward dissociation into ions. HCN H+ + CN− Dissociation constant is a constant whose numerical value depends on the equilibrium between the undissociated and dissociated forms of a molecule. A higher value indicates greater dissociation. The equilibrium constant of such a dissociation is called the acid dissociation constant or acidity constant, given by Ka = [H+ ] ⋅ [CN− ] [HCN] … (2.49)

The concentration of water [H2O] can be taken as constant. Similarly, for a base, the equilibrium in following reaction is also dissociation; NH3 NH + 4 + OH− The base dissociation constant or basicity constant, given by Kb = [NH + 4 ] ⋅ [OH− ] [NH4] … (2.50) where, Ka or Kb is the measures of the strength of the acid (base).

The acid-base dissociation constant is a measure of the tendency of a molecule or ion to keep a proton (H+ ) at its ionization center(s), and is related to the ionization ability of chemical species. Since water is a very polar solvent (ε = 80), ionization will increase the likelihood of a species to be taken-up into aqueous solution. If a molecule does not readily ionize, it will tend to stay in a non-polar solvent such as cyclohexane (ε = 2) or octanol (ε =10). Dissociation constant is the core property of substance that defines its chemical and biological behaviour. In biological terms, dissociation constant is important in determining whether a molecule will be taken-up by aqueous tissue components or lipid membranes. The scientists require an understanding of dissociation constant because it impacts the choice of techniques used to identify and isolate the compound of interest. Dissociation constant is also closely related to the concepts of pH (the acidity of solution) and log P (the partition coefficient of a neutral compound between immiscible liquids).

States and Properties of Matter and Physicochemical Properties of Drug Molecules

States and Properties of Matter and Physicochemical Properties of Drug Molecules

States of Matter and Properties of Matter

States of Matter

Changes in the State of Matter

Latent Heat

Vapour Pressure

Critical Point

Eutectic Mixtures

Gases

Aerosols

Inhalers

Relative Humidity

Liquid Complexes

Liquid Crystals

Glassy State

Solids

Physicochemical Properties of Drug Molecules: Determinations and Application

Refractive Index

Optical Rotation

Measurement of Optical Activity:

Dielectric Constant

Dipole Moment

Dissociation Constant

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