Physical Pharmacy PDF

 Force ➔ A push or pull required to set a body in motion ➔ The larger the mass of the body and the greater the...

 Force ➔ A push or pull required to set a body in motion ➔ The larger the mass of the body and the greater the required acceleration, the greater the force that one must exert. Terminologies Pressure Physical Pharmacy ➔ Defined as the force per unit area; the unit ➔ Area of pharmacy that dealt with the quantitative commonly used in science is dyne/cm. and theoretical principles of physicochemical science as they applied to the practice of pharmacy Energy Pharmaceutical Science ➔ Energy is frequently defined as the condition of a ➔ Biomedical aspects of the practice of pharmacy. body that gives it the capacity to do work ➔ Energy may be classified as kinetic energy or Dosage Form potential energy. ➔ An entity that is administered to the patients so that they receive an effective dose of a drug. Temperature Significant figures ➔ It is a physical property of matter that quantitatively ➔ Any digit used to represent a magnitude or a expresses the common notions of hot and cold.Its quantity in the place in which it stands assigned unit is degree. RULES: ➔ On the centigrade and the Kelvin or absolute ➔ Non-zero digits are always significant. scales, the freezing and boiling points of pure ➔ Zeros between non-zero are always significant water at 1 atmosphere (atm) pressure are ➔ Zeros at the beginning of a number are never separated by 100 degrees. significant, with a special exception: a decimal ➔ Zero degree on the centigrade scale is equals to point. 273.15 on the Kelvin scale. ➔ Zeros that fall at the end of a number after a decimal point are always significant. Volume ➔ The measurable quantity, also derived from length ➔ Reference std is CUBIC METER; its CGS unit is one millionth of this value or 1 cubic centimeter (cc orcm3). ➔ This was originally defined in terms of liter, the volume of kilogram water at 1 atm and 4°C. Mass ➔ The reference standard is kilogram It is the mass of Density a platinum-iridium block preserved at the Bureau of ➔ This is a derived quantity since it combines the units Weights and Measures. The practical unit of mass of mass and volume. in the CGS system is gram (g),which is one ➔ Defined as the mass per unit volume at a fixed R.A. Cuarteros PPAR211LEC │2  temp and pressure and is expressed in the cgs system in grams per cubic centimeter (g/cm3). 𝑀𝑎𝑠𝑠 𝑑𝑒𝑛𝑠𝑖𝑡𝑦 = 𝑉𝑜𝑙𝑢𝑚𝑒 Specific Gravity ➔ The ratio of the weight of a given substance to the weight of an equal volume of a substance is chosen as the standard. ➔ It is a means of determining the strength, purity, or volume of a substance. 𝑤𝑒𝑖𝑔ℎ𝑡 𝑜𝑓 𝑠𝑢𝑏𝑠𝑡𝑎𝑛𝑐𝑒 𝑠𝑝𝑒𝑐𝑖𝑓𝑖𝑐 𝑔𝑟𝑎𝑣𝑖𝑡𝑦 = 𝑤𝑒𝑖𝑔ℎ𝑡 𝑜𝑓 𝑒𝑞𝑢𝑎𝑙 𝑣𝑜𝑙𝑢𝑚𝑒 𝑜𝑓 𝑤𝑎𝑡𝑒𝑟 Precision ➔ A measure of the agreement among the values in a group of data. ➔ Accuracy ➔ The agreement between the data and the true value. Statistical Methods and the Analysis of Errors Determinate (Constant) ➔ Errors that although sometimes unsuspected, may be avoided or determined and corrected once they are uncovered Indeterminate (Accident or Chance) ➔ The results of a series of tests will yield a random pattern around an average or central value, known as the mean Pseudo accidental or variable determinate errors ➔ Arise from random fluctuations in the temperature or other external factors and from the variations involved in reading instruments. ➔ It can be corrected by careful analysis and refinement of techniques A Measures of Dispersion ➔ Variability is expressed in range, mean and standard deviation ◆ Range = the lowest and highest value ◆ Standard deviation = the average deviation from the mean ◆ Coefficient of variation percent = more useful in comparing variability R.A. Cuarteros PPAR211LEC │3  intramolecular & intermolecular forces of attraction Intramolecular ➔ Are forces within molecules that caused by the attractionand repulsion of charged particles. 1.Ionic or Electrovalent Bond ➔ Electrostatic force of attraction between ions of opposite charge. ➔ is a type of chemical bond formed through an Molecular Forces electrostatic attraction between two oppositely ➔ Knowledge of these forces and their equilibrium is charged ions. important in understanding the: ➔ Ionic bonds are formed between a cation, which is 1. properties of gases, liquids, and gases usually a metal, and an anion, which is usually a 2. interfacial phenomena nonmetal. 3. flocculation in suspensions ➔ Transfer of electrons between a non-metal and 4. stabilization of emulsions metal 5. compaction of powders in capsules ➔ a bond is formed by the complete transfer of 6. dispersion of powders or liquid droplets in valence electron(s) between atoms. aerosols ➔ generates two oppositely charged ions. 7. compression of granules to form tablets ➔ metal loses electrons to become a positively ➔ Atoms aggregate to form molecules and lattices. charged cation, whereas the nonmetal accepts Molecules aggregate to form condensed phases of those electrons to become a negatively charged matter. anion. ➔ The aggregation of atoms, a positively charged Na+Cl→Na+ + Cl− → NaCl ions, and molecules are a consequence of electrical Properties of Ionic or Electrovalent Compounds forces exerted on the electrons of one particle by ➔ Crystalline solids- rigidity and strength the nucleus (or nuclei) of the other. ➔ High melting andboiling points ➔ There are 2 types of attraction in molecules: ➔ Conduct electricity in a molten and aqueous state ➔ They are hard ➔ They are brittle ➔ Soluble in polar solvents such as water (solute-solvent interactions) R.A. Cuarteros PPAR211LEC │4  2.Covalent Bonds or bronze. ➔ Formed between atoms with a small difference in Influences the strength of the bond: electronegativity. ➔ Availability of electrons ➔ is a form of chemical bonding ➔ More available delocalized electrons, the stronger characterized by the sharing the electrostatic attraction, the stronger the of one or more pairs of metallic bond. electrons between atoms, in ➔ Size of the charge on metal ion order to produce a mutual ➔ Larger the charge size, the stronger the metallic attraction, which holds the bond. resultant molecule together. Intermolecular ➔ are the forces of attraction between A. Polar Covalent Bonds molecules. ➔ A polar covalent bond is formed when atoms of ➔ Cohesion is the molecular attraction or joining of slightly different electronegativities share electrons. the surfaces of two pieces of the same B. Non-polar Covalent Bonds substance. ➔ A nonpolar covalent bond is formed between same ➔ Adhesion is any attraction process between atoms or atoms with very similar dissimilar molecules that can potentially bring electronegativities— the difference in them in “direct contact". electronegativity between bonded atoms is less Repulsive and Attractive Forces than 0.5. ➔ When molecules interact, both repulsive and attractive forces operate. ➔ Attractive forces are inversely proportional to the distance of separation. ➔ As two atoms or molecules are brought closer Properties of Covalent Compounds together, the opposite charges and binding forces ➔ Liquids and gasses at room temperature in the two molecules are closer together than the ➔ Relatively low boiling point. similar charges and forces, causing the molecules ➔ Do not conduct electricity to attract one another. ➔ Insoluble in polar solvent ➔ The negatively charged electron clouds of ➔ Soluble in non-polar solvent molecules largely govern the balance (equilibrium) 3. Metallic Bonds forces between the two molecules. ➔ Positive ions surrounded by a sea of mobile 1. Van Der Waals Forces (delocalized) electrons. Strong electrostatic force of ➔ Relate to non-ionic interactions between molecules attraction binds the system together. yet they involve charge-charge interactions. ➔ This type of covalent bonding specifically occurs ➔ Molecules frequently tend to align themselves with between atoms of metals, in which the valence their neighbors, so that the negative pole of one electrons are free to move through the lattice. molecule points toward the positive pole of the ➔ This bond is formed via the attraction of the mobile next. electrons—referred to as sea of electrons—and the ◆ Keesom fixed positively charged metal ions. ◆ Debye ➔ Metallic bonds ◆ London are present in A. Keesom or Dipole-Dipole Forces samples of pure ➔ Dipole-dipole forces exist between polar molecules elemental metals, where the positive end of one molecule is attracted such as gold or to the negative end of another molecule. aluminum, or ➔ The greater the polarity (difference in alloys, like brass electronegativity of the atoms in the molecule), the R.A. Cuarteros PPAR211LEC │5  stronger the dipole-dipole attraction. ➔ The more electrons a molecule has, the stronger ➔ Dipole-dipole attractions are very weak and the London dispersion forces are. substances held together by these forces have low ➔ For example, bromine, Br2. melting and boiling point temperatures. 2. Ion-Dipole Interaction ➔ Generally, substances held together by dipole ➔ Attractions that occur between ions and polar dipole attractions are gases at room temperature. molecules. ➔ Named after Willem Hendrik Keesom. ➔ These types of interactions ➔ Keesom forces is also known as account in part for the “Orientation Effect” solubility of ionic crystalline ➔ InKeesom forces, permanent dipoles interact substance in water. with one another with an ion-like fashion. ➔ Forces occur when the partially positively charged 3. Ion-Induced Dipole Interaction part of a molecule interacts with the partially ➔ This is a weak attraction that results when the negatively charged part of the neighboring approach of an ion induces a dipole in an atom or molecule. in a nonpolar molecule by disturbing the ➔ The prerequisite for this type of attraction to exist is arrangement of electrons in the nonpolar species. partially charged ions—for example, the case of ➔ Ion-induced dipole forces are presumably involved polar covalent bonds such as hydrogen chloride. in the formation of the iodide complex. ➔ Dipole-dipole interactions are the strongest ➔ Solubility of iodine in KI is an example. intermolecular force of attraction. I2 + KI = KI3 4. Hydrogen Bond ➔ The interaction between a molecule containing a hydrogen atom and a strongly electronegative atom such as fluorine, oxygen, or nitrogen is of B. Debye or Dipole-Induced Dipole Forces particular interest. ➔ Permanent dipoles are capable of inducing an ➔ Because of the small size of a hydrogen atom and electric dipole in nonpolar molecules (which are its electrostatic field, it can move in close to the easily polarizable) in order to produce electronegative atom and form an electrostatic dipole-induced dipole type of union known as a hydrogen bond or ➔ This interaction is called the Debye force after Peter hydrogen bridge. J.W. Debye ➔ The partially positive end of hydrogen is attracted C. London or Induced Dipole-Induced Dipole Forces to the partially negative end of the oxygen, ➔ They result from the movement of the electrons in nitrogen, or fluorine of the molecule which generates temporary positive another molecule. and negative regions in the molecule. ➔ Hydrogen bonding is a ➔ Caused by correlated movements of the electrons relatively strong force of in interacting molecules. The electrons, which attraction between belong to different molecules, start "feeling" and molecules, and avoiding each other at the short intermolecular considerable energy is distances, which is frequently described as required to break formation of "instantaneous dipoles" that attract hydrogen bonds. each other. ➔ Nonpolar molecules can induce polarity in one another by induced dipole-induced dipole, ➔ The weakest intermolecular forces and exist between all types of molecules, whether ionic or covalent—polar or nonpolar. R.A. Cuarteros PPAR211LEC │6  ➔ Hydrogen bonding plays an important role in biology; hydrogen bonds are responsible for holding nucleotide bases together in DNA and RNA. Relative Strength of Intermolecular Forces of Attraction Intermolecular Occurs between Relative strength force Dipole-dipole Partially charged attraction oppositely ions Strong charged ions Hydrogen H, atom and O, Strongest bonding N/ or F atom of the dipole-dipole attractions London Temporary Weakest dispersion or induced attraction dipoles R.A. Cuarteros PPAR211LEC │7  Ideal Gas Vs. Real Gas Ideal Real Made of small particles that SAME have MASS Gases are mostly empty SAME space Gas Low Density SAME ➔ Have kinetic energy rapid motion Gas particle are in constant SAME ➔ Weak intermolecular forces random straight line motion ➔ Capable of filling all available spaces ➔ Compressible There are NO attractive or There are VERY SMALL ➔ Many are invincible repulsive forces between attractive and repulsive Kinetic Molecular Theory particles forces between particles 1. Gases are composed of particles called atoms POr Particles have NO volume Particles have a very small molecules, the total volume of which is so small as volume to be negligible in relation to the volume of the Collisions are ELASTIC (no Collisions are inelastic (NOT space in which the molecules are confined. This loss in total kinetic energy) Elastic) (When gas particles condition is approxi- mated in actual gases only at collide they will lose energy) low pressures and high tem- peratures, in which case the molecules of the gas are far apart. Ideal gases obey all gas laws Real gases obey gas laws 2. The particles of the gas do not attract one another, under all conditions of only at low pressure and but instead move with complete independence; temperature and pressure high temperature again, this statement applies only at low pressures, The volume occupied by the The volume occupied by the 3. The particles exhibit continuous random motion molecules is negligible as molecules is not negligible as owing to their kinetic energy. The average kinetic compared to the total compared to the total energy, E, is directly proportional to the absolute volume occupied by the gas volume of the gas temperature of the gas, or E = (k)RT. 4. The molecules exhibit perfect elasticity; that is, The force of attraction The force of attraction are there is no net loss of speed or transfer of energy among the molecules are not negligible at all after they collide with one another and with the negligible temperatures and pressures molecules in the walls of the confining vessel, wlch Obeys ideal gas equation Obeys Van der Waals latter effect accounts for the gas pressure. PV=nRT equation Although the net velocity, and therefore the (𝑃 + ) (𝑉 − 𝑛𝑏) = 𝑛𝑅𝑇 2 𝑎𝑛 2 average kinetic energy, does not change on 𝑉 collision, the speed and energy of the individual The Gaseous State molecules may differ widely at any instant. More ➔ The Ideal Gas Law simply stated, the net velocity can be an average ➔ PV=nRT is the ideal gas equation velocity of many molecules; thus, a distribution of ➔ Boyle, Charles, and Gay-Lusaac formulated the individual molecular velocities can be present in the gas laws. system. R.A. Cuarteros PPAR211LEC │8  Boyles Law The Liquid State ➔ Inverse relationship pressure and volume at ➔ Liquefaction of Gasses constant temperature ➔ When a gas is cooled, it loses some of its kinetic energy in the form of heat, and the velocity of the molecules decreases. ➔ If pressure is applied to the gas, the molecules are brought within the sphere of the van der Waals interaction forces and pass into the liquid state. ➔ Because of these forces, liquids are considerably Charles Law denser than gases and occupy a definite volume. ➔ Direct relationship volume and temperature at ➔ The transitions from a gas to a liquid and from a constant pressure liquid to a solid depend not only on the temperature, but also on the pressure to which the substance is subjected Aerosols ➔ Gases can be liquefied by increasing pressure, provided we work below the critical temperature. ➔ When the pressure is reduced, the molecules Gay Lussac Law expand and the liquid reverts to a gas. This ➔ Direct relationship pressure and temp at constant reversible change of state is the basic principle volume involve in the preparation of pharmaceutical aerosols. ➔ In such products, a drug is dissolved or suspended in a propellant ➔ The container is so designed that, by depressing a valve, some of the drug-propellant mixture is expelled owing to the excess pressure inside the container. ➔ If the drug is nonvolatile, it forms a fine spray as it Henry’s Law of Gas Solubility leaves the valves orifice; at the same time, the ➔ "At a constant temperature, the amount of a given liquid propellant vaporizes off. gas that dissolves in a given type and volume of ➔ The propellant used in these products is frequently liquid is directly proportional to the partial pressure a mixture of fluorinated hydrocarbons, although of that gas in equilibrium with that liquid.“ other gases, such as nitrogen and carbon dioxide, ➔ amount of gas dissolve in plasma is proportional to are increasingly used. the partial pressure of the gas in equilibrium with Clausius Clapeyron Equation the plasma ➔ Relationship between the vapor pressure and the Dalton’s Law absolute temperature of a liquid 𝐿𝑜𝑔𝑃2 𝐻𝑣 (𝑇2−𝑇1) ➔ the total pressure exerted is equal to the sum of 𝑃1 = 2.303 𝑅𝑇1𝑇2 the partial pressures of the individual gases. ➔ States that the total pressure in a mixture of gases is equal to the sum of the partial pressure of each gas R.A. Cuarteros PPAR211LEC │9  The Solid State Crystalline Solids ➔ The structural units of crystalline solids, such as ice, sodium chloride, and menthol, are arranged in fixed geometric patterns or lattices. Crystalline solids, unlike liquids and gases, have definite shapes and an orderly arrangement of units. ➔ Gases are easily compressed, whereas solids, like Latent Heat of Vaporization liquids, are practically incompressible. Crystalline ➔ Heat taken up when the liquids vaporize and are solids show definite melting points, passing rather lost or liberated when vapor condense to liquids sharply from the solid to the liquid state. The various crystal forms are divided into six crystal Molar Heat of Vaporization systems ➔ ∆Hv is the heat absorbed by 1 mole of liquid when ➔ Systems: it passes into the vapor state ◆ cubic (sodium chloride) Boiling Point ◆ tetragonal (urea) ◆ hexagonal (iodoform) ➔ If a liquid is placed in an open container and ◆ rhombic (iodine) heated until; the vapor pressure equals the ◆ monoclinic (sucrose) atmospheric pressure, the vapor is seen to form ◆ triclinic (boric acid) bubbles that rise rapidly through the liquid and ➔ Melting Point and Heat of Fusion escape into the gaseous state. ◆ The temperature at which a liquid passes into ➔ boiling point refers to the temperature at which the a solid state is known as the freezing point. vapor pressure of the liquid equals the external or ◆ It is also the melting point of a pure crystalline atmospheric pressure compound. ➔ The pressure at sea level is about 760mm Hg; at ◆ The freezing point or melting point of a pure higher elevations, the atmospheric pressure crystalline solid is strictly defined as the decreases and the boiling point is lowered. At a temperature at which the pure liquid and solid pressure of 700 mm Hg, water boils at 97.7 °C, at exist in equilibrium. 17.5 mm Hg, it boils at 20 °C. ◆ In practice, it is taken as the temperature of the equilibrium mixture at an external pressure of 1 atm, this is sometimes known as the normal freezing or melting point. ➔ Melting Point and Intermolecular Forces ◆ The heat of fusion: the heat required to increase interatomic or intermolecular distances in crystals, thus allowing melting point to occur. ◆ A crystal that is bound together by weak forces generally has a low heat of fusion and a low melting point, whereas one bound together by strong forces has a high heat of fusion and a high melting point. ➔ Latent Heat of Fusion ◆ Heat absorbed when 1 g of solid melts or the heat liberated when it freezes ◆ Molar heat of fusion: amount of heat absorbed when 1 mole of solid changes to 1 R.A. Cuarteros PPAR211LEC │10  mole of liquid ➔ Antipsychotic haloperidol and bromperidol ➔ Types of Crystal Cacao Butter Polymorphism ◆ Atomic (strong carbon covalent bond) ➔ 33 deg C is the lowest possible temp to heat cacao ◆ Metallic (strong metal bond) butter in preparing suppositories so as not to ◆ Molecular (van der waals) destroy crystal nuclei. ◆ Ionic (electrostatic ionic bond) ➔ The melting point of the beta polymorph is 34.5 Lattice Defects deg C. ➔ Lattice defects dictate the hardness and strength of ➔ If theobroma is heated to about 35 deg C, the metals. stable beta nuclei will be destroyed. ➔ The presence of the electron gas in metals provides ➔ Theobroma oil is a polymorphous natural fat. its conductivity ➔ It has a narrow temperature range (34-36 deg C) ➔ Lattice defects = kind of imperfections of crystals Four Polymorphic Forms ➔ Metallic crystals are positively charged ions in a ➔ Unstable gamma form: melts at 18 deg C freely moving electrons e- ➔ Alpha form: melts at 22 deg C Polymorphism ➔ Beta prime form: melts at 28 deg C ➔ Some elemental substances, such as carbon and ➔ Stable beta form: melts at 34.5 deg C sulfur, may exist in more than one crystalline and 2 Properties of Polymorphism are said to be polymorphic. ➔ Monotropic ➔ Polymorphs generally have different melting points, ◆ Polymorphism x-ray diffraction patterns, and solubility’s, even involves a change though they are chemically identical. in one direction ➔ Difference is solubility will produce a drug with and is usually from different rate of dissolution. a metastable to ➔ As a result of this, one polymorph may be more stable form biologically active than another polymorph of the ◆ Examples: same drug. long-chain organics ➔ Polymorphs have different stabilities and may eg, fatty acids, glycerides and fats spontaneously convert from the metastable form at ➔ Enantiotropic temperature to a stable form ◆ Polymorphism ➔ The formation of polymorphs of a compound may involves a change in depend upon several variables pertaining to the various direction and crystallization process, including: is usually associated ◆ Solvent differences with changes in ◆ Impurities temperature or ◆ Level of supersaturation solvent Examples: ◆ Temperature at which crystallization is carried Sulfur, water, aspirin out Some Polymorphic Drugs ◆ Geometry of the covalent bonds ➔ Acetaminophen, caffeine, chloramphenicol ◆ Attraction and repulsion of cations and anions palmitate, cimetidine, nifedipine, phenobarbital Na, ◆ Temperature phenytoin, progesterone, theophylline ◆ Pressure Pseudopolymorphs Pharmaceutical Importance: ➔ Also called as solvates ➔ Polymorphic state of chloramphenicol palmitate ➔ Crystals containing solvent molecules has significant influence on biological activity Classes of Solid According to Characteristic ➔ 5 forms of cortisone acetate, 4 unstable forms causes caking of the crystals ➔ Anisotropic solids: Are those showing different ➔ Beta form of antiestrogenic tamoxifen citrate – characteristics in various directions along crystal stable ➔ These crystals have unlike light properties in R.A. Cuarteros PPAR211LEC │11  different sites properties of liquids. ➔ (light is refracted at different site angles from ◆ At the same time they possess the property of different sites) being birefringent, a property associated with ➔ Isotropic solids: Crystals that exhibit similar crystals. properties and characteristics in all directions Birefringence ➔ light properties in all directions ➔ The property of crystals and mesophase to divide ➔ These are amorphous solids and cubic crystals passing light into two components with different Amorphous Solids velocities and refractive index ➔ Amorphous solids 2 Types of Liquid Crystal may be ➔ Smectic liquid crystals: considered as ◆ known as the Soaplike or Grease-like crystals super cooled ◆ mobile in two directions liquids in which ◆ Considered as the mesophase of most the molecules are pharmaceutical significance arranged in a random manner somewhat as in the ◆ it is the phase that usually forms a ternary liquid state. (complex) mixtures with other additives. ➔ Non-crystalline ◆ They are usually used as components of ➔ More soluble, therefore more bioavailable - (more emulsions and are employed in the prompt action) solubilization of water-insoluble materials ➔ Example are antibiotics. owing to their highly viscous nature. ➔ Other includes beef, pork, human insulin ➔ Thread-like crystals ➔ Substance such as glass, pitch, and many synthetic ◆ Known as the nematic form of liquid crystals plastics are amorphous solids. ◆ mobile in three directions/dimensions ➔ They differ from crystalline solids in that they tend ◆ Include Cholesterictype to flow when subjected to sufficient pressure over ◆ Lyotropic liquid crystals : Crystals derived from a period of time, and they do not have definite the action of certain solvents on solids. melting points. ◆ Thermotropic liquid crystals : Crystals obtained Polymeric Solids by heating solids to obtain mesophase ➔ Most polymers used in pharmaceuticals ◆ Cholesteryl benzoate: The first thermotropic ➔ These are carbon-based formed by the hybrodized liquid crystal that was recorded studied by carbon atom with S and P orbitals to give four Reinitzer in 1888 valency bonds at fairly well defined angles. Superfluid State ➔ Examples: Polyvinyl alcohol, methylcellulose, ➔ Is a mesophase formed from the gaseous state carbomer, carbopol, polymethylmethacrylate where the gas is held under a combination of (PMMA), PEG, PP, PVP temperature and pressure that exceed critical The Liquid Crystalline State points ➔ A fourth state of matter is the liquid crystalline Phase Rule state or mesophase ➔ The phase rule, which ➔ Structure of Liquid Crystals was c. ◆ Molecules in the liquid state are mobile in ➔ This is expressed by the ff three directions and can also rotate about equation three axes perpendicular to one another. F=C–P+2 ➔ Properties and Significance of Liquid Crystals ◆ Because of their intermediate nature, liquid ➔ Where: F: is the number crystals have some of the properties of liquids of degrees of freedom in and some of solids. the system ◆ For example, liquid crystals are mobile and ➔ C: is the number of thus can be considered to have the flow component P: is the number of phases present R.A. Cuarteros PPAR211LEC │12  ➔ In the Gibbs’ phase rule, the number of phases will represent the number of homogenous, physically distinct portion of a system that is separated from other portions of the system by bounding surfaces ➔ Ex. Water and its vapor is a 2-phase system ➔ In the Gibbs’ phase rule, the degrees of freedom is the least number of intensive variables that must be fixed to describe the system completely System ➔ A system is defined as bound space or a definite quantity of substance that is under observation and experimentation ➔ From the equation of the Gibbs’ phase rule, we can see that F and C are directly proportional, hence an increase in the number of components will also mean an increase in the degrees of freedom. ➔ On the other hand, F is equal to –P, this means than an increase in the number of phases will cause a decrease in the degrees of freedom. We can therefore generalize that as the system becomes more complex we have to fix more variables to define it. ➔ Example (1) ◆ Determine the degrees of freedom for a system containing liquid water, liquid ethanol in equilibrium with their vapors ◆ F=2–2+2=2 ➔ Example (2) ◆ Determine the F containing either ice, water, or water vapor ◆ Since the system is made up of one chemical entitiy (H2O), the number of component is 1. If the system is to contain either ice, water or water vapor, then there is only a single phase to consider. ◆ Substituting in the formula for the Gibbs’ phase rule, we compute F as follows ◆ F = 1 – 1 + 2 = 2, the system is said to be bivariant, wherein two variables must be fixed in defining the system ◆ If P = 1 : System is bivariant (F=2) ◆ If P = 2 : System is univariant (F=1) ◆ If P = 3 : System is invariant (F=0) R.A. Cuarteros PPAR211LEC │13  Solute Solvent Example solid solid Gold-Silver mixture solid liquid Paraffin in Mineral Oil gas solid Hydrogen in Palladium Definition of Terms liquid liquid Aqueous Sodium Chloride Solution System gas liquid Carbonated Water ➔ is generally considered to be a bounded space or solid gas Iodine Vapor in Air an exact quantity of a material substance. Dispersion liquid gas Water in Oxygen ➔ consists of at least two phases with one or more gas gas air dispersed (internal) phases contained in a single continuous (external) phase. ➔ Solutions can be classified as: Phase ◆ A saturated solution contains the maximum quantity of solute that dissolves at that ➔ a distinct homogenous part of a system separated temperature. by a definite boundaries from other parts of a ◆ An unsaturated solution contains less than the system maximum amount of solute that can dissolve True Solutions at a particular temperature. ➔ defined as a mixture of two or more components ◆ A supersaturated solution contains more than that form a homogeneous molecular dispersion. the maximum amount of solute that a solvent ➔ A solution is a homogenous mixture of 2 or more can dissolve at a given temperature. substances in a single phase. Ionic Compounds ➔ One constituent is usually regarded as the SOLVENT and the others as SOLUTES. ➔ Many reactions involve ionic compounds, especially reactions in water — aqueous solutions. Parts of Solution Classes of Solutes ➔ SOLUTE – the part of a solution that is being dissolved (usually the lesser amount) ➔ Nonelectrolytes - are substances that do not ionize ➔ SOLVENT – the part of a solution that dissolves the when dissolved in water and therefore do not solute (usually the greater amount) conduct an electric current through the solution ➔ Solute + Solvent = Solution ➔ Electrolytes - are substances that form ions in solution, conduct electric current, and show apparent “anomalous” colligative properties. Aqueous Solutions ➔ How do we know ions are present in aqueous solutions? ◆ They are called ELECTROLYTES ◆ HCl, MgCl2 , and NaCl are strong electrolytes. R.A. Cuarteros PPAR211LEC │14  ◆ Ephedrine and Phenobarbital are Weak electrolytes ➔ Some compounds dissolve in water but do not conduct electricity. They are called nonelectrolytes. ➔ Examples: ◆ sugar ◆ glycerine ◆ naphthalene ◆ urea Expressions of Concentration ➔ A molar solution contains a mole (one gram molecular weight) in a liter of solution or one millimole per milliliter of solution. ➔ A mole is the molecular weight expressed in grams. ➔ A millimole is one thousandth part of a mole. 𝑛𝑜. 𝑜𝑓 𝑚𝑜𝑙 𝑠𝑜𝑙𝑢𝑡𝑒 𝑔 𝑠𝑜𝑙𝑢𝑡𝑒/ 𝑚𝑜𝑙𝑎𝑟 𝑚𝑎𝑠𝑠 𝑔 𝑠𝑜𝑙𝑢𝑡𝑒 𝑀 = 𝑛𝑜. 𝑜𝑓 𝐿 𝑜𝑓 𝑠𝑜𝑙𝑛 = 𝐿 𝑜𝑓 𝑠𝑜𝑙𝑛 = (𝑚𝑜𝑙 𝑤)(𝐿 𝑜𝑓 𝑠𝑜𝑙𝑛) Molal ➔ solution contains a mole (one gram molecular weight) in one thousand grams of solution. 𝑛𝑜. 𝑜𝑓 𝑚𝑜𝑙 𝑠𝑜𝑙𝑢𝑡𝑒 𝑔 𝑠𝑜𝑙𝑢𝑡𝑒/ 𝑚𝑜𝑙𝑎𝑟 𝑚𝑎𝑠𝑠 𝑔 𝑠𝑜𝑙𝑢𝑡𝑒 𝑚 = 𝑛𝑜. 𝑜𝑓 𝑘𝑔 𝑜𝑓 𝑠𝑜𝑙𝑛 = 𝑘𝑔 𝑜𝑓 𝑠𝑜𝑙𝑛 = (𝑚𝑜𝑙 𝑤)(𝑘𝑔 𝑜𝑓 𝑠𝑜𝑙𝑛) Normal Concentration Units ➔ solution contains one gram equivalent weight of solute in a liter of solution or one gram ➔ An IDEAL SOLUTION is one where the properties milliequivalent weight in a milliliter of solution. depend only on the concentration of solute. 𝑛𝑜. 𝑜𝑓 𝑔−𝑒𝑞 𝑤𝑡 𝑔 𝑠𝑜𝑙𝑢𝑡𝑒/ 𝑒𝑞 𝑤𝑡 𝑔 𝑠𝑜𝑙𝑢𝑡𝑒 𝑁 = = = Molality 𝐿 𝑜𝑓 𝑠𝑜𝑙𝑛 𝐿 𝑜𝑓 𝑠𝑜𝑙𝑛 (𝑒𝑞 𝑤𝑡)(𝐿 𝑜𝑓 𝑠𝑜𝑙𝑛) Molar ➔ Moles of solute in 1000g of solvent 𝑚𝑜𝑙𝑒𝑠 𝑠𝑜𝑙𝑢𝑡𝑒 𝑚𝑜𝑙𝑎𝑙𝑖𝑡𝑦 = ➔ Moles (gram molecular weights) of solute in1 liter 𝑘𝑔 𝑜𝑓 𝑠𝑜𝑙𝑛 𝑔𝑟𝑎𝑚𝑠 𝑠𝑜𝑙𝑢𝑡𝑒 of solution % 𝑏𝑦 𝑚𝑎𝑠𝑠 = 𝑔𝑟𝑎𝑚𝑠 𝑠𝑜𝑙𝑛 𝑛𝑜. 𝑜𝑓 𝑚𝑜𝑙 𝑠𝑜𝑙𝑢𝑡𝑒 𝑔 𝑠𝑜𝑙𝑢𝑡𝑒/ 𝑚𝑜𝑙𝑎𝑟 𝑚𝑎𝑠𝑠 𝑔 𝑠𝑜𝑙𝑢𝑡𝑒 𝑀 = 𝑛𝑜. 𝑜𝑓 𝐿 𝑜𝑓 𝑠𝑜𝑙𝑛 = 𝐿 𝑜𝑓 𝑠𝑜𝑙𝑛 = (𝑚𝑜𝑙 𝑤)(𝐿 𝑜𝑓 𝑠𝑜𝑙𝑛) 𝑚𝑜𝑙𝑒𝑠 𝑠𝑜𝑙𝑢𝑡𝑒 𝑀𝑜𝑙𝑎𝑟𝑖𝑡𝑦 = 𝐿 𝑜𝑓 𝑠𝑜𝑙𝑛 R.A. Cuarteros PPAR211LEC │15  in a redox reaction. ➔ a.k.a. “Gram Equivalent Weight” 𝐴𝑡𝑜𝑚𝑖𝑐 𝑊𝑒𝑖𝑔ℎ𝑡 𝐸𝑞𝑢𝑖𝑣𝑎𝑙𝑒𝑛𝑡 𝑊𝑒𝑖𝑔ℎ𝑡 = 𝑁𝑢𝑚𝑏𝑒𝑟 𝑜𝑓 𝐸𝑞𝑢𝑖𝑣𝑎𝑙𝑒𝑛𝑡𝑠 𝑝𝑒𝑟 𝑎𝑡𝑜𝑚𝑖𝑐 𝑤𝑒𝑖𝑔ℎ𝑡 (𝑉𝑎𝑙𝑒𝑛𝑐𝑒) Gram Equivalent Weight 𝑀𝑊 𝑔𝐸𝑊 = 𝐹𝑎𝑐𝑡𝑜𝑟 𝐹𝑎𝑐𝑡𝑜𝑟 = 𝑡𝑜𝑡𝑎𝑙 𝑝𝑜𝑠𝑖𝑡𝑖𝑣𝑒 𝑐ℎ𝑎𝑟𝑔𝑒 Milliosmoles ➔ The milliosmolar value of separate ions of an electrolyte may be obtained by dividing the concentration, in milligrams per liter, of the ion by its atomic weight. ➔ Osmolarity expressed in mOsmol/L Normality 𝑤𝑡 𝑖𝑛 𝑔/𝐿 = 𝑀𝑊 𝑖𝑛 𝑔 𝑥 # 𝑜𝑓 𝑠𝑝𝑒𝑐𝑖𝑒𝑠 (𝐹𝑎𝑐𝑡𝑜𝑟) 𝑥 1000 ➔ gram equivalent weights of solute in 1 liter of solution. 𝐸 𝑁𝑜𝑟𝑚𝑎𝑙𝑖𝑡𝑦 = 𝑉 Percentage expressions ➔ Percentage by weight (% w/w) ◆ grams of solute in 100 g of solution. ➔ Percentage by volume (% v/v) ◆ milliliters of solution in 100 ml of solution. ➔ Percentage weight in volume (% w/v) ◆ grams of solute in 100 ml of solution. Equivalents Weights Mole Fraction ➔ it is the mass of one equivalent, that is the mass of a given substance which will: ➔ Ratio of the moles of one constituent (e.g., the ◆ supply or react with one mole of hydrogen solute) of a solution to the total moles of all cations H+ in an acid–base reaction; or constituents (solute and solvent) 𝑀𝑜𝑙𝑒𝑠 𝑜𝑓 𝐴 ◆ supply or react with one mole of electrons e − 𝑋𝑎 = 𝑀𝑜𝑙𝑒𝑠 𝑜𝑓 𝐴 + 𝑀𝑜𝑙𝑒𝑠 𝑜𝑓 𝐵 + 𝑀𝑜𝑙𝑒𝑠 𝑜𝑓 𝐶.... R.A. Cuarteros PPAR211LEC │16  𝑀𝑜𝑙𝑒𝑠 𝑜𝑓 𝐵 𝑋𝑏 = 𝑀𝑜𝑙𝑒𝑠 𝑜𝑓 𝐴 + 𝑀𝑜𝑙𝑒𝑠 𝑜𝑓 𝐵 + 𝑀𝑜𝑙𝑒𝑠 𝑜𝑓 𝐶.... R.A. Cuarteros PPAR211LEC │17 R.A. Cuarteros PPAR211LEC │18  liquids in liquids Legends: gasses in liquids Topic ◆ Parts of Solution SOLUTE – the part of a solution that is being Sub topic dissolved (usually the lesser amount) Sub Sub Topic SOLVENT – the part of a solution that dissolves the solute (usually the greater amount) Solute + Solvent = Solution Definition of Terms ◆ System generally considered to be a bounded space or an exact quantity of a material substance. ◆ Dispersion consists of at least two phases with one or more dispersed (internal) phases contained in a single continuous (external) phase. ◆ Phase a distinct homogenous part of a system separated by a definite boundaries from other parts of a system. ◆ True Solutions Importance of Studying the Phenomenon of Solubility defined as a mixture of two or more ◆ Select the best solvent for a drug or a mixture of components that form a homogenous drugs. molecular dispersion. ◆ Overcome problems arising during preparation of Liquid Dosage Forms pharmaceutical solutions. ◆ Depending on the size of the dispersed particle, ◆ Have information about the structure and liquid dosage forms are classified as: intermolecular forces of the drug. true solutions ◆ Many drugs are formulated as solutions, or added colloidal solution as powder or solution forms to liquids. disperse systems ◆ Drugs with low aqueous solubility often present ◆ If sugar is dissolved in water, it is supposed that the problems related to their formulation and ultimate sugar particle is of molecular dimensions bioavailability and that a true solution is formed. Solution ◆ Between these two extremes lie colloidal solutions, ◆ a chemically and physically homogeneous mixture the dispersed particles of which are larger than of two or more substances. those of true solutions but smaller than the ◆ The term solution generally denotes a particles present in suspensions. homogeneous mixture that is liquid, even though it ◆ If very fine sand is mixed with water, a suspension is possible to have homogeneous mixtures that are of comparatively large particles, each consisting of solid or gaseous. many molecules, is obtained. ◆ Most important in pharmacy: solids in liquids (Note: eto yung difference nung tatlo) R.A. Cuarteros PPAR211LEC │19  USP Descriptive Terms of Solubility True Solution Colloidal Coarse Solution Dispersion Description Volume (mL) solvent that Homogenous Heterogenous Heterogenous dissolves 1 g or 1 part solute Solution in Solution in Solution in which which which Very soluble 10,000 be atoms, ions, dispersed; can suspended; can Insoluble molecules be combined or be large large molecules particles or (Note: Tandaan daw) combined ◆ Solubility Expressions particles The USP lists the solubility of drugs as: the number of ml of solvent in which 1g of solute Do not separate Do not separate Particles settle will dissolve. on standing on standing out E.g. 1g of boric acid dissolves in 18 mL of water, and in 4 mL of glycerin. (to interpret: Do not scatter Scatter light May scatter boric acid is soluble in water, boric acid is light (Tyndall effect) light, but are freely soluble in glycerin) not transparent Solubility ◆ The extent to which the solute dissolves is referred They may pass They may pass They can’t pass to as its solubility. through through through ◆ Solubility refers to the maximum amount of solute ordinary as well ordinary as well ordinary as well that can be dissolved in a given amount of solvent as ultra filters. as ultra filters. as ultra filters ◆ Extent to which dissolution proceeds under Examples: Examples: Examples: experimental conditions -- amount of solid passing 1. Solution of 1. Solution of 1. in solution with equilibrium (form saturated NaCl in water Scratch Pharmaceutical solution) between solution & excess (undissolved) 2. Solution of 2. Milk suspensions and substance glucose in water 3. Solutions of emulsions Gums 2. Grain of sand Classification of Solutions in water ◆ A. Based on Elemental Composition ORGANIC – compounds containing carbon (except CO2, CO, carbonates and cyanides) INORGANIC – compounds of other elements including acids, bases and salts ◆ B. Based on Ionization/ Electrolytic Property of Solute Electrolytic property – the ability of the solution to conduct electricity R.A. Cuarteros PPAR211LEC │20  1. Electrolytes (3) solvent and solute ○ substances whose aqueous solutions ◆ Steps of solid going into solution: contain ions and thus conduct electricity Step 1: Hole open in the solvent ○ Strong electrolytes – substances which Step 2: Onemolecule of the solid breaks away completely dissociates into ions e.g. salts, from the bulk strong acids, strong bases NaCl → Na+ Step 3: The solid molecule is enter into the (aq) + Cl- (aq) hole in the solvent ○ Weak electrolytes – substances which produce small amounts of ions; partially dissociated into ions e.g. weak acids, weak bases CH3COOH ↔ CH3COO- (aq) + H+ (aq) CH3COOH ionizes to form CH3COO- and H+. While this happens, some ions also combine to form back CH3COOH. This results to partial dissociation 2. Nonelectrolytes substances that does not dissociate into ions; form non conducting solutions e.g. most molecular compounds are nonelectrolytes Examples include sugar, glycerin, naphthalene Saturated Solution and urea ◆ A solvent can not dissolve any more solute Types of Solution ◆ A saturated solution is at equilibrium. Particles are ◆ Dilute – solution with low solute concentration dissolving and precipitating at the same rate ◆ Concentrated – one with high solute concentration Factors Affecting Solubility Degree of Saturation ◆ 1. Temperature ◆ Types of Solution Based on the Degree of Saturation Temperature will affect solubility. If the Saturated solution – a solution that contains solution process absorbs energy then the the maximum amount of solute the solvent solubility will be increased as the temperature can dissolves; no more solute can dissolve in it is increased. (solution in equilibrium) If the solution process releases energy then Unsaturated solution – solution containing the solubility will decrease with increasing amount of solute less than its solubility; more temperature. solute can dissolve in it Generally, an increase in the temperature of Supersaturated solution – solution containing the solution increases the solubility of a solid an amount of solute greater than the solute. solubility; unstable solution A few solid solutes are less soluble in warm Solubility Curve solutions. n For all gasses, solubility decreases ◆ Any solution can be made saturated, unsaturated, as the temperature of the solution increases. or supersaturated by changing the temperature. Generally, Thermodynamic Solubility of Drugs Gases: ↑ Temp, ↓ solubility, ↑ KE ↑ KE= ↑ motion ◆ The thermodynamic solubility of a drug in a solvent in molecules= break intermolecular bonds and is the maximum amount of the most stable escape from solution crystalline form that remains in solution in a given Solid: ↑ Temp, ↑ solubility volume of the solvent at a given temperature and pressure under equilibrium conditions. Endothermic reaction (ΔHsolvation>0) ◆ The equilibrium involves a balance of the energy of Increasing the temperature results in a stress three interactions against each other: on the reactants side from the additional heat. (1) solvent with solvent Le Châtelier's principle predicts that the (2) solute with solute R.A. Cuarteros PPAR211LEC │21  system shifts toward the product side in order ○ volatile oils & alcohols to form spirits, to alleviate this stress. By shifting towards the elixirs product side, more of the solid is dissociated Liquid-liquid systems may be divided into2 when equilibrium is again established, categories: resulting in increased solubility. 1. Systems showing complete miscibility such Exothermic reaction ((ΔHsolvation