PHYSICAL-PHARMACY-copy-compressed.pdf

 P M R E L I M S M CHAPTER I INTERPRETIVE TOOLS PHYSICAL PHARMACY -application of physical and chemical principles and laws in pharmaceutical sciences -to understand and develop dosage froms and drug delivery systems -study of physico-chemical properties of subs...

 P M R E L I M S M CHAPTER I INTERPRETIVE TOOLS PHYSICAL PHARMACY -application of physical and chemical principles and laws in pharmaceutical sciences -to understand and develop dosage froms and drug delivery systems -study of physico-chemical properties of substances used in drug formulation. DIMENSIONAL ANALYSIS also called “factor- label method or unit factor Example: method” How many seconds are there in 1 year? -is a problem-solving method that uses the fact that conversion factors: 365 days= 1 year,. 24 hours= 1 any number or expression can be multipliedby 1 day,. 60 minutes 1hour,. 60 seconds= 1 minute without changing its value. MEASUREMENT AND DATA TYPES ACCURACY PRECISION -closeness to true value/literature value -closeness of actual values -indicates how well a measurement -indicates how closely different agrees with the accepted or true value. measurements of a quantity agree SIGNIFICANT FIGURES -are consecutive figuresthat express the value of a denominate number accurately enough for a given purpose. the accurary varies with the number of significant figures, which are all absolute in value except the last, and this is properly called uncertain. Writing or Interpreting Significant Figures in Numbers RULE EXAMPLE All nonzero digits are 98.513 has five considered significant numbers SIGNIFICANT Leading zeros are NOT 0.00361 has three SIGNIFICANT significant numbers Trailing zeros in a 998.100 has six number containing a significant numbers decimal point are SIGNIFICANT All zeros between 607.123 has six other significant significant number figures are SIGNIFICANT  ERROR ERROR AND -defined as a deviation from DESCRIBING the absolute value or from the true average of a large number DETERMINATE ERRORS VARIABILITY resluts. -identifiable causes -aka systematic errors or bias 2 types: 1. Determinate errors 2. Indeterminate erros Causes: C.Apparatus errors due to poor A.Personal errors made by the individual analyst construction or calibration Example: inability to judge color changes sharply, (instrumental) resulting in habitual reading of end points in titration too late. Example: B.Errors of method caused by faulty procedure >inaccuracy in the calibration of (Methodic) buret or pipets Example: > incorrect sampling >inequality in the length of the >Contamination of precipitates arms of the balance >Improper selection of indicators >incorrect weights INDETERMINATE ERRORS -uncontrollable causes Errors that arise from random -aka random errors fluctuations in temperature or other external factors and from the variations -manifest themselves by slight variation in a series of observations made by the same involved in reading instruments are not observer under identical conditions. to be considered accidental or random. Instead, they belong to the class of >are intangible and their elimination by the determinate erros and are often called analyst is impossible pseudoaccidental or variable >eg. A differences in the judgement and determinate errors. skill of the analyst CENTRAL TENDENCY: MEAN, MEDIAN, MODE CENTRAL TENDENCY MEAN -is a single value thta attempts to describe a set of data by identifying the central position within that set of data. the arithmetic mean is obtain by adding together the results of a series of measurements and dividing the total by the number N of the measurements. EXAMPLE: ATTEMPT WEIGHT (g) 1 1.05 2 0.98 3 0.95 4 1.00 5 1.02 6 1.00 7 1.10 8 1.03 9 0.96 10 0.98 MEDIAN MODE -is the middle value of a range of values when -is te value in the data set that occurs most they are arranged in rank order(e.g., from often. lowest to highest). So, the median value of the list (1,2,3,4,5) is the number 3. e.g. 1,1,1,2,3,4,5,6,6 the mode number is 1.  METALLIC Bond CHAPTER II -exist as a collection of many atoms as +ions arranged in a well-defined 3D arrangement called crystal lattice with STATES OF MATTER some of the outermost electrons roaming around in the whole piece of the metal, I. FORCES OF ATTRACTION forming a sea of electrons around the INTRAmolecular forces metal atoms -forces within a molecule 3 major types 1. Metallic bond 2. Ionic Bond 3. Covalent bond IONIC BOND -affinity between oppositely charged particles -presents in salts/ionic compounds -forces that hold ions together in the crystal lattice of salt. -this bond is formed by the complete transfer of valence electron(s) between atoms. A type of chemical bond that generates two oppositely charged ions.  COVALENT BOND -made by sharing electrons -this bond is formed between atoms that have similar eleactronegativities >Nonpolar (CI2, CO2, CCI2) -no sIgnificant difference of electronegativity >Polar (HCl,HCHO) - has significant difference of electronegativity Nonpolar Covalent Bond -this bond is formed between atoms that have similar electronegativities- the affinity or desire for electrons. Polar Covalent Bond -is formed when atoms of slightly differenct electronegativities share electrons. INTERmolecular forces -forces of attraction between molecules -forces of attraction or repulsion which act between neighborin particles(atoms, molecules or ions). -they are weak compared to the intramolecular forces Types: A. Binding forces B. Attractive forces Cohesion - similar molecules van der waals Adhesion - different molecules Hydrogen bond Repulsive - prevent molecules from Ion-dipole annihilating each other Ion-induced Dipole VAN DER WAALS -weak forces that involve the dispersion of charge 1. Keesom Forces (orientation across a molecule called dipole effect) ▪ Dipole-dipole ▪ molecules are polar with permanent polar dipoles ▪ 1-7 kcal/mole ▪ ex. water, HCl, ethanol, acetone, phenol 2. Debye Forces (induction effect) 3. London Forces (dispersion effect) ▪ dipole-induced dipole ▪ induced dipole-induced dipole ▪ transient dipole induced by a ▪ induce polarity between non polar permanent dipole molecules ▪ polar molecules produce temporary ▪ responsible for liquefaction of gases electric dipole in nonpolar molecules ▪ 0.5-1 kcal/mole ▪ 1-3 kcal/mole ▪ ex. carbon disulfide, CCl2, hexane ▪ ex. ethyl acetate, methylene chloride, ether HYDROGEN BOND -electrostatic interaction of H with highly electronegative atoms (N, O, F, S, Cl) -accounts for unusual properties of water ION-DIPOLE INTERACTION - polar molecules are attracted to either positive or negative charges - occurs when salt is dissolved in a polar solvent solubility of crystalline substances in H2O -quaternary ammonium + tertiary amine  ION-INDUCED DIPOLE -induced by close proximity of a charged ion to a non-polar molecule -responsible for the solubility of non-polar molecules -ex. Iodine complex with salts INTERMOLECULAR 1. What forces exist between 2. Which of the following can the following? form H-bonds with water? FORCES OF  HBr - H2S ▪ CH3OCH3 ATTRACTION Cl2 - CBr4 ▪ CH4 I2 - NO3 ▪ F- SAMPLE PROBLEMS NH3 - C6H6 ▪ HCOOH ▪ Na+ PHYSICAL PROPERTIES OF MATTER ADDITIVE depends on the total CONSTITUTIVE contribution of the atoms in the molecules depends on the arrangement of the number and kind of atoms ex. MW, mass within a molecules atoms = MW = Mass ex. refractive index, optical rotation COLLIGATIVE function of the number of ❑Osmotic pressure eleviation species or particles present in a given solution ❑Vapor pressure lowering ❑Freezing point depression ❑Boiling point elevation  TYPES OF PROPERTIES INTENSIVE/INTRINSIC EXTENSIVE/EXTRINSIC independent of the amount of the substances in the system depends on the quantity of ex. Temperature, Pressure, substance in the system Density, Viscosity, Surface tension, Specific gravity ex. Mass, Volume, Lenght DENSITY = mass per unit volume (M/V) SPECIFIC VOLUME= reciprocal of SPECIFC GRAVITY = density of specific gravity, opposite of sample/ density of standard density 1. Pycnometer method 2. Mohr Westphal balance 3. Plummet method (Archimedes principle: buoyancy - weight of immersed is equal to displaced) SAMPLE PROBLEMS DENSITY 3. The density of a substance is 1. A loaf bread has a mass of 500g 4.0g/cm3. If a sample of the and a volume of 2500cm3. What substance has a volume of 25cm3, is the density of the bread? then what is its mass? 2. A block of wood has a mass of 4. You have a lead ball with mass 6.0g and a volume of 12.0cm3. of 420g. The density of lead is What is the density of the block of 10.5g/cm3. What is the volume of wood? the ball? SPECIFIC GRAVITY 1. Calculate the specific gravity of 3. if the sp.gr. of ice is 0.92, then iron. The density of iron is 7850 wis the density of ice? kg/m3. The specific gravity ofers is 1000kg/m3. 2. If the density of gold is 19300kg/m3, what is the sp.gr. of gold? density of water: 1000kg/m2  THE GASEOUS STATE GAS LAWS 1. BOYLES LAW refers to an ideal situation where no intermolecular interactions exist and collisionsare perfectly aka “mariottes law” elastic. states that in a perfect gas where there is no energy exchanged mass and temperature are kept upon collision. constant, the volume of the gas varies inversely with the absolute pressure. formula P1V1=P2V2 2. GAY-LUSSAC LAW 3. CHARLES LAW states that the pressure of a temperature and the volume is given mass of gas varies directly directly proportional with the abolute temperature of pressure is constant the gas, when the volume is kept constant formula formula IDEAL GAS LAW COMBINED GAS LAW ▪ also called the general gas ▪ the ratio oduct of pressure and equation volume and the absolute ▪ is the equation pf state of a temperature of a gas is equal to a hypothetical ideal gas. constant ▪ formula: SAMPLE PROBLES GAS LAW BOYLES LAW Problem #2: If a gas at 25.0 °C occupies 3.60 liters at a pressure Problem #1: A gas occupies 12.3 of 1.00 atm, what will be its liters at a pressure of 40.0 mmHg. volume at a pressure of 2.50 atm? What is the volume when the pressure is increased to 60.0 mmHg? Problem #3: To what pressure must a gas be compressed in order to get into a 3.00 cubic foot tank the entire weight of a gas that occupies 400.0 cu. ft. at standard pressure? CHARLES Problem #1: Calculate the Problem #3: A gas occupies 900.0 decrease in temperature (in mL at a temperature of 27.0 °C. Celsius) when 2.00 L at 21.0 °C is What is the volume at 132.0 °C? compressed to 1.00 L. Problem #4: What change in Problem #2: 600.0 mL of air is at volume results if 60.0 mL of gas is 20.0 °C. What is the volume at cooled from 33.0 °C to 5.00 °C? 60.0 °C?  GAY-LUSAAC’S Problem #1: A 30.0 L sample of nitrogen inside a rigid, metal container at 20.0 °C is placed inside Problem #2: Determine the an oven whose temperature is 50.0 pressure change when a constant °C. The pressure inside the volume of gas at 1.00 atm is container at 20.0 °C was at 3.00 heated from 20.0 °C to 30.0 °C. atm. What is the pressure of the nitrogen after its temperature is Problem #3: A gas has a pressure increased to 50.0 °C? of 0.370 atm at 50.0 °C. What is the pressure at standard temperature? IDEAL GAS LAW Problem #1: Determine the Problem #3: At what temperature volume of occupied by 2.34 grams will 0.654 moles of neon gas of carbon dioxide gas at STP. occupy 12.30 liters at 1.95 atmospheres? Problem #2: A sample of argon gas at STP occupies 56.2 liters. Problem #4: A 30.6 g sample of Determine the number of moles gas occupies 22.414 L at STP. of argon and the mass of argon in What is the molecular weight of the sample. this gas?  COMBINED GAS LAW Problem #1: A gas has a volume of 800.0 mL at −23.0 °C and 300.0 torr. What would the volume of the gas Problem #3: 690.0 mL of oxygen are be at 227.0 °C and 600.0 torr of collected over water at 26.0 °C and a pressure? total pressure of 725.0 mm of mercury. What is the volume of dry oxygen at 52.0 °C and 800.0 mm pressure? Problem #2: 500.0 liters of a gas in a flexible-walled container are prepared at 700.0 mmHg and 200.0 Problem #4: What is the volume of gas °C. The gas is placed into a tank at 2.00 atm and 200.0 K if its original under high pressure. When the tank volume was 300.0 L at 0.250 atm and cools to 20.0 °C, the pressure of the 400.0 K. gas is 30.0 atm. What is the volume of the gas?  KINETIC MOLECULAR gases are composed of THEORY particles called atoms or molecules, the total volume of which is so small as to be the particles exhibit negligible in relation to the continuous random motion volume of the space in owing to their kinetic energy which the molecules are the molecules exhibit confined. perfect elasticity. the particles of the gas do not attract one another, but instead move with complete independence. THE LIQUID STATE BOILING POINT CRITICAL TEMPERATURE the temperature at which the vapor pressure of the liquid equals the external and temperature above which a atmospheric pressure. liquid can no longer exist CRITICAL PRESSURE pressure required to liquefy a gas at its critical temperature, which is also the highest vapor pressure of a liquid can have. THE SOLID STATE have fixed shapes nearly incompressible have strong intermolecular forces very little kinetic energy atoms vibrate fixed positions about an equilibrium position, and so there is very little transitional motion. CRYSTALLINE SOLIDS solids whose structural units are arranged in a fixed 6 Distinct Critical Systems Based geometric pattern or lattices. on Symmetry definite shape ❑CUBIC - NaCl orderly arrangement of units ❑TETRAGONAL - Urea definite and sharp melting ❑HEXAGONAL - Iodoform points ❑MONOCLINIC - Sucrose solvates - aka “pseudopolymorphs” crystals ❑RHOMBIC - Iodine having solvent molecules ❑TRICLINIC - Boric Acid  POLYMORPHISM condition where substances can exist in more than 1 crystalline THEOBROMA oil Polymorphs (MP) form unstable y form - 18*C polymorphs have different ᵅ form - 22*C melting points, x-ray crystals and β prime form - 28*C diffraction patterns and solubility β stable from - 34*C types of polymorphism ❖enantiontropic - reversible ❖monotropic - unidirectional transition POLYMORPHIC CHANGES AND PROPERTIES ENANTIOTROPIC ISOTROPIC - change is reversible - similar (identical) properties in all directions MONOTROPIC - change is unidirectional ANISOTROPIC - different properties in various directions along the crystal  AMORPHOUS CRYSTALLINE random unoriented fixed geometric patterns molecules ice and NaCl glass and plastics orderly arranged units, arranged in random incompressible manner Definite melting point no definite melting point pass sharply from solid to faster dissolution rate liuid FREEZING POINT LATENT HEAT of FUSION temperature at which liquid > solid melting point of a pure crystalline energy absorbed when 1 g of solid compound melt heat liberated when it freezes LE CHATELIER’S PRINCIPLE - states that a system at equilibrium readjusts so as to reduce the effect of an external stress LIQUID CRYSTALLINE STATE (MESOPHASE) Liquid cyrstals- itermediate between liquid and solid states may result from the heating of solids (thermotropic) or form the action of certain solvents on solids (lyotropic liquid crystals) Two main types of Liquid Crystals 1. SMECTIC 2. NEMATIC Soap like or grease like threadlike molecules are mobile in 2 molecules are mobile in 3 directions directions rotates in 1 axis rotates in 1 axis CHOLESTERIC speacial type of nematic  SUPERCRITICAL FLUIDS properties intermediate between those of liquids and gases formed from the gaseous state where the gas is held under a combination of temperature and pressures that exceed the critical point of a substance PHASE EQUILIBRIA AND THE PHASE RULE (GIBSS’S PHASE RULE) relates the effect of the least F = no. of degrees of freedom number of indipendent variables C = no. chemical components (T, P, and C) among the various phases (S,L and G) that can exist P = no. of phases in an equilibrium system X = variable dependent upon containing a given number of coniderations of the phase components diagram formula F=C-P+X F -number of degrees of freedom least number of intensive/independent variables that must be fixed to describe the system completely P - number of homogenous C - number of components physically distinct portion of smallest number of constituents a system that is separated by which the composition of each from other portions of the phase in the system at equilibrium can be expressed in system by bounding surfaces the form of a chemical formula or equation  APPLICATION of the phase rule to single-component system SYSTEM NUMBER OF DEGREES OF PHASES FREEDOM GAS, LIQUID OR 1 F=C-P+2 system is BIVARIANT (F=2) SOLID = 1 -1 + 2 = 2 GAS-LIQUID, 2 F=C-P+2 system is UNIVARIANT (F= 1) LIQUID-SOLOD, =1-2+2=1 OR GAS-SOLID GAS-LIQUID- 3 F=C-P+2 system is INVARIANT (F=0) SOLID =1-3+2=0 TWO COMPONENT SYSTEMS S and L phases only 2 components - liquid phase aka “condensed system” 2 components S & L - eutectic the vapor state is disregarded mixtures with an assumption of working 3 components at a pressure at 1atm TWO COMPONENT SYSTEM CONTAINING TWO LIQUIDS BINODAL CURVE UPPER CONSULATE/ CRITICAL SOLUTION TEMPERATURE - area within the curve which - maximmum temperature at represents a 2-phase system which two phase region in the phase diagram of a two- - any point beyond it is a single component system containing two phase liquids will exist TIE LINE - line from which a system CONJUGATE PHASES separates into phases of constant composition - approximates proportion of - phases of constant composition components in a particular that separate when a mixture is temperature prepared within the boundary of the 2-phase system. TWO COMPONENT SYSTEM CONTAINING SOlID AND LIQUID minimum temperature where both exist in liquid form point where solid A, solid B and the liquid phase co-exist EUTECTIC POINT - is the point where solid A, solid B and the liquid phase co-exist EUTEXIA - 3 phase co-exist - phenomenon of lowering the melting point due to combinations of components (thymol-salol, comphor-menthol) THREE COMPONENT SYSTEM ternary system PHASE RULE: F = C - P 2 liquids are miscible + 3rd component (co-solventt) with APEX - 100% of each component afffinity to both layers BASE - opposite of apex, 0% of has 4 degrees of freedom each component temperature and pressure are both made constant PHASE DIAGRAM represents that states of matter that exist as temperature and pressure are varied. it is a graphic way to summarize the conditions under which equilibria exist between the different states of matter. such a diagram also allows us to predict which phase of a substance is present at any given temperature and pressure. LATENT HEAT/ MOLAR HEAT heat necessary for 1 mole of a gas, solid, or liquid to change to another phase. either gained or lost. without latent heat, no phase transition CHAPTER III THERMODYNAMICS deals with the quantitaitve relationships of interconversion Surroundings - the rest of the of the various forms of energy universe from which the science of the relationship observations are made between heat, work, temperature and energy Boundaries - physical or virtual barriers that separate a system SYSTEM - a well-defined part of from the suroundings the universe under study  TYPES OF SYSTEM OPEN SYSTEM ISOLATED SYSTEMS -energy and matter can be exchanged with the surroundings -neither matter nor energy can be exchanged with the surroundings CLOSED SYSTEM - energy can be exchanged with the surrounding but not matter FIRST LAW OF THERMODYNAMICS law of conservation of energy ENTHALPY energy cannot be created nor - a property of a thermodynamic destroyed, it can only be system, is the sum of the system’s transformed into a different internal energy and the product of from its pressure and volume. formula: H = E + PV E - is the enternal energy P - pressure V - volume  Modified First-Law Equations for processes occuring under various conditions Specified Conditions Process Constant Heat Adiabatic Insulated vessel, such as Dewar Flask Revirsible process at a constant Isothermal constant-temperature bath temperature Ideal gas at a constant Isothermal constant-temperature bath temperature Constant volume Isometric (isochoric) closed vessel of constant volume, such as Bomb Calorimeter Constant pressure Isobaric reaction occuring in an open container at constant (atmospheric) pressure SECOND LAW OF THERMODYNAMICS refers to the probability of the ENTROPY occurence of a process based on -is the measure of disorder in a the tendency of a system to thermodynamic system approach a state of energy equilibruim entropy THIRD LAW OF THERMODYNAMICS the entropy of a pure crystalline substance is zero at absolute zero because the crystal arrangement must show the greatest orderliness at this temperature THE ZEROTH LAW OF THERMODYNAMICS When two systems are each in thermal equilibrium with a third system, the first two systems are in thermal equilibrium with each other. This property makes it meaningful to use thermometers as the thirds system and to define a temperature scale. CHAPTER IV AND V NONELECTROLTYES AND ELECTROLYTES LIQUIDS ▪ possess less kinetic energy that gases ▪ occupy a definite volume ▪ take the shape of the container ▪ incompressible SYSTEM -is generally considered to be a bounded space or an exact CLASSIFICATION OF DISPERSED quantity of a material substance SYSTEM True solution Colloidal dispersion Coarse dispersion TRUE SOLUTION COLLOIDAL DISPERSION -a mixture of two or more -may be considered as a two-phase compoonents that form a (heterogenous) system under some homogenous molecular dispersions circumstances. - one-phase system -1nm - 0.5nm - 0.5nm -two common pharmaceutical dispersions are emulsion and suspension  SOLUTE SOLVENT EXAMPLE TYPES OF SOLUTION Gas Gas Air Liquid Gas Water in oxygen Solution can be Solid Gas Iodine vapor in air classified according to Gas Liquid Carbonated water the states in which the solute and solvent Liquid Liquid Alcohol in water occur Solid Liquid Aqueous NaCl solution Gas Solid Hydrogen in palladium Liquid Solid Mineral oil in paraffin Solid Solid Gold- Silver mixture, mixture of alums CONCENTRATION EXPRESSION MOLALITY MOLARITY - no. of moles of solute in 1 kg of - no. of moles of solute in 1L of solvent solution  PERCENTAGE STRENGHT - signifies the no. of grams of solute per NORMALITY 100 g of solution -is the number of grams equivalent per liter of the solution. MOLE FRACTION (X,N) - ratio of the moles of one constituents (ex. the solute) of a solution to the total MOLE PERCENT moles of all constituents - moles of one constituent in (solvent and solute) 100 moles of the solution; mole percent is obtained by multiplying mole fraction by 100 PERCENT BY WEIGHT (%w/w) - grams of solute in 100 g of solution PERCENT WEIGHT-IN- VOLUME (%w/v) -grams of solute in 100 ml of PERCENT BY VOLUME (% v/v) solution - milliliters of solute in 100ml of solution ELECTROLYTES form ions in solution STRONG ELECTROLYTES electrical conductance completely ionized in solution shows apparent “anomalous” NaCl, HCl, H2SO4 colligative properties, that is they produce a considerably greater freezing point WEAK ELECTROLYTES depression and boiling point elevation partial ionization CH3COOH, ephedrine and ex. HCl, sodium sulfate, phenobarbital ephedrine, and phenobarbital IONIZATION ANODE - positively charge ELECTRODE -is the process by which an atom or a molecule acquires a negative CATHODE or positive charge by gaining or - negatively charge ELECTRODE losing electrons, often in conjunction with other chemical changes. The resulting electrically ANions charged atom or molecule is - negatively charge IONS called ION. CATions - positively charge IONS NON-ELECTROLYTES substances that do not ionize when dissolved in water and therefore do not conduct an electric current through solution. ex. glycerin, naphthalene, urea, sucrose  COLLIGATIVE PROPERTIES OF SOLUTIONS VAPOR PRESSURE LOWERING RAOULTS LAW -lowering of a vapor pressure f a -the addition of a non-volatile solute solvent is equal to the product of lowers the VP of a liquid the mole fraction of the solute - a liquid in a closed container will and vapor pressure of the solvent establish an equilibrium with its vapor - when equilibrium is reached, vapor exerts a pressure ()vapor pressure Temperature is constant VOLATILE - exhibits VP NONVOLATILE - no measurable VP BOILING POINT ELEVATION BP -temperature at which liquid pressure is equal to atmospheric pressure (1atm = 760mmHg) The boiling point of a solution containing a non-volatile solute would be higher that the pure solvent because the solute would lower the vapour pressure of the solvent  FREEZING POINT DEPRESSION FP- temperature at which solid and liquid phases are in eqlibrium under an external pressure In general, solutions have lower freezing point than the pure solvent Applications: salt is spread on roads to melt ice ethylene glycol as “snti-freeze” OSMOTIC PRESSURE OSMOSIS - movement of water across a semipermeable membrane from low to high conceentration. This is the pressure required to offset the movement of solvent thru a semipermeable membrane. Also defined as the pressure required to prevent osmosis in solutions. CHAPTER VI IONIC EQUILIBRIA SOLVENTS PROTOPHILIC OR BASIC SOLVENTS phase of the solution, usually -proton accepting (acetone, ehter, constitutes the largest proportion of and liquid ammonia) the system. types: PROTOGENIC SOLVENTS protophilic or basic solvent -proton-donating (formic acid, acetic protogenic solvent acid, sulfuric acid, liquid HCl, and amphiprotic solvent liquid HF) aprotic solvents AMPHIPROTIC SOLVENTS -act as both (water and alcohols) DISSOLUTION APROTIC SOLVENTS - neither accept or donate -transfer of molecules or ions protons, neutrals (hydrocarbon) from a solid state into solution  THEORIES OF ACIDS AND BABSES THEORY ACID BASE IONIZATION OF WEAK ACIDS ARRHENIUS liberates Liberates OH AND BASES H20 in in aqueous aqueous solution solution IONIZATION - complete separation of ions in a crystal LEWIS E acceptor E donor lattice when a salt is dissolved BRONSTED- P donor P acceptor LOWRY DISSOCIATION - separation of ions in solution when the ions are associated by interionic attraction SOLUBILITY concentration of a standard solution in which the dissolved solute is in equilibrium with its Factors Affecting Solubility solid phase at constant temperature Intrinsic Solubility pH Apparent Solubility addition of salt Kinetic Solubility pressure (Henry’s Law) Thermodynamic Solubility salt formation particle zie Temperature ▪ ENDOTHERMIC DISSOLTION - heat is absorbed, > temp, > Salts Formation soLubility ▪ SALTING-IN - a salts increases hydrophilicity of ▪ EXOTHERMIC DISSOLUTION the solution - heat is released, < temp, >solubility (CaOH) ▪ SALTING-OUT - added salt reduces the available amount of water > solute precipitates  - the negative logarithm of the H+ pH concentration pH = -log [H+] CHAPTER VII BUFFERS AND ISOTONIC SOLUTION compound or a mixture of compounds whcih has the BUFFER CAPACITY ability to resist changes in pH - buffer efficiency or buffer index when small amounts of acids and bases are added this property results from the B = represents the small presence of a buffer pair increment in gram equivalents per which consist f either: liter of strong base or acid added to the buffer solution to produce a - weak acid and some salt or its change in pH conjugate base - weak base and some salt or its conjugate acid HENDERSON-HASSELBALCH EQUATION aka pH or buffer equation WEAK ACIDS preparationof drug solutions at a desired pH using both the neutral and the salt forms of a drug determine percentage of neutral WEAK BASES and ionized forms at a given pH determination of pKa of an acid or a base BUFFERED ISOTONIC SOLUTIONS Pharmaceutical solutions that are meant for application delicate Isotonic solutions cause no membranes of the body should swelling or contraction of the also be ajusted to approximately tisseus with which they come in the same osmotic pressure that of contact and produce no the body fluids. discomfort when instilled in the eye, nasal tract, blood or other body tissues. TONICITY OF SOLUTIONS ISOTONIC SOLUTIONS - living cell does not gain or loss HYPERTONIC SOLUTIONS water - sameosmotic pressure with body - more solutes compared to cell fliuds concentrations - 0.9% NaCl solution, normal - freeze lower than 0.52 *C saline, D5W - cauases creantion of the cell - 5% NaCl solution  HYPOTONIC SOLUTIONS - less solutes compared to cell concentrations - freeze higher than -0.52*C - causes lysis of the cell - distilled water TONICITY - is the concentration of the solutes that cannot cross the OSMOLALITY and membrane since these solutes OSMOLARITY exert an osmotic pressure on that membrane. - are colligative properties that measure the concentration of the solutes Tonicity is not the difference independently of their between the two osmolarities abilityto cross a cell on opposing sides of the membrane membrane. MEASUREMENT OF TONICITY 1. HEMOLYTIC METHOD - the effect of various solutions of the drug is A quantitative method observed on the developed by Hunter was appearance of red blood used based on the fact that a hypotonic solutions cells suspended in the liberates oxyhemoglobin solutions. in direct proportion to the number of cells hemolyzed. 2. Measurement of the slight temperature differences 3. CALCULATING TONICITY using -the second method is based Liso values on a arising measurement of - because freezing point the slight temperature depressions for solutions of differences from differences in electrolyts of both the weak and the vapor pressure of thermally strong types are always greater than those calculated from insulated samples contained in equation ΔTf= Kfc, a new factor, L constant-humidity chambers. = i Kf, is introduced to overcome this difficulty. The L value can be obtain from the freezing oint lowering of the solutions of representative compounds of a given ionic type Example. at acentration c that is isotonic with body fluids. The specific The Liso value for a 0.9% (0.154M) value of L is writtn as Liso solution of sodium chloride, which has a freezing point depression of 0.52*C and is isotonic with body fluids, is?  METHODS OF ADJUSTING TONICITY AND pH CLASS 1 METHODS - NaCl or some other substances is added to the solution of the drug to make it isotonic 1.Freezing Point Depression/Crysoscopic Method - FPD used to calculate the amount of solute to add in making an isotonic solution 2. Sodium Chloride Equivalent Method/ E VALUE E value - gram of NaCl equivalent to 1 gram of a substance Step 3. Subtract step 1 from step 2 STEPS: Step 4. if agent other than Step 1. Calculate the amount of NaCl (boric acid, dextrose, Na NaCl represented by the or K nitrate), divide the ingredients amount of NaCl (step 3) by Step 2. Calculate the amount of the NaCl equivalent (E value) NaCl that would make the volume of the other substance of solution specified isotonic  CLASS II METHODS - water is added to the drug > isotonic solutions White Vincent Method V = wt x E x 111.1 Sprowls Method V = 0.3g x E x 111.1 CHAPTER VIII SOLUBILITY AND DISTRIBUTION PHENOMENA SOLUBILITY - is defined as a concentration SOLUTION of solute in a saturated solution at certain temperature. It is - is a system in which intrinsic property of solute. molecules of a solute are dissolved in a solvent. - the spontaneous interaction of two or more substance to form a homogenous molecular dispersion. Solubility canbe categorized as: 1. Unbuffered UNBUFFERED SOLUBILITY 2. Buffered 3. Intrinsic - typically measure in water, describes the solibility of a solution saturated with a compoud at the terminal pH of a solution. BUFFERED SOLUBILITY -apparent solubility - describes the solubility at a specific pH. Solubility is INTRINSIC SOLUBILITY dependent on the concentration and nature of - refers to the solubilityof an ions within the solvent. ionizable compound in its neutral form. This parameter that is typically independent of the ionic properties of the medium used. SATURATED SOLUTIONS -the solute is in equilibrium with the solid phase SUPERSATURATED SOLUITON - solute concentration is above necessary or it contains UNSATURATED SOLUTION more of the dissolved solute than it would normally - or subsaturated solution, contain in a definite solute concentration is below temperature necessary for complete saturation at a definite temperature  SOLUBILITY DEFINITION IN THE USP VERY SOLUBLE less than 1 part FREELY SOLUBLE 1-10 SOLUBLE 10-30 SPARINGLY SOLUBLE 30-100 SLIGHTLY SOLUBLE 100-1,000 VERY SLIGHTLY SOLUBLE 1,000-10,000 PRACTICALLY INSOLUBLE >10,000 SOLUTE-SOLVENT INTERACTIONS “LIKE DISSOLVES LIKE” - polar substances tto dissolve in polar solvents, while nonpolar TYPE OF SOLVENTS: subrstances tend to dissolve in nonpolar solvents. 1. Polar solvents - the more similar the 2. Nonpolar solvents intermolecular attractions are, the 3. Semipolar solvents more likely one substance is t be soluble in another. POLAR SOLVENTS Polar solvents dissolve ionic substances and polar substances. The solubility of a drug in polar Water dissolves phenols, alcohols, solvents depends on aldehydes, ketone amines, and the polarity of the solute and other oxygen- and nitrogen- the solvent. containing compounds that can form hydrogen bonds with water. the ability of the solute to form hydrogen bonds  the ratio of polar to nonpolar groups of the molecule. - when additional polar groups are present in the molecule (e.g., propylene glycol, glycerin, and tartaric acid), water solubility increases greatly. As the length of a nonpolar Branching of the carbon chain chain of an aliphatic alcohol reduces the nonpolar effect and increases, the solubility in leads to increased water water decreases (e.g. straight solubility (e.g. tertiary butyl chain monohydroxy alcohols, alcohol is miscible in all aldehydes, and acids with proportions with water, more than 4 carbons cannot whereas n-butyl alcohol enter into the hydrogen- dissolves only to a small extent.) bonded structure of water and hence are slightly soluble.  NONPOLAR SOLVENTS Ionic and polar solutes are not soluble or are only slightly in nonpolar solvents (e.g. hydro- Nonpolar solvents cannot break carbons) because: covalent bonds and ionize weak Nonpolar solvents are unable electrolytes, because they are to reduce the attraction aprotic (no hydrogen). between the ions of strong and weak electrolytes because of Nonpolar solvents cannot form the solvents' low dielectric hydrogen bridges with constants. nonelectrolytes. SEMI-POLAR SOLVENTS ▪ semipolar solvents induce certain degree of polarity in nonpolar solvents ▪ semipolar solvents acts as intermediate solvents that generate miscibility between polar and non-polar liquids (e.g. acetone increase the solubility of ether in water)  IDEAL AND REAL SOLUTIONS IDEAL SOLUTONS REAL SOLUTIONS  ideal solution is one in which  in real solutions the “cohesive” force there is no attraction between of attraction between A and B exceeds solute and solvents molecules the “adhesive” force of atttraction ideal solution is one in which existing between A and B. there is no change in the  alternatively, the attractive forces properties of the components, between A and B may be greater that other than dilution those between A and A or B and B. they obey Raoult’s Law  this may occur even though the liquids form solution in all proportions. Such mixtures are real or non-ideal  they do not obey Raoult;s Law  RAOULT’S LAW ▪ in ideal solutions partial vapor ▪ in which ƤA and ƤB are the pressure of each volatile constituent is equal to the vapor partial vapor pressures of the pressure of the pure constituent constituents over the multiplied by its le fraction in the solution when the mole solution. fraction concentrations are ▪ thus, for two constituents A and B Xa and Xb respectively. Vapor pressure composition curve for an ideal solution (binary mixture) AZEOTROPIC BINARY MIXTURE it is the mixture of liquids that has a constant boiling point because the vapour the components of the has the same composition solution cannot be as the liquid mixture. separated by simple the boiling point of an distillation. azeotropic mixture may be higher or lower than that of any its compenents. SOLUBILITY OF LIQUID IN LIQUID Frequently two or more liquids are mixed together in the preparation of Liquid-liquid systems can be pharmaceutical products (e.g divided into two categories aromatic waters, spirits, according to the solubility of elixirs, lotions, aprays, and the substances in one medicated oils.) another: 1. Complete miscibility 2. Partial miscibility COMPLETE MISCIBILITY polar and semipolar solvents, such as water and alcohol, alcohol and acetone, are said to be completely miscible These liquids are miscible because they mix in all because the broken attractive proportions. forces in both pure liquids are re-established in the mixture. Nonpolar solvents such as benzene and CCl4 are also completely miscible  PARTIAL MISCIBILITY When water and phenol are mixed, two liquid layers are formed each containing some of the other liquid in the dissolved state. It is possible to calculate the composition of each component in the two conjugate phases and the relative amount of each phase from the tie lines that cut the binodal curve. Partially miscible liquids are influenced by temperature. The two conjugate phases changed to a homogenous single phase at the critical solution temperature (or upper consolute temperature.) Some liquid pairs (e.g. trimethylamine and water) exhibit a lower consolute temperature, below which the two members are soluble in all proportions and above which two separate layers form. EUTECTIC MIXTURE: SALOL + THYMOL PHASE DIAGRAM SALOL-THYMOL SYSTEM ▪ a single liquid phase if salol-thymol combinations are ▪ a region containing solid salol to be dispensed as a dry powder, and a conjugate liquid phase the ambient temperature must be below its eutectic point of 13*C. ▪ a region in which solid thymol is Above this temperature it exists in equilibrium with a conjugate in liquefied form. liquid phase ▪ a region in which both components are present as pure solid phases 3 COMPONENT SYSTEM DIAGRAM TERNARY DIAGRAM SOLUBILITY OF GAS IN LIQUID Solubility of gas in liquids depends on: ▪ the mass of gas molecules ▪ Pressure ▪ Temperature ▪ Presence of Salts ▪ Chemical reactions with solvent MASS OF GAS MOLECULES the solubility of gas PRESSURE molecules typically increase with increasing mass of the gas molecules.  gases increase in solubility  the larger the mass of gas with an increase in pressure molecules, the stronger  increasing the pressure London and Debye forces is results in more collisions of between gas and solvent the gas molecules with the molecules surface of the solvent (more solvation); hence greater solubility. HENRY’S LAW the effect of the pressure on the the partial pressure of the gas is solubility of a gas is expressed by obtained by subtracting the Henry’s Law vapour pressure of the solvent “in a very dilute solution at from the total pressure above constant temperature, the the solution. concentration of dissolved gas is if C2 is the concentration of the proportional to the partial dissolved gas in grams per liter pressure of the gas above the of solvent and p is the partial solution at equilibrium” pressure in ml of the undissolved gas above the solution, TEMPERATURE ▪ gases decrease in solubility with an increase in temperature ▪ increasing temperature causes an increase in kinetic energy of PRESENCE OF SALTS gas molecules which leads to breakdown of intermolecular dissolved gases are often bonds and gas escaping from liberated from solutions by solution. the introduction of an ▪ e.g. carbon dioxide gases escape electrolyte (e.g. NaCl) and faster from a carbonated drink sometimes by a non as the temperature increases. electrolyte (e.g. sucrose) this phenomenon is known as SALTING OUT pH ꙳systems of solids in liquids ꙳acidic drugs (e.g. NSAIDS) include the most frequent are more soluble in basic and important type of solutions where the ionized pharmaceutical solutions form is the predominant ꙳most drugs belong to the ꙳basic drugs (e.g. ranitidine) class of weak acids and are more soluble in acidic bases. They react with solutions where the ionized strong acids or bases to form form of the drug is water soluble satls. predominant.  SUBSTITUENTS ▪ substituents can influence solubility by affecting the solute molecular cohesion and its interaction with water molecules. ▪ polar groups scuh as -OH are capable of hydrogen binding (high solubility) ▪ e.g. hydroxy acids, such as tartaric and citric acids, are quite soluble in water because they are solvated through their hydroxyl groups ▪ Non-polar groups such as - CH3 and -Cl are hydrophobic (low solubility) ▪ e.g. the OH group of salicylic ▪ the position of the acid cannot contribute to the sustituent on the molecule solubility because it is involved can effect the solute in an intramolecular hydrogen molecular cohesion and its bond. interaction with water molecules, and hence its solubility. SOLVENT frequently solute is more soluble in mixture of solvents than single solvent the solvent, which in CRYSTAL CHARACTERISTICS combination with the main ▪ different crystalline forms of the solvent increase solubility is same substance (polymorphs) known as cosolvent possess different lattice energies ▪ the polymorphic form with the lowest free energy will be the most stable. ▪ less stable (metastable) forms with the highest energy will be the most soluble one. They tend totransform into the most stable form over time. ▪ the solubility of a crystalline material and its rate of dissolution can be increased by using a metastable polymorph ▪ the interaction between the ▪ many drugs exhibit polymorphish, substance and water that occurs e.g. barbiturates and i cyrstal phase reduces the sulphonamides amount of energy liberated ▪ incorporation of solvent molecules when the solid interact with the into the lattice structure of solvent cyrstalline material during ▪ therefore unsolvated crystals will crystallization will result in solids dissolve faster. that are called solvates. If water is the solvating molecule, the solids are called hydrate COMPLEXATION complexation can increase or e.g. tetracycline can form water decrease the solubility of drugs insoluble complexes with various depending whetehr the formed metal cations. Therefore complex is water soluble or tetracycline solubility is insoluble. decreased in the presecne of e.g. cyclodectrin can form water those metals. soluble complexes with most drugs, thus increasing their water solubility. BOILING AND MELTING POINT ▪ in general, aqueous solubility decreases with increasing boiling and melting point PARTICLE SIZE ▪ this is because the higher the ❑solubility increase with boiling point of liquids and decreasing particle size, due melting point of solids, the to the increased particle stringer the interactions surface area, meaning more between the molecules in of the solid is in contact with the pure liquid or the solid the solvent state. PARTITION COEFICIENT if a liquid or solid substance is added to a mixture of two immiscible liquids, it will become the equilibrium constant, K, distributed between the two is known as the distribution layers inefiinite concentration ratio, dirstribution coeficient ratio. or partition coeficient if C1 and C2 are the equilibrium concentrations of the substance in solvent1 and solvent2, respectively, the equilibrium expression becomes C1 / C2 = K ▪ Partition coeficinet (P) is a parameter that characterizes ▪ determination of P (or log P) the relative affinity of a values involves the placing of compound in its unionized a drug compound along with form for water and an the two immiscible solvents in immiscible model lipid solvent a separation funnel (octanol) ▪ molecules of the solute will ▪ octanol was chosen as the distribute in each phase until model lipid phase because it equilibrium is established most closely stimulates the ▪ the ratio of the two properties of biological concentration is the partition membranes coefficient or distribution coefficient P, P = Co/Cw. INTERPRETATION P>1 or log P> 0 implies that the drug has affinity for lipid membranes ▪ structure affect the value of P=1 or log P=0 there is equal partition coefficinet distribution between the ▪ the substituent that increase water and oil layer P value are -alkyl, -aryl, and P