Lecture 1.pdf

Physical Pharmacy Lecture One Dr. Safaa Fakher Ground Rules Be on time Way of contact Cell phone Evaluation Syllabus Introduction to physical pharmacy State of Matter Two component systems containing liquid phases. Three component systems. Tie linear for three componen...

Physical Pharmacy Lecture One Dr. Safaa Fakher Ground Rules Be on time Way of contact Cell phone Evaluation Syllabus Introduction to physical pharmacy State of Matter Two component systems containing liquid phases. Three component systems. Tie linear for three component systems. Expression of concentrations in pharmaceutical preparations. Ionic Equilibrium Buffer solutions Some basic definitions Physical pharmacy has been associated with the area of pharmacy that dealt with the quantitative and theoretical principles of physicochemical science as they applied to the practice of pharmacy. Physical pharmacy attempted to integrate the factual knowledge of pharmacy through the development of broad principles of its own, and it aided the pharmacist and the pharmaceutical scientist in their attempt to predict the solubility, stability, compatibility, and biologic action of drug products. A dosage form is the entity that is administered to patients so that they receive an effective dose of a drug. Drug delivery system is outlined as a technological formulation or a tool that permits the introduction of a drug molecule within the body by controlling the rate, extent and the location of drug release. Binding Forces Between Molecules Repulsive and attractive forces Intramolecular forces Intermolecular forces Bond energy 5  Repulsive and Attractive Forces For molecules to exist as aggregates in gases, liquids, and solids, intermolecular forces must exist These intermolecular forces involve both attractive and repulsive forces. These forces must be balanced in an energetically favored arrangement for the molecules to interact. 6 Repulsive and Attractive Forces When two atoms or molecules are brought closer together, the opposite charges and binding forces in the two molecules are closer together than the similar charges and forces, causing the molecules to attract one another. When the molecules are brought so close that the similar charges touch, they repel each other like rigid elastic bodies. At a certain equilibrium distance, (about 3–4 Å), the repulsive and attractive forces are equal. At this position, the potential energy of the two molecules is a minimum and the system is most stable. Repulsive and Attractive Forces Too close together, repulsive forces dominate Too far apart, weak attractive forces, no bond Atoms at right distance for stable bond 8  Intramolecular forces Atoms within a molecule are attracted to one another by sharing of electrons, it can be categorized into: 1. Ionic bonding: 2. Covalent bonding: 3. Metallic bonding: between positively and between atoms that share between positively charged ions negatively charged ions the outer shell of electrons. and delocalized outer electrons of (metals and non-metals). metal elements, it is weaker than both ionic and covalent bonding. 9  Intermolecular forces Intermolecular forces occur between molecules. It can be divided into: 1. Van der waals forces 2. Ion–dipole interaction 3. Ion–induced dipole interaction 4. Hydrogen bonds 10  Intermolecular forces Van der Waals Forces Van der Waal interactions are weak forces that involve the dispersion of charge across a molecule called a dipole. Van der Waal interactions can be classified into: Dipole–dipole interaction, orientation effect, or Keesom force Dipole-induced dipole interaction, induction effect, or Debye force Induced dipole–induced dipole interaction, dispersion effect, or London force 11  Intermolecular forces Van der Waals Forces: Keesom forces Keesom forces occur between polar molecules in which the permanent dipoles interact with one another (dipole- dipole interactions) or (orientation effect). Polar molecules have polar covalent bonds which are unevenly distributed in space due to the difference in the electronegativity of the atoms forming the bond. e.g. HCl. The nucleus of the chlorine atom pulls the electron pair involved in the chlorine–hydrogen bond closer to itself and creates a permanent partial positive charge on the hydrogen and a permanent partial negative charge on the chlorine (Permanent dipole), 12  Intermolecular forces Van der Waals Forces: Keesom forces The Partial opposite charges (permanent dipoles) attract one another (dipole-dipole interactions) δ+ δ- δ+ δ- Keesom forces The dipole-dipole forces increases as the polarity of the molecule increases. Keesom forces are much weaker than ionic bonds because the charges involved in bonding are partial. 13  Intermolecular forces Van der Waals Forces: Debye forces Debye forces occur between a polar and a nonpolar molecule in which the permanent dipole in the polar molecule induce an electric dipole in the nonpolar one (dipole-induced dipole interactions) or (induction effect). e.g. water and oxygen Oxygen molecule Water molecule (nonpolar) (polar; permanent dipole) 14  Intermolecular forces Van der Waals Forces: Debye forces The oxygen molecule is nonpolar. However, when it comes close to the oxygen atom in a water molecule, the partial negative charge on the oxygen atom repels the electrons in the oxygen molecule. This causes (induces) a temporary partial positive charge in the end closest to the water molecule and a buildup of a partial negative charge in the end furthest away. This induced dipole is temporary and forms only when the two molecules are extremely close to one another. 15  Intermolecular forces Van der Waals Forces: Debye forces The strength of Debye forces increases with the ease of distortion of the electron cloud of the nonpolar molecule (i.e. polarizability of the molecule) Debye forces is weaker than Keesom forces because the dipole in the nonpolar molecule is temporary (induced) and forms only when the two molecules is extremely close to each other. 16  Intermolecular forces Van der Waals Forces: London forces London forces occur between two nonpolar (neutral) molecules in which molecules can induce polarity on each other (induced dipole-induced dipole interactions) or (dispersion effect). e.g. Helium 17 Intermolecular forces Van der Waals Forces: London forces Two Helium atoms are nonpolar and posses no dipoles. The average distribution of electrons around each nucleus is spherically symmetrical At a given instant in time, the distribution of electrons around an individual atom may not be perfectly symmetrical (the electron cloud may be on one side of the nucleus). This atom will have an instantaneous dipole (temporary partial charges) causing a neighboring atom to distort due to the electrostatic attractions/repulsions of their electron clouds Attraction between opposite partial charges of neighboring induced dipoles cause atoms to stick together for a very short time 18  Intermolecular forces Van der Waals Forces: London forces London forces occur between all atoms and molecules (between polar/polar and polar/nonpolar molecules as well) The larger the atom or molecule, the greater its polarizability (easier to induce a momentary dipole) and the stronger the dispersion forces become. This is because: The electrons are farther from the positive nucleus and so are held less strongly The number of electrons is greater London Force is the weakest of all the intermolecular forces. 19  Intermolecular forces Ion-dipole forces Ion-dipole forces occur between a charged ion and a polar molecule (i.e. a molecule with a dipole) Cations are attracted to the negative end of a dipole, while anions are attracted to the positive end of a dipole 20  Intermolecular forces Ion-dipole forces These types of interactions account in part for the solubility of ionic crystalline substances in water; the cation, for example, attracts the relatively negative oxygen atom of water and the anion attracts the hydrogen atoms of the dipolar water molecules. 21  Intermolecular forces Ion-induced dipole forces Ion-Induced dipole forces occur between a charged ion and a nonpolar molecule. Theses forces result when the approach of an ion induces a dipole in an atom or in a nonpolar molecule by disturbing the arrangement of electrons Ion-induced dipole forces are presumably involved in the formation of the iodide complex I2 + K+I- = K+ I3 - 22  Intermolecular forces Hydrogen bond Hydrogen bond is a strong type of dipole-dipole interaction that occurs between a molecule containing a hydrogen atom and a strongly electronegative atom such as fluorine, oxygen, or nitrogen In order to create the bond, the hydrogen atom must be covalently attached to another electronegative atom. A perfect example of hydrogen bond is water. Hydrogen bonds can also exist between alcohol molecules, carboxylic acids, aldehydes, esters, and polypeptides. 23  Intermolecular forces Hydrogen bond In a water molecule, the hydrogen ion electron is pulled by the covalently attached oxygen atom, creating a naked nucleus on the side of the hydrogen atom facing away (partial positive charge). This side of the hydrogen atom can get close to neighboring oxygen atom (with a partial negative charge) and interact strongly with it. 24  Intermolecular forces Hydrogen bond Hydrogen bonds are responsible for many unusual physical properties of water including its abnormally low vapor pressure, high boiling point, and the greater volume of ice water. Hydrogen bonding is stronger than all Van der Waals intermolecular forces, but are still weaker than ionic and covalent bonds. 25  Bond energy Bond energy is a measure of bond strength. It is the heat required to break one mole of molecules into their individual atoms The relative strength of forces from strongest to weakest: Intramolecular forces Intermolecular forces 1. Ion-dipole forces 1. Ionic 2. Ion-induced dipole 2. Covalent 3. Hydrogen Bond 3. Metallic 4. Van der Waals’ forces (Keesom forces > Debye forces > London forces) 26  States of matter 1. Gaseous state 2. Liquid state 3. Solid and crystalline state 4. Liquid crystalline state 27 The gaseous state Gas general properties Ideal gas Real gas 28 Gas general properties Gases can be expanded infinitively; therefore, gases can fill containers and take their volume and shape. Gases diffuse and mix evenly and rapidly. Gases have much lower densities than liquids and solids (There is a lot of free space in a gas, therefore; It is the most compressible state of matter). 29  Gas general properties Gas molecules travel in random paths and collide with one another and with the walls of the container in which they are confined A gas exerts a pressure (a force per unit area) expressed in dynes/cm2, atmospheres or in mmHg (1 atm = 760 mmHg = 760 Torr). Gases have volumes that is expressed in liters or cubic centimeters (1 cm3 = 1 mL). The temperature involved in the gas equations is expressed by the absolute or Kelvin scale (0°C=273.15 K (Kelvin)). 30  Ideal gas Ideal gas is a gas where no intermolecular interactions exist and collisions are perfectly elastic, and thus no energy is exchanged during collision. The properties of the ideal gas can be described by the general ideal gas law, which are derived from Boyle, Charles and Gay- Lussac laws 31  Ideal gas Boyle’s law Boyle’s law states that the volume and pressure of a given mass of gas is inversely proportional (i.e. when the pressure of a gas increases, its volume decreases). P α 1/V or P =K/V P1V1 = P2V2 P: pressure, K: constant, V: volume  Ideal gas Charles law Charles law states that the volume and absolute temperature of a given mass of gas at constant pressure are directly proportional (i.e when the temperature of a gas increases, its volume increases as well). V α T or V = k T 𝑽𝟏 𝑽𝟐 = 𝑻𝟏 𝑻𝟐 T: temperature in Kelvin 33  Ideal gas Gay-Lussac law The law of Gay-Lussac states that the pressure and absolute temperature of a given mass of gas at constant volume are directly proportional (i.e when the temperature of a gas increases, its pressure increases as well). P α T or P = k T 𝑷𝟏 𝑷𝟐 = 𝑻𝟏 𝑻𝟐 34  Ideal gas Combined gas law Boyle, Gay-Lussac and Charles law can be combined to obtain the familiar relationship: 𝑷 𝟏 𝑽𝟏 𝑷 𝟐 𝑽𝟐 = 𝑻𝟏 𝑻𝟐 35  Ideal gas Combined gas law: Example 1 A gas occupies a volume of 30.0 mL at a temperature of 20°C and a pressure of 740 mm Hg. Assuming the gas is ideal, what is the volume at 0°C and 760 mm Hg? 𝐏𝟏 𝐕 𝟏 𝐏𝟐 𝐕 𝟐 = 𝐓𝟏 𝐓𝟐 740 × 30 760 × 𝐕𝟐 = 𝐕 = 27.2 mL 273 + 20 273 36  Ideal gas Combined gas law: Example 2 A sample of methane CH4 has a volume of 7.0 dm3 at a temperature of 4°C and a pressure of 0.848 atm. Calculate the volume of methane at a temperature of 11°C and a pressure of 1.52 atm. 𝐏𝟏 𝐕 𝟏 𝐏𝟐 𝐕 𝟐 = 𝐓𝟏 𝐓𝟐 0.848 × 7 1.52 × 𝐕𝟐 = 𝐕 = 4 dm3 273 + 4 273 +11 37  Ideal gas General ideal gas law General ideal gas law (also called equation of state) relates the specific conditions, that is, the pressure, volume, and temperature of a given mass of gas. 𝐏 𝟏 𝐕𝟏 𝐏 𝟐 𝐕𝟐 𝐏𝐕 = ⟹ =R 𝐓𝟏 𝐓𝟐 𝐓 R: the molar gas constant value for the PV/T ratio of an ideal gas. For n moles it becomes: PV = nRT 38  Ideal gas General ideal gas law: Molar gas constant The volume of 1 mole of an ideal gas under standard conditions of temperature and pressure (i.e., at 0°C and 1 atm) has been found by experiment to be 22.414 liters. Substituting this value in general ideal gas law: 𝐏𝐕 𝟏 ×𝟐𝟐.𝟒𝟏𝟒 R= = = 0.08205 atm L/mole K 𝐓 𝟐𝟕𝟑.𝟏 The molar gas constant can also be expressed by energy units: R = 8.314 Joules/mole K or R = 8.314 × 107 erg/mole K 1 joule = 107 erg 39  Ideal gas General ideal gas law: Example What is the volume of 2 moles of an ideal gas at 25°C and 780 mm Hg? 780 mmHg / 760 mmHg = 1.0263 atm 25 °C + 273 = 298 K PV = nRT nRT 2 × 0.08205 × 298 𝐕= = = 47.65 L P 1. 062 40  Ideal gas General ideal gas law: Molecular weight The approximate molecular weight of a gas can be determined by use of the ideal gas law: PV = nRT since n = g/M Then 𝐠 PV = RT M 𝐠𝐑 M= 𝐏𝐕 41  Ideal gas General ideal gas law: Molecular weight Example If 0.30 g of ethyl alcohol in the vapor state occupies 200 mL at a pressure of 1 atm and a temperature of 100°C, what is the molecular weight of ethyl alcohol? 100 °C + 273 = 373 K 200 mL ÷ 1000 mL = 0.2 L 𝐠𝐑𝐓 0.3 × 0.082 ×373 g 𝐌 = = = 46 𝐏𝐕 1 ×0.2 mole 42  Ideal gas Kinetic Molecular Theory Kinetic molecular theory explains the behavior of gases according to the ideal gas law: 1. Gases are composed of particles called atoms or molecules, the total volume of which is so small (negligible) in relation to the volume of the space in which the molecules are confined. 2. Gas molecules exert neither attractive nor repulsive forces on one another. 43  Ideal gas Kinetic Molecular Theory 3. The particles exhibit continuous random motion owing to their kinetic energy. The average kinetic energy, E, is directly proportional to the absolute temperature of the gas, or E = (3/2) RT. 4. The molecules exhibit perfect elasticity; there is no net loss of speed or transfer of energy after they collide with one another and with the walls of the confining vessel. 44  Real gas Real gases do not interact without energy exchange, and therefore do not follow the laws of Boyle, Charles, and Gay-Lussac. Real gases are not composed of infinitely small and perfectly elastic non-attracting spheres. They are composed of molecules of a finite volume that tend to attract one another. The significant molecular volume and the intermolecular attractions between gas molecules affect both the volume and the pressure of a real gas respectively. 45  Real Gas Van der Waals Equation The van der Waals equation is a modified ideal gas equation that takes into account the factors that affect the volume and pressure of a real gas. 𝒂 For 1 mole of gas: 𝑷+ 𝑽𝟐 𝑽 − 𝒃 = 𝑹𝑻 𝟐 For n moles of gas: 𝑷 + 𝒂𝒏 𝑽𝟐 𝑽 − 𝒏𝒃 = 𝒏𝑹𝑻 The term a/V2 accounts for the internal pressure per mole resulting from the intermolecular forces of attraction between the molecules; b accounts for the excluded volume, which is about four times the molecular volume. 46  Real Gas Van der Waals Equation The influence of non-ideality is greater when the gas is compressed (At high pressure and low temperature). When the volume of a gas is large (At low pressure and high temperature), the molecules are well dispersed and far apart. Under these conditions, a/V2 and b become insignificant with respect to P and V, respectively, and the van der Waals equation for the real gas reduces to the ideal gas equation: PV = nRT At low pressures, real gases behave in an ideal manner. 47 THANKS