lec1 physical pharmacy.pdf

Ministry of Higher Education and Scientific Research Al- Amal College for Specialized Medical Sciences Pharmacy department, undergraduate, 2nd class Year 2024-2025 Physical pharmacy I Lecture 1...

Ministry of Higher Education and Scientific Research Al- Amal College for Specialized Medical Sciences Pharmacy department, undergraduate, 2nd class Year 2024-2025 Physical pharmacy I Lecture 1 Dr. Ahmed AL-mouswy Reference text: Physical Pharmacy by Alfred Martin et al. 1 Is the branch of pharmacy that concentrates on the application of physics and chemistry to the study of pharmacy. It forms the basis for design, manufacture, and distribution of drug products and serves as the foundation for the stable and proper use of medical drugs. It covers areas such as solubility, pharmacokinetics and drug delivery. 2  Contents OF PART 1 In this lecture, you'll learn: What is “Matter” ? what is “Matter made from” ? What are the “States of Matter”? What are the “Forces which attach components of matter to each other “? 3 1. What is “Matter” ? Matter can be defined as anything that has a mass and a volume. 4 5 6 7  States of Matter The states of matter: Gases, liquids and Solids. For molecules to exist as aggregates in gases, liquids and solids, intermolecular forces must exist. 8 W hat is “Matter” made from ? All matter is composed of small particles (atoms, molecules, or ions). 9 10 11 12 13  Intermolecular 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. 14 15 16 17  Intermolecular forces Types of intermolecular forces Ion-dipole (between ions and polar molecules) dipole-dipole (between 2 polar molecules) dipole-induced dipole (between a polar molecule and a non-polar molecules) induced dipole - induced dipole (between 2 nonpolar molecules). Water uses ion-dipole forces to dissolve salts 18 19 Intermolecular forces Van der W aals 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: A. Dipole–dipole interaction, orientation effect, or Keesom force B.Dipole-induced dipole interaction, induction effect, or Debye force C.Induced dipole–induced dipole interaction, dispersion effect, or London force 20 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), one another (dipole-dipole interactions) The Partial opposite charges (permanent dipoles) attract. 21 Van der Waals Forces: 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. 22 Intermolecular forces Van der Waals Forces: 2- 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 23  Van der W aals 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. 24 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. 25 Intermolecular forces Van der Waals Forces: 3- 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 26  Intermolecular forces 1-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. 27 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. 28 Intermolecular forces 2- 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 29 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 (they are given their own classification), but are still weaker than ionic and covalent bonds. 30 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. 31 States of matter 1. Gaseous state 2. Liquid state 3. Solid and crystalline state 4. Liquid crystalline state 32  CONTENT OF PART 2 The gaseous state What are Gas general properties? What is Ideal gas? What are Laws of gas? What is Real gas? What are differences of ideal gas and real gas ? 33 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). 34 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)). 35 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 GayLussac laws 36 Ideal gas 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). 37 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). 38 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). 39 Ideal gas Combined gas law Boyle, Gay-Lussac and Charles law can be combined to obtain the familiar relationship: 40 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? 41 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. 42 Ideal gas 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. R: the molar gas constant value for the PV/T ratio of an ideal gas. For n moles it becomes: 43 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: The molar gas constant can also be expressed by energy units: 44 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? 45 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: 46 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? 47 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. 3. The particles exhibit continuous random motion. 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. 48 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 49 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. 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. 50 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 these conditions, real gases behave in an ideal manner. 51 52 53