Physical Chemistry Lecture 2 PDF

Prof. Dr. Hossieny Ibrahim Badr University in Assiut School of Biotechnology [email protected] Office number: Bio-326 1 Unit 1: Thermodynamics and Thermochemistry  Thermodynamics:  The science of energy and its transformations.  Thermochemistry The...

Prof. Dr. Hossieny Ibrahim Badr University in Assiut School of Biotechnology [email protected] Office number: Bio-326 1 Unit 1: Thermodynamics and Thermochemistry  Thermodynamics:  The science of energy and its transformations.  Thermochemistry The branch of thermodynamics specifically focused on the changes in energy and transfer of heat related to chemical reactions.  Temperature and Heat Temperature: is a measure of the average kinetic energy of the molecules in a substance.  Heat is a measure of the internal energy of the molecules that make up a substance.  Heat may be transferred from one substance to another.  We say that substances can absorb or radiate heat. 2  The difference between Heat and Temperature  Heat is a form of energy that is transferred between two bodies because of a temperature difference existing between them.  Processes of Heat Transfer are: ❶ Conduction, ❷ Convection ❸ Radiation Convection: The transfer of heat by the actual motion of a fluid (liquid or gas) in the form of currents. Conduction: The transfer of heat by direct contact of particles of matter. Radiation: Heat transfer by electromagnetic waves. 3  Internal energy  Internal energy (E): the total energy of a thermodynamic system i.e., the sum of all the different forms of energy contained by all components of the system.  Internal energy includes TWO broad categories: ❶ Kinetic Energy: Translational, rotational, and vibrational motions of the atoms and molecules making up the system. ❷ Potential Energy: Binding energy of the atoms and molecules that make up the system.  primarily chemical bonds (covalent, ionic)  also intermolecular forces, binding energy of nuclei, etc. 4 Unit 1: Calorimetry: The Measurement of Heat Flow  Calorimetry: The device used experimentally to determine the heat associated with a chemical reaction is called a calorimeter.  Calorimetry, the science of measuring heat, is based on observing the temperature change when a body absorbs or discharges energy as heat.  The change in temperature, ΔT, of the calorimeter is proportional to the heat that the reaction releases or absorbs. A bomb calorimeter 5 Unit 1: Units of Energy Changes  Units of Energy:  The SI unit of energy is the joule (J). 1 kJ = 1000 J  An older, non-SI unit is still in widespread use:  The calorie (cal), which is the heat required to raise the temperature of one gram of water 1 ºC.  1 cal = 4.184 J 1 kcal = 4.18 kJ Problem:  Make the following conversions. (a) Convert 725 cal to kJ (b) 444 calories to Joules (c) 850 Joules to calories 6 Unit 1: Specific Heat Capacity  Heat Capacity : the amount of heat needed to increase the temperature of an object exactly one degree (K or ºC).  Specific Heat Capacity: The amount of heat required to raise the temperature of 1 g of a substance by one degree (K or ºC).  Units are: J/ (g.ºC) or J/ (g.K), cal/ (g.ºC) or cal/ (g.K),  Molar Heat Capacity: The amount of heat required to raise the temperature of one mole of substance by one degree (K or ºC).  Units are: J/ (mol.ºC) or J/ (mol.K), cal/ (mol.ºC) or cal/ (mol.K) 7 Unit 1: Specific Heat Capacity (J/g.K) Note the large difference in Specific Heat. Water’s The Higher the specific value is heat the More energy VERY HIGH. needed to change the temperature 8 Unit 1: Solved Problem If 25.0 g of Al cool from 310 oC to 37 oC, how many joules of heat energy are lost by the Al? (sp.ht of Al = 0.897 J/g.K) where ∆T = Tfinal - Tinitial q = (0.897 J/g.K)(25.0 g)(37 - 310)K q = - 6120 J 9 Unit 1: Solved Problem Solved Problem: The temperature of a piece of copper with a mass of 95.4 g increases from 25 degrees Celsius to 48.0 degrees Celsius when the metal absorbs 849 J of heat. What is the specific heat capacity of copper? q = m. C. ΔT , where ( q ) the amount of heat absorbed, ( m ) the mass of the sample (C ) the specific heat of the substance, (ΔT ) the change in temperature, Quiz: A piece of iron with a mass of 75.0 g at 125.0 °C is allowed to cool to room temperature of 25.0 °C. Determine the magnitude and sign of q. The specific heat capacity of iron is 0.449 J/g.°C. 10 Unit 1: System and Surroundings  System = the molecules we want to study (here, the hydrogen and oxygen molecules).  Surroundings = everything else (here, the cylinder and piston). Piston H2 & O2 Cylinder 11 Unit 1: Work: Pressure-Volume Work  We know that energy transfer can occur via Heat or Work. We have seen how to determine the amount of Heat transfer by measuring the change in temperature.  Now we will look at how to determine the amount of Work by measuring the change in volume.  Chemical reactions can do several different types of work.  We are only going to consider pressure-volume (or PV) Work.  Work is a force (F) acting through a distance (D).  So we define work as Force x Distance or F x D. 12 Unit 1: Work: Pressure-Volume Work  PV work occurs when the force is caused by a Volume change against an external Pressure.  Like in the engine of a car – when the pistons are pushed outward against the external atmospheric pressure.  Pressure define as: P = F/A which we can rearrange to F = P x A.  Substitute (P x A) for F into the definition for work and we get: w=PxAxD  w = P x A x ∆h  w = P∆V (where A x ∆h = ∆V) Negative because an increase in Volume means that  w = – P∆V the system is doing work on the surroundings. 13 Unit 1: Work: Pressure-Volume Work Solved Problem 14 Unit 1: Work: Pressure-Volume Work Solved Problem 15 Unit 1: Internal Energy (E)  The internal energy (E) of a system is the sum of all kinetic and potential energies of all components of the system.  The internal energy of a system is a state function:  A state function is a mathematical function whose result depends only on the initial and final conditions, not on the process used.  By definition, the change in internal energy, E, is the final energy of the system minus the initial energy of the system. E = Efinal − Einitial 16 Unit 1: Measuring ∆E  So we know that systems exchange energy with their surroundings via Heat and Work.  Change in internal energy during a reaction is a sum of both the heat and work. ∆E = q + w  and w = – P∆V ∆E = q – P∆V  So if we carry out the reaction at constant volume then ∆V = 0 and w = 0 so  ∆E = q  This is written as ∆Erxn = qv 17 Unit 1: Sign for Heat Thermodynamics Symbol of heat is (q) Surrounding Heat flows from Surrounding to System System Energy of the System is Raised So Heat is consider as Positive (+q) Surrounding Heat flows from System to Surrounding Energy of the System is lowered System So Heat is consider as Negative (-q) 18 Unit 1: Sign for Work in Thermodynamics Symbol of work is (w) Surrounding Work is done on the System System Energy of the System is Raised So Work is consider as Positive (+w) Surrounding Work is done by the System Energy of the System is lowered System So Work is consider as Negative (-w) 19 Unit 1: Types of Thermodynamic Systems  Open System: can exchange both Matter and Heat with the surroundings.  Closed System: can exchange Heat with the surroundings but not matter  Isolated System: No transfer of mass or Heat. 20 Unit 1: Types of Thermodynamic Systems 21 Unit 1: Intensive and Extensive Properties  Intensive Properties: A property which does not depend on the quantity of matter present in the system.  Extensive Properties: A property that does depend on the quantity of matter present in the system. 22 Unit 1: Thermodynamic Processes  When a thermodynamic system changes from one state to another, the operation is called a Process.  These processes involve the change of conditions (temperature, pressure and volume).  The various types of thermodynamic processes are : (1) Isothermal Processes:  Processes in which the temperature remains Constant, are termed isothermal processes.  For an isothermal process dT = 0 or (T2-T1) = 0  Isothermal expansion  Isothermal compression 23 Unit 1: Thermodynamic Processes (2) Adiabatic Processes: Those processes in which no heat can flow into or out of the system.  Adiabatic conditions can be approached by carrying the process in an insulated container such as ‘thermos’ bottle.  For an adiabatic process dq = 0 Adiabatic Expansion Adiabatic Compression (P and T decreases) (P and T increases) Because work is done but no heat enters the system, the internal energy decreases, and therefore the temperature of the working gas also decreases. 24