CHM1102 Lecture 25: Introduction to Thermodynamics PDF

CHM1102 Lecture 25 Introduction to Thermodynamics  Introduction to Thermodynamics The study of the change of energy accompanying a “process” within a system a change of state (expansion), a change of physical state or phase (melting) or a chemical reaction Systems t...

CHM1102 Lecture 25 Introduction to Thermodynamics  Introduction to Thermodynamics The study of the change of energy accompanying a “process” within a system a change of state (expansion), a change of physical state or phase (melting) or a chemical reaction Systems tend to prefer to be in a lower energy level Spontaneous chemical reactions tend to move from a high energy level to a lower energy level The particular part of the universe that is being studied is referred to as the “system”, which has distinct physical boundaries – e.g. the reactants and the products The “surroundings” is every other part of the universe outside the system Universe = System + Surroundings  Types of Systems Three types of systems 1. Open system: Allows exchange of matter and energy between the system and its surroundings – e.g. an open bottle of soft drink 2. Closed system: Allows only the exchange of energy between the system and its surroundings, no matter is exchanged – e.g. A corked bottle of soft drink 3. An isolated system: Allows neither matter nor energy to be exchanged with its surroundings – e.g. A thermos flask used to store cold or hot liquids/foods  State Functions & Path Functions A state function is used to describe the condition (state) of a system property of a system that is dependent on the value of the property and not the means by which the system attained this state, e.g. Altitude The change in state functions is only dependent on the initial and final states not the intermediate values of getting to this state A path function depends on the process or path of getting from one value to the next value – e.g. Work State functions are given symbols using upper case letters State and Path Functions  Work & Heat Energy is required to do work – “Energy is the capacity to do work” Work is the transfer of energy to change motion of the atoms in a uniform manner Work is done when a force is moved over a distance w = F x d (Joules) When a system expands against a pressure it does work w = d x PexA (Pex – external pressure; A – Area) d x A = ∆V (change of volume) w = - Pex ∆V (-ve sign indicates work is being done by the system, using energy) Heat (q) is the transfer of energy to cause a chaotic motion of the atoms  Work ΔV Work Gas expands + - Gas compresses - +  Heat (q) The transfer of thermal energy between two bodies at different temperatures. Heat flow (a redundant term) increases the thermal energy of one body and decreases the thermal energy of the other.  First Law of Thermodynamics Energy cannot be created nor destroyed but only changed from one form to another (Conservation of Energy Law) The internal energy of an isolated system is constant (First Law of Thermodynamics) ∆U = w + q U = internal energy; w = work; q = heat The sign indicates if internal energy is lost or gained by the system Internal energy is an example of a state function (or state variable), whereas heat and work are not state functions.  Enthalpy(H) Now if the system does some PV work w = -P∆V ∆U = q -P∆V q = ∆U + P∆V If the process is done at constant pressure qp = ∆U + P∆V qp is given a special name Enthalpy (H)(a state function) ∆H = ∆U + P∆V (at constant pressure) H = U + PV Enthalpy is heat change at constant pressure ΔH = Hproducts– Hreactants  Calorimetry When a reaction occurs it may give off heat to the surroundings or absorb heat from the surroundings. The process of giving heat to the surroundings is called an exothermic process (-∆H, or ∆H < 0) The process of taking heat from the surroundings is called an endothermic process (+∆H or ∆H > 0) The method used to measure this heat change is called Calorimetry done by carefully tracking the temperature change during the reaction process Calorimetry  Calorimetry Regular calorimetry is done at constant pressure so enthalpy is determined Each material will increase in temperature when a certain amount of heat is applied or withdrawn This is the heat capacity of the material The amount of heat required to cause a change of one degree per unit mass is called the specific heat capacity (Cs) ‘coffee cup’ calorimeter e.g. 10 kJK-1 mol-1 or 12 kJK-1 g-1 – used to experimentally determine heat changes Cwater = 4. 184 J/g.K  Calorimetry If the temperature change is ∆T ( T2-T1) for a process The amount of heat q evolved is q = -m. C. ∆T m = mass of solution C = specific heat capacity of solution ∆T = Tfinal - Tinitial  Enthalpy of Phase Changes When a substance changes state the energy used for that phase change is called the enthalpy of the phase change Enthalpy of fusion - ∆Hf (kJ mol-1) Enthalpy of vaporisation - ∆Hvap (kJ mol-1)  Enthalpy of Formation Standard Enthalpy of formation (∆Hf°) of a substance is the standard reaction enthalpy for the formation of the substance from its elements in their most stable form under standard conditions (1atm, 298K, 1M) The standard state of a substance is when the substance is in its pure form at a pressure of 1 atm 2C(s) + 3H2(g) + ½ O2(g) → C2H5OH(l) ∆Hf°= -277.69 kJmol-1 The above reaction is a thermochemical equation The standard enthalpy of formation of an element in its most stable form is 0 kJmol-1.e.g. ∆Hf° {O2} = O kJmol-1