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Thermochemistry Lecture Notes Thermodynamics It is the science that studies the relationship between heat and other forms of energy. First Law of thermodynamics It is also known as the law of conservation of energy: “the total energy of the Universe is constant” or sometimes expressed as "Energy can neither be created nor destroyed". Another approach is to say that the total energy of an isolated system remains constant. Energy • Energy (E or U) - is the capacity to do work (W) or to produce heat (Q). Energy is a state function, does not depend on a pathway, it depends only on the present state. ΔE = Q + W Internal Energy = Heat + work Work & Heat • Work is defined as force acting over a distance. • Heat is the energy that flows between objects that become in contact if they are at different temperatures. It is an extensive property of matter, depends on the amount of matter. Energy Units Name Symbol Unit conversion 1 Joule J 1 kg.m2∕ s2 1 kilojoule kJ 1000 J 1 calorie cal 4.184 J 1 Calorie Cal 1000 cal Other forms of Energy Kinetic Potential Thermal energy; all matter has thermal energy at the submicroscopic level due to motion of submicroscopic particles. Chemical potential energy; contained in chemical bonds, observed in chemical reactions. Mechanical energy; all moving objects have it. Gravitational energy; caused by the effect of gravity over matter. Electric energy; generated by moving Electrostatic energy; attraction electrons. between particles of opposite charges placed at a certain distance. Sound; due to compression & expansion of spaces between molecules. Heat and Temperature • Temperature – T It is the property of matter that determines whether heat energy can be transferred from one body to another and in what direction. It is an “intensive” property, it does not depend on the amount of matter. • Heat – q It is thermal energy. It is directly proportional to the motion of molecules in matter. It is an extensive property. It depends on the amount of matter. Second Law of Thermodynamics • The second law states that heat energy cannot be transferred from a body at a lower temperature to a body with a higher one without the addition of energy. Heat flows from matter at a higher temperature to matter at a lower temperature. Third law of Thermodynamics The Third Law of Thermodynamics states that the entropy of any pure substance in thermodynamic equilibrium approaches zero as the temperature approaches zero (Kelvin), or conversely the temperature (Kelvin) of any pure substance in thermodynamic equilibrium approaches zero when the entropy approaches zero. Entropy is the measure of disorder in a system. System & Surroundings “Thermo” Vocabulary • System – designated part of the universe under study. • Surroundings - Everything outside the system, the rest of the universe. In practice one considers an “immediate surroundings” and then surroundings. • Heat flow - the direction at which thermal energy flows. It always goes from the matter at higher temperature to the one at a lower temperature. • Thermal equilibrium – It is the point at which all heat transfer stops. • Law of conservation of energy – The quantity of heat lost is equal to the quantity of heat gained. Heat Transfer ENDOTHERMIC EXOTHERMIC SURROUNDINGS ↓ SYSTEM qf > q i SURROUNDINGS ↑ SYSTEM qf<qi The system absorbs heat from the surroundings The system releases heat to the surroundings Specific Heat Capacity or Specific Heat • It is the quantity of heat required to raise the temperature of 1 gram of a substance by one kelvin degree. Its units are joules per gram per kelvin. • Heat = joules (energy unit) • Mass = grams • Temperature = kelvin scale Specific Heat Capacity, C (J/g.K) C = ___q____ m x (Tf-Ti) q = total heat transferred in Joules m = amount of matter in grams Tf = Final temperature in Kelvin degrees Ti = initial temperature in Kelvin degrees Sign Conventions (page 209) ΔTsys = Tf –Ti Δqsys= qf - qi Direction of heat transfer + Tf > Ti + qf>qi Endothermic process Tf<Ti qf<qi Exothermic process Example for calculating C • If 1160 Joules of heat were used to raise the temperature from 25 0C to 325 0C of 10.0 g of a metal sample, what is its specific heat capacity? Identify this metal using Table 6.1* page 210. DATA q = 1160 J Tf = 325 0C + 273 = 598 K Ti = 25 0C + 273 = 298 K m = 10.0g C=? SOLUTION C = ___ q____ = _____1160 J_____ = 0.385_ J_ (*copper) m . (Tf-Ti) 10.0 g . (598-298)K g.K Heat Transfer and specific heat capacity q sys + q surr = 0 This assumption is used to calculate experimentally the specific heat capacities of unknown metals. System = piece of metal , with known mass and initial high temperature. Immediate surrounding = water, with known specific heat, known mass and initial cool temperature. For the metal q sys = q metal= m metal. Cmetal. (Tf –Ti metal ) For the water q surr = q water = m water . (4.184 J/g.K) . (Tf-Ti water) q water + q metal = 0 Note that Tf is the same for both. Tf is the temperature at which thermal equilibrium has been reached and no more heat transfer occurs. Calorimetry • Heat transferred at constant pressure is called enthalpy, H. q rxn + q sol = 0 q rxn = - q sol • Heat transferred at constant volume is called internal energy , ΔE q rxn + q bomb + q sol = 0 q rxn = -(q bomb + q sol) Enthalpy Changes in Chemical Reactions • Enthalpies are measured taking the system (chemical reaction) as its reference point. • ΔH > 0 means the system has absorbed heat, it is an endothermic reaction. • ΔH < 0 means the system has released heat, it is an exothermic reaction. • ΔH of the reverse reaction is equal to - ΔH of the forward reaction. • ΔH is molar dependent. Hess’s Law “ If a reaction is the sum of two or more other reactions, ΔH for the overall process is the sum of the ΔH values of those reactions” ΔH3 = ΔH1 + ΔH2 + … Enthalpy Changes in Chemical Reactions • Enthalpies are measured taking the system (chemical reaction) as its reference point. • ΔH > 0 means the system has absorbed heat, it is an endothermic reaction. • ΔH < 0 means the system has released heat, it is an exothermic reaction. • ΔH of the reverse reaction is equal to - ΔH of the forward reaction. • ΔH is molar dependent. Standard Molar Enthalpy of Formation • Standard Molar Enthalpy of Formation, ΔHf, is the enthalpy change for the formation of 1 mol of a compound directly from its component elements in their standard states, usually measured at 298 K and 1 at. • Example ΔHf (CO2, 1at, 298 K) is • C(s) + O2(g) CO2(g) ΔH f = -393.5 kJ Enthalpy Change for a Chemical Reaction • Reactants Products • ΔH rxn = ΣΔH (products) – ΣΔH (reactants) • Exothermic Reactions are product favored, because the final enthalpy is lower, more stable. • Endothermic reactions are reactant favored.