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CHM 113- Week 7 Notes

by: Andrew Notetaker

CHM 113- Week 7 Notes CHM 113

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These cover the lecture material from week 7, Chapter 5.
General chemistry 1
Class Notes
hess's law heats of reactions general chemistry janet degrazzia engineering engineers thermochemistry, thermochemistry, calorimetry, Enthalpy, thermal physics entropy enthalpy heat engine refrigerator
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This 4 page Class Notes was uploaded by Andrew Notetaker on Friday September 30, 2016. The Class Notes belongs to CHM 113 at Arizona State University taught by Cabirac in Fall 2016. Since its upload, it has received 8 views. For similar materials see General chemistry 1 in Science at Arizona State University.


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Date Created: 09/30/16
Chapter 5: Thermochemistry Monday, September 26, 2016 2:56 PM 5.1 The Nature of Energy Most chemical reactions give off or absorb energy, generally as heat. Thermochemistry-Study of heat changes in chemical reactions Energy- The capacity to do work or transfer heat. Work (w): the transfer of energy when a force movesan object - w = F x d (force x distance) Heat(q): the transfer of energy from a hotter to a coolerobject Forms of energy: Potential & Kinetic Kinetic: Energy of motion Ek= 1/2 mv ( T) Potential: Stored energy from attractions/repulsions Two important concepts: • Breaking chemical bonds requires input of energy • Forming chemical bonds releases energy The Joule (J) is the SI unit for energy 1 calorie is the amount of energy required to increase the temperature of 1.0g of water by 1°C 1 cal = 4.184J 1Cal = 1kcal = 1000 cal System: The part of the universe we are interested in studying Surroundings: the rest of the universe When the system loses energy it has a - sign When it gains energy it has a + sign System: chemical reaction Surroundings: all else Energy can energy/leavesystem as heat (q) or work (w) done by piston. Heat and work are not forms of energy but ways energy is transferred A closed system is one in which energy can be exchanged (through q or w) between system and surroundings but not matter Energy can be transferred: • Through work: force x distance -(F x d) • Through heat: transfer of thermal (kinetic) energy between two systemsat different temperatures • Through heat: transfer of thermal (kinetic) energy between two systemsat different temperatures 5.2 First Law of Thermodynamics Energy is neither created or destroyed. If a system loses energy, the surroundings gain energy. Internal energy ( E ) = Sum of E k sum of E p The change in internal energy is written as: ΔE = E finalEinitial If ΔE is negative (<0) Energy is lost to surroundings and final state has less energy that initial. If ΔE is positive(>0) Energy is gained from surroundings and final state has more energy than initial. ΔE = q + w to determine internal energy For q + means gains heat -Means loses heat For w + means work done on system -Work is done by system For ΔE + means net gain of energy by system Means net loss of energy by system Exothermic process indicates heat is lost by the system. Endothermic process indicated heat is gained by system. 5.3 Enthalpy ΔE= q + w = q - PΔV For most chemical reactions:ΔE=q p Which means the change in energy during a chemical reaction under constant pressure is approximately equal to the heat involvedin the reaction. Q ps called Enthalpy (H) When does ΔH =ΔE? The enthalpy change for a reaction equals the change in internal energy when work is negligible. Therefore ΔE approximatelyequals ΔH ΔH rxn products (finalreactants (initial) The enthalpy change is a measure of heat gained or lost ΔH is also a good estimate of internal energy ΔH <0 (negative) for exothermic reactions ΔH >0 (positive) for endothermic reactions 5.4 Enthalpies of Reaction General considerationsfor enthalpies: 1. ΔH for a reaction in the forward direction is equal in size, but opposite in sign, to ΔH for the reverse reaction ΔH for a forward reaction = -890 kJ ΔH for the reversed reaction = 890kJ 2. Enthalpy is an extensive property Na + Cl -> NaCl ΔH =30 kJ 2Na + 2Cl -> 2NaCl ΔH = 60 kJ 3. ΔH for a reaction depends on the state of the products and reactants Enthalpies of phase change Solid phase to liquid: ΔH = + Liquid phase to a gas: ΔH=+ Solid phase to a gas ΔH = + Amount of heat required to melt: Heat of fusion, ΔH fus Amount of heat required to evaporate: Heat of vaporization, Δh vap Amount of heat required to sublimate: Heat of sublimation, Δh sub 5.5 Calorimetry The value of ΔH can be determined by measuring the heat flow accompanying a reaction at constant pressure. Calorimetersare used to determine this Calorimetry is an experimental method for measuring heat flow (q or ΔH) by measuring temperature change (ΔT) of the system or the surroundings. The greater the amount of heat, q, the greater the change in temperature. The heat needed to raise the temperaturevaries. Heat capacity ( C ): amount of heat required to raise the temperature of a given quantity of a substance by 1°C; C=q/ΔT Units of J/°C Specific heat (Csor s): amount of heat required to raise 1 gram of a substance by 1°C. Molar heat capacity (C m: amount of heat that can be absorbed by 1 mole of substance when Molar heat capacity (C ):mamount of heat that can be absorbed by 1 mole of substance when temperatureincreases 1°C. q= (molesof substance) x (C ) x(ΔT) m Q = m x s x ΔT q= heat (J) M = mass (g) S = specific heat (J/g*°C) ΔT=(T final Tinitial For chemical reactions performed in a calorimeter, Heat released/consumedby a reaction = heat gained/ lost by a calorimetersolution ΔH rxn= qp= q rxn= -qsoln


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