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by: Caitrín Hall

Thermochemistry CHEM 1127Q 001

Caitrín Hall
GPA 3.9

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About this Document

These notes cover chapter 5 of the textbook.
General Chemistry
Fatma Selampinar (TC), Joseph Depasquale (PI)
Class Notes
Chemistry, thermochemistry, outline
25 ?




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This 4 page Class Notes was uploaded by Caitrín Hall on Wednesday March 30, 2016. The Class Notes belongs to CHEM 1127Q 001 at University of Connecticut taught by Fatma Selampinar (TC), Joseph Depasquale (PI) in Spring 2016. Since its upload, it has received 25 views. For similar materials see General Chemistry in Chemistry at University of Connecticut.


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Date Created: 03/30/16
Chapter 5: Thermochemistry 5.1 Energy Basics  Thermochemistry – the study of measuring the amount of heat absorbed or released during a  chemical reaction or a physical change Energy – the capacity to supply heat or do work  Work (w) – the process of causing matter to move against an opposing force  Potential energy – the energy an object has because of its relative position, composition,  or condition  Kinetic energy – the energy an object possesses because of its motion  When one substance is converted to another, there is a conversion of energy from one  form to another (heat, light, electrical energy, chemical energy, mechanical energy)  Law of conservation of energy: energy can be converted from one form to another during a chemical or physical change but not created nor destroyed  Law of conservation of matter: there is no detectable change in the total amount of matter during a chemical change  At the molecular level, these laws are combined: the total quantity of matter and energy  in the universe is fixed Thermal Energy, Temperature, and Heat  Thermal energy – kinetic energy associated with the random motion of atoms and  molecules  Temperature – a quantitative measure of “hot” or “cold” o Rapid movement  high KE  “hot” o Slow movement  low KE  “cold” o Increasing energy will increase temperature and the substance expands o Decreasing energy will decrease temperature and the substance contracts  Heat – the transfer of thermal energy between two bodies at different temperatures o Increases thermal energy of one body and decreases thermal energy of the other o Flows spontaneously from high temperature substance to low temperature  substance; heat flow continues until both substances are the same temperature o Exothermic process – a reaction or change that releases heat (ex: combustion); q  is negative o Endothermic process – a reaction or change that absorbs heat (ex: when the  substances in a cold pack (water & salt) come together, heat is absorbed leading  to the sensation of cold; q is positive  A calorie is the amount of energy required to raise one gram of water by 1 kelvin o Depends on the atmospheric pressure and starting temperature of the water  A joule is the amount of energy used when a force of 1 newton moves an object 1 meter o SI unit of heat, work, and energy  Heat capacity (C) – the quantity of heat (q) a body of matter absorbs or releases when it  experiences a temperature change (T) of 1 kelvin C = q/T o Determined by both the type and amount of substance that absorbs/releases heat o Ex: a small frying pan has a lower C than a large frying pan does because it  requires less energy to raise the temperature  Specific heat capacity (c) – the quantity of heat required to raise the temperature of 1  gram of a substance by 1 kelvin c = q/(mT) o Depends only on type of substance absorbing/releasing heat (intensive property) o Ex: a small frying pan requires less energy to heat but also weighs less so has the  same specific heat as a large frying pan o Derived from a ratio of 2 extensive properties (heat and mass) o Molar heat capacity (another intensive property) – heat capacity per mole of a  particular substance and has units of J/molC o Specific heat of liquid water = 4.184 J/molC 5.2 Calorimetry   Calorimetry – process of measuring the amount of heat involved in a chemical or  physical process o Heat is exchanged with calibrated object (calorimeter)  System – the substance or substances undergoing the chemical or physical change  Surroundings – the other components of the measurement apparatus that serve to either  provide heat to the system or absorb heat from the system  Calorimeter – a device used to measure the amount of heat involved in a chemical or  physical process o When an exothermic reaction occurs in solution in a calorimeter, heat produced  by the reaction is absorbed by the solution (solution temperature increases) o When an endothermic reaction occurs, heat required is absorbed from the thermal  energy of the solution (solution temperature decreases)  Heat produced or consumed in the system plus heat absorbed or lost by the surroundings  must add up to zero q + q = 0 reaction solution q reaction = -(q solution)  Bomb calorimeter – type of calorimeter that operates at constant volume to measure  energy produced by reactions that yield large amounts of heat and gaseous products (ex:  combustion reactions) o Require calibration to the heat capacity of the calorimeter and to ensure accuracy o Accomplished using a reaction with known q and m o Temperature change produced by the known reaction is used to determine the heat capacity of the calorimeter 5.3 Enthalpy   Chemical thermodynamics – the science that deals with the relationships between heat,  work, and other forms of energy in the context of chemical and physical processes  Energy is stored in a substance when the kinetic energy of its atoms/molecules is raised  Internal energy (U) – the total of all possible kinds of energy present in a substance  (sometimes symbolized as E)  Energy is transferred into a system when it absorbs heat (q) from the surroundings or  when the surroundings do work (w) on the system  First law of thermodynamics – internal energy of a system changes through heat flow  into (positive q) of out of (negative q) the system or work done on (positive w) or by the  (negative w) system ΔU = q + w   State function – state variable; depends only on the state that a system is in, and not on  how that state is reached; (ex: internal energy, pressure, volume, enthalpy)  Enthalpy (H) – the sum of a system’s internal energy (U) and the mathematical product  of its pressure (P) and volume (V) H = U + PV  Enthalpy change (ΔH) – enthalpy changes for chemical or physical processes can be  determined, but enthalpy values for specific substances cannot be measured directly ΔH = H products Hreactants  The following conventions apply: 1. The ΔH value indicates the amount of heat associated with the reaction involving  the number of moles of reactants and products as shown in the chemical equation 2. The enthalpy change of a reaction depends on the physical state of the reactants  and products of the reaction 3. ΔH > 0 for endothermic reactions ΔH < 0 for exothermic reactions 4. Products and reactants are at the same temperature (standard = 298 kelvins)  Standard state – set of conditions used as a reference point for the determination of  properties under other different conditions  Standard enthalpy of combustion (ΔHc) – the enthalpy change when 1 mole of a  substance burns under standard state conditions (25C and 1 atm)  Standard enthalpy of formation (ΔHf – the enthalpy change when 1 mole of a pure  substance is formed from free elements in their most stable states under standard state  conditions Hess’s Law: If a process can be written as the sum of several stepwise processes, the enthalpy  change of the total process equals the sum of the enthalpy changes of the various steps 1. ΔH is directly proportional to the quantities of reactants or products 2. ΔH for a reaction in one direction is equal in magnitude and opposite in  sign to ΔH for the reaction in the reverse direction 3. Elements in their standard states have enthalpies of formation of zero ΔH° reaction = ∑ n × ΔH° (pfoducts) − ∑ n × ΔH° (feactants)


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