Chemistry 1210, Chapter 5 Notes
Chemistry 1210, Chapter 5 Notes CHEM 1210
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This 3 page Class Notes was uploaded by Grace Campbell on Friday September 30, 2016. The Class Notes belongs to CHEM 1210 at Ohio State University taught by in Fall 2016. Since its upload, it has received 4 views.
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Date Created: 09/30/16
Chapter 5 Notes: Thermochemistry 9/22/16-9/28/16 I. Thermochemistry- the study of energy changes that relate to chemical reactions a. Nature of energy i. Energy = capacity to do work ii. Work (J) = force x distance iii. Heat = energy causes the temperature of something to increase iv. Calorie (cal) – amount of energy needed to raise the temperature of 1g of water by 1°C 1. 1 cal = 4.184 J b. Kinetic and potential energies i. Kinetic Energy (J) = 1/2mv 2 1. Motion energy ii. Potential energy 1. Stored energy 2. Chemical energy 3. Electrostatic PE a. The interaction between charged particles c. System (what we single out to study) and surroundings (everything else) i. EX: contents of a flask(system) and the flask and everything outside of it (surroundings) II. The First Law of Thermodynamics a. The Law of Conservation of Energy i. Energy can neither be created nor destroyed, but can be converted into a different form ii. If energy is lost from a system, it must be gained by the surroundings b. Internal energy (E) i. Total energy of system ii. Actual value cannot be determined, but its change can c. State function- property that depends only on its initial and final value i. Does NOT depend on the path that it took during a reaction ii. Extensive d. Heat (q) and work (w) i. Are not state functions ii. q is positive = endothermic (heat gained to system) iii. q is negative = exothermic (heat lost to surroundings) iv. w is positive = work done ON the system v. w is negative = work done BY the system e. change in internal energy = q + w f. exothermic reaction releases heat g. endothermic reaction absorbs heat i. requires input of energy III. Enthalpy a. Is a state function b. At constant pressure i. w = -P∆V 1. negative because work is done BY the system ii. ∆E = q p P∆V iii. qp= ∆E + P∆V c. At constant volume i. ∆E = q v d. Enthalpy (H) = E + PV e. ∆H = ∆E + P∆V f. ∆H = q p IV. Enthalpies of reaction a. ∆H rxn H products reactants b. endothermic reactions are the reverse of the exothermic reactions c. ∆H is negative for an exothermic reaction and is positive for an endothermic reaction V. Calorimetry a. Experimental method utilizing the fact that any heat evolved or absorbed by the system will be reflected in the surroundings b. Heat capacity (C) i. C = q / ∆T ii. Molar heat capacity = heat capacity per mole iii. Specific heat = heat capacity per gram c. Calorimeter = ∆H at constant pressure d. Bomb calorimeter = ∆E at constant volume VI. Hess’s Law a. ∆H = ∑ ∆H rx steps b. add the chemical equations for all of the steps to get to the overall reaction c. add the ∆H of all the steps to get the ∆H of the overall reaction d. Tips when solving Hess’s Law problems: i. if you reverse an equation, the sign of that ∆H changes ii. if you multiply an equation by a factor, that ∆H gets multiplied by the same factor iii. any substances that don’t appear in the final equation must cancel out VII. Enthalpy of Formation - The enthalpy change for the formation of a compound from its elements (∆H ) f a. Standard enthalpy change- the enthalpy change when all reactants and products are in their standard states (∆H°) i. Standard state- the stable state of a substance in its pure form at standard pressure and temperature 1. Solid or liquid a. Pure substance at 1 atm 2. gas a. pressure at 1 atm 3. species in solution a. concentration of 1 M b. standard enthalpy of formation (∆H °)- f mole of a compound is formed from its elements with all of the substances in their standard states i. element in its standard state = 0
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