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## CHEM 1030 Week 14 Notes

by: Alyssa Anderson

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# CHEM 1030 Week 14 Notes CHEM 1030

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Alyssa Anderson
AU

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These notes cover we went over in class on 4/19/16. It includes several diagrams and tables from the chapter.
COURSE
Fundamentals Chemistry I
PROF.
Dr. Streit
TYPE
Class Notes
PAGES
6
WORDS
KARMA
25 ?

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This 6 page Class Notes was uploaded by Alyssa Anderson on Wednesday April 20, 2016. The Class Notes belongs to CHEM 1030 at a university taught by Dr. Streit in Spring 2016. Since its upload, it has received 12 views.

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Date Created: 04/20/16
CHEM 1030 Week 14 Thermochemical Equations A. The following guidelines are useful when considering thermochemical equations 1. Always speciﬁc the physical states of reactants and products because they help determine the actual enthalpy changes (different states have different enthalpies) 2. When multiplying an equation by a factor (n), multiply the delta H blue by the same factor 3. Reversing an equation changes the sign but not the magnitude of delta H (multiply by -1) B. Worked example 10.3 1. Calculate the solar energy required to produce 75.0 g of C6H12O6. 2. Calculate the number of moles needed, then multiply it by the entire equation Calorimetry A. The measurement of heat changes B. Heat changes are measured i a device called a calorimeter C. The speciﬁc heat (s) of a substance is the amount of heat required to raise the temperature of 1 g of the substance by 1*C. D. The heat capacity (C) is the amount of heat required to raise the temperature of an object by 1*C E. The “object” may be a given quantity of a particular substance F. Heat capacity of water = 4.184 J / (1g x *C) x 1000g = 4184 J/*C G. Speciﬁc heat capacity has units of J/(g • °C) H. Heat capacity has units of J/°C Speciﬁc Heat and Heat Capacity A. The heat associated with a temperature hcnage ay be calculated B. q = smΔT C. q = CΔT D. Calculate the amount of heat required o heat 1.01 kg of water from 0.05*C to 35.81*C. Constant-Pressure Calorimetry A. Concepts to consider for coffee-cup calorimetry: p = delta H B. In an exothermic reaction, the system loses heat C. Worked example 10.5 D. Constant volume calorimetry is carried out in a device known as a constant-volume bomb E. A constant-volume calorimeter is an isolated system. F. Bomb calorimeters are typically used to determine heats of combustion. G. q = −q cal rxn Hess’s Law A. Hess’s law states that the change in enthalpy for a stepwise process is the sum of the enthalpy changes for each of the steps B. When applying Hess’s Law: 1. Manipulate thermochemical equations in a manner that gives the overall desired equation 2. Remember the rules for manipulating thermochemical equations: a. Always specify the physical states of reactants and products because they help determine the actual enthalpy changes. b. When multiplying an equation by a factor (n), multiply the ΔH value by same factor. c. Reversing an equation changes the sign but not the magnitude of ΔH. 3. Add the ΔH for each step after proper manipulation. 4. Process is useful for calculating enthalpies that cannot be found directly. Standard Enthalpies of Formation A. The standard enthalpy of formation (ΔH° ) isfdeﬁned as the heat change that results when 1 mole of a compound is formed from its constituent elements in their standard states B. The superscripted degree sign denotes standard conditions 1. 1 atm pressure for gases 2. 1 M concentration for solutions D. “f” stands for formation E. ΔH f° for an element in its most stable form is zero. F. ΔH f° for many substances are tabulated in Appendix 2 of the textbook G. The standard enthalpy of reaction (ΔH °rxn ) is deﬁned as the enthalpy of a reaction carried out under standard conditions. H. ΔH °rxn = [cΔH f°(C) + dΔH f°(D) ] – [aΔH f°(A) + bΔH f°(B)] I. ΔH °rxn = ΣnΔH f°(products) – ΣmΔH f°(reactants) J. n and m are the stoichiometric coefﬁcients for the reactants and products. Bond Enthalpy and the Stability of Covalent Molecules A. The bond enthalpy is the enthalpy change associated with breaking a bond in 1 mole of gaseous molecule B. H (g) → H(g) + H(g) // ΔH° = 436.4 kJ/mol 2 C. The enthalpy for a gas phase reaction is given by: ΔH° = ΣBE(reactants) – ΣBE(products) D. ΔH° = total energy input – total energy released E. Bond enthalpy change in an exothermic reaction F. Bond enthalpy change in an endothermic reaction Lattice Energy and the Stability of Ionic Compounds A. A Born-Haber cycle is a cycle that relates the lattice energy of an ionic compound to quantities that can be measured B. Na(s) + 1/2 Cl (g) → Na (g) + Cl (g) - 2 Comparison of Ionic and Covalent Compounds A. Ionic and covalent compounds differ in their general physical properties because of the differences in the nature of their bonds. B. Check out table 10.6

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