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Chem 161-2011 Lecture 18, Chapters 10,11

Chem 161-2011 Lecture 18, Chapters 10,11 - CHEMISTRY...

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Chem 161-2011 Lecture 18 1 CHEMISTRY 161-2011 LECTURE 18 CHAPTERS 10,11 ENERGY CHANGES IN CHEMICAL REACTIONS GASES ANNOUNCEMENTS E-MAIL EXAMS Exam III, Wed, 11/30, Chapters 8.3 (8.4?) – 10.8 MISCELLANEOUS Tue, 11/22, no class (replaced by Thu class) Thanksgiving recess: Wed 11/23 – Sun 11/27
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Chem 161-2011 Lecture 18 2 Today’s topics Finish chapter 10 Chapter 11 Properties of Gases The Kinetic Molecular Theory of Gases Molecular Speed Diffusion and Effusion Gas Pressure Definition and Units of Pressure Calculation of Pressure Measurement of Pressure The Gas Laws Boyle’s Law: The Pressure-Volume Relationship Charle’s and Gay-Lussac’s Law: The Temperature-Volume Relationship Avogadro’s Law: The Amount-Volume Relationship The Gas Laws and Kinetic Molecular Theory The Combined Gas Law: The Pressure-Temperature-Amount-Volume Relationship The Ideal Gas Equation Applications of the Ideal Gas Equation Real Gases Factors that Cause Deviation from Ideal Behavior The van der Waals Equation van der Waals Constants Gas Mixtures Dalton’s Law of Partial Pressures Reactions with Gaseous Reactants and Products Calculating the Required Volume of a Gaseous Reactant Determining the Amount of Reactant Consumed Using Change in Pressure Using Partial Pressures to Solve Problems
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Chem 161-2011 Lecture 18 3 CHAPTER 10 RELEVANT EQUATIONS E = Δ U = q + w 1 Latm = 101.3 J Pressure-volume work: w = -P Δ V; therefore, E = q –P V Non-pressure-volume work: w = F x d (note, using a piston and cylinder, the piston has to move); F = m x a; w = m x a x d = m x g x d Units of non-pressure-volume work: m in kg, g = 9.81 m x s -2 , d in m; Units: kg x ms -2 x m = kgm 2 s -2 = J Heat Capacity = C = q/ Δ T molar heat capacity = C/mol = q/( Δ T x mol) Specific Heat Capacity = SHC = SH = C/g = (q/ Δ T)/g = q/ Δ Tg q = SHC x Δ T x g SHC H2O = 4.18J/g o C First law of thermodynamics: Energy is neither created nor destroyed. Therefore, heat lost by one = the heat gained by other. -q 1 = +q 2 Hess’s Law : Δ H is a state function, so path not important. Rules for changing enthalpy equations: o If the enthalpy equation is reversed, sign of Δ H is reversed. o If the enthalpy equation is halved, quantity of Δ H is halved. o If the enthalpy equation is doubled, quantity of Δ H is doubled. Enthalpy of Formation (more correctly, “Standard enthalpy of formation”) = Δ H o f o is for formation of one mole of product o from its elements o all reactants (i.e., elements) and products in their standard states (1 atm & 298K) o all reactants (i.e., elements) in their most stable form, e.g., oxygen = O 2(g) ; mercury = Hg (l) ; sodium = Na (s) o Δ H o f of elements = 0 Enthalpy of reaction : o Δ H o Rx = n p Δ H o f,products - n r Δ H o f,reactants o Δ H o Rx = n p Δ H o bf,products - n r Δ H o bf reactants
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Chem 161-2011 Lecture 18 4 Δ H Rx We already know two ways to determine Δ H Rx .
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