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Exam over Chapters 8 & 9 Chapter ten notes Section 10.1 Section 10.2 Section 10.3 Section 10.4 Section 10.5 Section 10.6 Section 10.7 Section 10.8 System: Part we care about Reactants & Products Surroundings: Everything else in the universe Aopen system (mass and heat pass through) Bclosed system (heat only pass through) Cisolated system (no heat or mass transfer) For chemical reactions to happen spontaneously, the final products must be more stable than the starting reactants Higher energetic substances Typically less stable, more reactive Lower energetic substances Typically more stable, less reactive Thermal energy flows from warmer to cooler H2O(s) H2O(l) 2H2(g) + O2(g) 2H2O(l) Study of heat and its transformations into other energies Thermochemistry is a part of this Thermodynamics studies changes in the state of a system State functions are properties that are determined by the state of the system, regardless of how it was achieved Final – Initial Ex: ▪ Energy ▪ Pressure ▪ Volume ▪ Temperature Has 2 components: Kinetic energy: various types of molecular and electron motion Potential energy: attractive and repulsive interactions between atoms and molecules ΔU = U(products) – U(reactants) ΔU = q + w q = heat (absorbed or released by the system) w = work (done on or by the system) Calculate the overall change in internal energy (ΔU) for a system that absorbs 188 J of heat and does 141 J of work on its surroundings. Convert 723.01 J into calories SKETCH and LABEL what an exothermic and endothermic energy vs. time graph would look like. Calculate the overall change in internal energy for a system that releases 43 J in heat and has 37 J of work done on it by its surroundings Reactions can be carried out in two ways: In a closed container (constant volume): qv = ΔU In an open container (constant pressure): qp = Δ H Combustion of propane gas: ΔH = H(products) – H(reactants) “+” = endothermic “—” = exothermic H2O(s) H2O(l) ΔH = +6.01 kJ/mol CH4(g) + 2O2(g) CO2(g) + 2H2O(l) ΔH = -890.4 kJ/mol CH4(g) + 2O2(g) CO2(g) + 2H2O(l) ΔH = -890.4 kJ/mol How much energy is release from 18.4 g of methane being burned? If 924.3 kJ of energy was released, how many grams of water was produced? If you change the AMOUNTS in a balanced equation, you change the enthalpy the same way 1) Ex: if coefficients are doubled, so is the enthalpy If you reverse the equation, you reverse the sign of the ΔH 2) Ex: H2O(s) H2O(l) ΔH = +6.01 kJ/mol H2O(l) H2O(s) ΔH = -6.01 kJ/mol Measurement or heat changes within a system Using a calorimeter Specific Heat (s): amount of heat required to raise the temperature of 1 g of a substance by 1°C (ex: liquid water is 4.184 J/(g*°C) q = (s)(m)(ΔT) Heat Capacity (C): amount of heat required to raise the temperature of an object by 1°C q = (C)(ΔT) What is the amount of heat (in kJ) required to heat 255 g of water from 25.2 °C to 90.5 °C? Can calculate changes in heat using styrofoam cups and known mass of water Assuming constant pressure Therefore… qp = msΔT = ΔH System: reactants and products (the reaction) Surroundings: water in calorimeter For an exothermic reaction: The system loses heat The surroundings gain (absorb) heat A 30.4-g piece of unknown metal is heated up in a hot bath to a temperature of 92.4°C. The metal is then placed in a calorimeter containing 100. g of water at 25.0°C. After the calorimeter is capped, the temperature of the calorimeter raises to 27.2°C. What was the specific heat of the unknown metal? Ex: 50.0 mL of 1.00 M HCl and 50.0 mL of 1.00 M NaOH are mixed in a calorimeter with 100 g of water and capped at room temp (25°C). The reaction reaches a max of 31.7°C. What is the ΔH°rxn? 125.0-g of a metal is heated to 100.0°C. It is then placed into a calorimeter containing 100.0 mL (100.0 g) of water at 25.0°C and capped. The energy is transferred and the max temperature of 34.1°C is reached. What is the specific heat of the metal? Given the following, determine the ΔH for 3H2(g) + O3(g) 3H2O(g) Standard Enthalpy of Formation (ΔH°f): heat change that results when 1 mole of a compound is formed from its constituent elements in their standard states “Standard State” means “stable form” 1 atm and 25°C typically Example: O(g) (249.4), O2(g) (0), O3(g) (142.2) ΔH°rxn: enthalpy of a reaction under standard conditions When we know reactions go to completion or can be done in one step, we can use a direct method Ex: Calculate ΔH°rxn for 2SO(g) + 2/3O3(g) 2SO2(g) From Appendix 2: SO(g): (5.01), O3(g): (142.2), SO2(g): (-296.4) When a reaction is too slow or side reactions occur, enthalpy of reaction can be calculated using Hess’s Law Recall: when bonds are made, energy is given off (exo); when bonds break, energy is needed (endo) Bond Enthalpy: the measure of stability of a molecule Enthalpy change associated with breaking a particular bond in 1 mole of gaseous molecules ▪ H2(g) H(g) + H(g) ΔH = 436.4 kJ/mol The higher the bond enthalpy, the stronger the bond The bonds in different compounds have different bond enthalpies Ex: O—H bond in water vs. O—H bond in methanol are different Therefore, we speak of AVERAGE bond enthalpy Recall: amount of energy required to convert 1 mole of ionic solid to its constituent ions in the gas phase Ex: NaCl(s) Na+(g) + Cl-(g)