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Question of the Day 1. A 1.0 mole sample of ethanol (C2H5OH) has a heat capacity of 110.4 J/°C. Calculate the specific heat of ethanol (hint: What is the relationship between heat capacity and specific heat?) Day 6 2-22 = 2.4 J/g°C Thermodynamics Energy Enthalpy Entropy Spontaneity 2 State Functions • Property determined under specified conditions of T, P, location, etc • Independent of HOW conditions are reached 3 State Functions Examples Energy d B B _____________________________________ DE = difference between A & B _____________________________________ A A 4 State Functions: Hess’s Law • For a reaction that occurs in several steps, the sum of the D’s for each step is the D for the net reaction. • Thus, it is possible to calculate D for a complex reaction by adding the D’s of each simple reaction that leads to the complete reaction. 5 State Functions: Hess’s Law • The Equation DTOTAL = Dstep 1 + Dstep 2 + Dstep 3 + … + Dstep n 6 Energy Transfers: • Possibilities: • From system to surroundings • From surroundings to system • How heat transfer is made: • Heat = q • Positive change = heat added to system • Negative change = heat released from system • Work = w • Positive change = work done ON system • Negative change = work done BY system 7 Energy Transfers • First Law of Thermodynamics • Energy lost by the system EQUALS energy gained by the surroundings • Energy gained by the system EQUALS energy lost by the surroundings 8 Energy Transfers • Exothermic Change system • Energy moves from ______________ to surroundings ________________ Loss of q • _________ • Temperature of surroundings increases ________________ 9 Energy Transfers • Endothermic Change surroundings • Energy moves from ________________ system to _____________ Gain of q • _________ • Temperature of surroundings decreases ________________ 10 Energy Transfer Equations • Internal Energy: DE = Ef – Ei • Ef = final energy • Ei = initial/starting energy • Energy Change due to Heat & Work: DE = q + w • • • • Use Use Use Use + q if endo – q if exo + w if work is done ON system – w if work is done BY system 11 Assignment Review section 17.1 and complete #s 5-11 on page 561. - Due Monday 2-22 Assignment Review section 17.1 and complete #s 42, 43, 45, 47, 48, and 50-52 on page 586. - Due Thursday 2-25 Question of the Day 1. The specific heat of ice is 2.1 J/g°C. What is the heat capacity for a 2.0 mole sample? Day 1 2-23 = 75.6 J/°C Enthalpy • Thermodynamic property that measures/describes the changes in heat content of a system under constant pressure. • Symbol: H, DH • Units of measure: cal, kcal, J, kJ • State function • Equations and usage DH = Hf - Hi +DH Hf > Hi Heat gained & endo -DH Hf < Hi Heat lost & exo15 BOOK - Calorimetry •Constant-Pressure Calorimeters The enthalpy (H) of a system accounts for the heat flow of the system at constant pressure. • The heat absorbed or released by a reaction at constant pressure is the same as the change in enthalpy, symbolized as ΔH. BOOK - Calorimetry •Constant-Pressure Calorimeters The value of ΔH of a reaction can be determined by measuring the heat flow of the reaction at constant pressure. • In this textbook, the terms heat and enthalpy change are used interchangeably. • In other words, q = ΔH. Calculating Enthalpy Change in Reactions • Formula: DHRXN = DHPRODUCTS - DHREACTANTS 2 H2(g) + O2(g) Higher E means __________________________ Less stability + 483.6 – Lower E means __________________________ more stability 2 H2O(g) 18 Calculating Enthalpy Change in Reactions CH4(g) + 2O2(g) + H CH4(g) + 2O2(g) 802 – CO2(g) + 2H2O(g) + 890 – CO2(g) + 2H2O(l) 19 Calculating Enthalpy Change in Reactions •Phase matters!!!!!!! •Heat In = Heat Out 20 Thermochemical Equations •Thermochemical Equations How can you express the enthalpy change for a reaction in a chemical equation? In a chemical equation, the enthalpy change for the reaction can be written as either a reactant or a product. Thermochemical Equations In the equation describing the exo reaction of calcium oxide and water, the enthalpy change can be considered a product. CaO(s) + H2O(l) → Ca(OH)2(s) + 65.2 kJ Calcium oxide is one of the components of cement. Thermochemical Equations A chemical equation that includes the enthalpy change is called a thermochemical equation. CaO(s) + H2O(l) → Ca(OH)2(s) + 65.2 kJ Thermochemical Equations •Heats of Reaction The heat of rxn is the ΔH for the chemical equation exactly as it is written. • Heats of reaction are reported as ΔH. • The physical state of the reactants and products must also be given. • The standard conditions are that the reaction is carried out at 101.3 kPa (1 atm) and 25°C. Thermochemical Equations •Heats of Reaction Each mole of calcium oxide and water that reacts to form calcium hydroxide produces 65.2 kJ of heat. CaO(s) + H2O(l) → Ca(OH)2(s) ΔH = –65.2 kJ • In exothermic processes, the chemical potential energy of the reactants is higher than the chemical potential energy of the products. Thermochemical Equations •Heats of Reaction Baking soda decomposes when it is heated. This process is endothermic. 2NaHCO3(s) + 85 kJ → Na2CO3(s) + H2O(l) + CO2(g) The carbon dioxide released in the reaction causes muffins to rise while baking. Thermochemical Equations •Heats of Reaction 2NaHCO3(s) + 85 kJ → Na2CO3(s) + H2O(l) + CO2(g) Remember that ΔH is positive for endothermic reactions. Therefore, you can write the reaction as follows: 2NaHCO3(s) → Na2CO3(s) + H2O(l) + CO2(g) ΔH = 85 kJ Thermochemical Equations •Heats of Reaction The amount of heat released or absorbed during a rxn depends on the number of moles of the reactant involved. • The decomposition of 2 mol of sodium bicarbonate requires 85 kJ of heat. • Therefore, the decomposition of 4 mol of the same substance would require twice as much heat, or 170 kJ. Thermochemical Equations •Heats of Reaction To see why the physical state of the reactants and products must be stated, compare the following two equations. H2O(l) → H2(g) + H2O(g) → H2(g) + O2(g) ΔH = 285.8 kJ 1 2 ΔH = 241.8 kJ 1 2 O2(g) difference = 44.0 kJ Sample Problem 17.4 Using the Heat of Reaction to Calculate Enthalpy Change 2NaHCO3(s) + 85 kJ → Na2CO3(s) + H2O(l) + CO2(g) •Calculate the amount of heat (in kJ) required to decompose 2.24 mol NaHCO3(s). Assignment Review section 17.1 and complete #s 42, 43, 45, 47, 48, and 50-52 on page 586. - Due Thursday 2-25 Assignment: #s 14 and 15 on page 567 (show work), #s 16, 17, 20, and 21 on page 568 (show work) - Due Wednesday 2-17 Question of the Day 1. The appropriate unit for specific heat = 2. Ice subliming (solid gas) is exo or endo? Day 5 2-13 PRACTICE: 2NaHCO3(s) + 85 kJ → Na2CO3(s) + H2O(l) + CO2(g) 1. Calculate the amount of heat (in kJ) required to produce 3 mols of Na2CO3(s). = 255 KJ 1. If a sample of NaHCO3(s) requires 127.5 kJ to completely decompose, what was the mass of the original sample? = 252 g Question of the Day: 3C2H2(g) C6H6(l) + 630 kJ Is this reaction endo OR exo? How much heat is involved if you start with 2 grams of acetylene (acetylene = C2H2)? Day 6 217 EXO 16.2 kJ Thermochemical Equations •Heats of Combustion The heat of combustion is the heat of reaction for the complete burning of one mole of a substance. Thermochemical Equations •Heats of Combustion Small amounts of natural gas within crude oil are burned off at oil refineries. CH4(g) + 2O2(g) → CO2(g) + 2H2O(l) + 890 kJ • This is an exothermic reaction. • Burning 1 mol of methane releases 890 kJ of heat. • The heat of combustion (ΔH) for this reaction is –890 kJ per mole of methane burned. Heats of Combustion at 25°C Interpret Data ΔH (kJ/mol) Substance Formula Hydrogen H2(g) –286 Carbon C(s, graphite) –394 Methane CH4(g) –890 Acetylene C2H2(g) –1300 Ethanol C2H6O(l) –1368 Propane C3H8(g) –2220 Glucose C6H12O6(s) –2808 Octane C8H18(l) –5471 Sucrose C12H22O11(s) –5645 Like other heats of reaction, heats of combustion are reported as the enthalpy changes when the reactions are carried out at 101.3 kPa and 25°C. •Which of the following thermochemical equations represents an endothermic reaction? A. Cgraphite(s) + 2 kJ B. 2H2(g) + O2(g) Cdiamond(s) 2H2O + 483.6 kJ Vocab. from the book: Heats of reaction Heats of combustion Molar heat of fusion (melting) Molar heat of solidification (freezing) ΔHfus = -ΔHsolid Vocab. from the book: Molar heat of vaporization Molar heat of condensation ΔHvap = -ΔHcond Molar heat of solution (dissolving) Spontaneity • Main question of chemistry: Will a reaction GO! when reactants are put together? • Spontaneous reactions • Reaction occurs by itself, activation E may be needed • May have fast, moderate, or slow rate • Non-spontaneous reactions • Require constant E supply to occur • Reaction stops when E supply is removed • Reversing reactions • If a reaction is spontaneous under specified conditions, the reverse is non-spontaneous at those same conditions 42 Predicting Sponaneity • Why is DH a pretty good predictor of spontaneity? • -DH : heat out • +DH: heat supply needed • When might DH not indicate spontaneity well? 43 Predicting Spontaneity • H2O(s) → H2O(l) ∆H = + 6.0 kJ; 1 atm, 0°C 44 Predicting Spontaneity • H2O(l) → H2O(g) ∆H = + 40.7 kJ; 1 atm, 100°C 45 Predicting Spontaneity • CaCO3(s) → CaO(s) + CO2(g) ∆H = +178.0 kJ; 298 K vs. 1100 K 46 Predicting Spontaneity 47 Question of the Day: Day 1 2-18 NaCl(s) + 3.88kJ NaCl(aq) Is this reaction endo OR exo? Draw an enthalpy diagram representing the change above: NaCl(aq) + 3.88 kJ NaCl(s) – Entropy • Entropy is the measure of the or randomness of a system. disorder • Entropy is a state function. • Entropy changes follow Hess’ Law. • A system has high entropy if it • _____________________________________________ (Very disorganized) • _______________________________________ Is Very dsorigniaezd Has freedom of motion 49 Entropy & Spontaneity • Reactions will be spontaneous etnopry if ________ is increased. • Examples • ice cubes melting vs. putting water in freezer to make ice cubes • room gets messy vs. doing work to put things in order 50 Standards & Constants • Conditions & Units • T = 25oC = 298K • P = 1 atm • Units: J/K = Joules/Kelvin • What is the relationship between entropy and temperature? • As T 0 K, S 0 • T↓, S↓ and T↑, S↑ 51 Standards & Constants • Why are the entropy values for elements and compounds always positive? • T always > 0K • Therefore everything is always moving • How do entropies compare among the phases of matter? • So solid < So liquid < So gas 52 Standard Entropy Changes • The equation ∆So = SSoproducts – SSoreactants • Example: 2HCI(aq) + 2Ag(s) → 2AgCl(s) + H2(g)1 atm, 25oC • Predict if this reaction will have + ∆So or – ∆So. • Calculate ∆So (NEED STANDARD VALUE HANDOUTS). 53 Standard Entropy Changes • The equation DSo = SnSoproducts – SmSoreactants • Example: 2 H2(g) + O2(g) → 2 H2O(g) 1 atm, 25oC • Predict if this reaction will have + ∆So or – ∆So. • Calculate ∆So. 54 Quick Talk: Enthalpy ΔS q State function Entropy ΔH Joules 4.184 J Potential Chemical Specific Energy heat Assignment: #s 14 and 15 on page 567 (show work), #s 16, 17, 20, and 21 on page 568 (show work) - Due Thursday 2-19 (pd 1) Wednesday 2-18 (pd 3)