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Unit 1: Energy Changes and Rates of Reaction

Unit 1: Energy Changes and Rates of Reaction . Chapters 5: Thermochemistry Chapter 6: Chemical Kinetics. Different forms of energy. Law of Conservation of Energy. energy cannot be created or destroyed, only transformed from one form to another (1 st Law of Thermodynamics). Thermochemistry.

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Unit 1: Energy Changes and Rates of Reaction

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  1. Unit 1: Energy Changes and Rates of Reaction Chapters 5: Thermochemistry Chapter 6: Chemical Kinetics

  2. Different forms of energy

  3. Law of Conservation of Energy • energy cannot be created or destroyed, only transformed from one form to another (1st Law of Thermodynamics)

  4. Thermochemistry • study of energy changes in physical, chemical and nuclear changes

  5. Energy and Molecules Kinetic Energy: energy of motion/work (i.e. collision theory). Substances may have Thermal Energy (energy from motion of molecules- amount of substance matters). Temperature is the average kinetic energy of the particle (amount of substance negated). Potential Energy: (due to position and composition) stored in molecules because of the arrangement of nuclei and electrons in its atoms, electrical forces b/w charged particles

  6. Endo vs. Exothermic • Which process requires energy? releases energy? breaking bonds or forming bonds • What is the net E change if: a) E (required to break bonds) > E (released when bonds are formed)? net change = E is absorbed (endothermic) reactants + energy  products b) E (required to break bonds) < E (released when bonds are formed)? net change = E is released (exothermic) reactants  products + energy

  7. IMPORTANT TERMINOLOGY Surroundings System Exchanges matter and energy Universe • Chemical System: substances (reactants and products) we are studying, represented by the chemical equations Open System: Energy and matter can move in or out Isolated System: an ideal system where no energy or matter can move in or out Closed System: only energy can move in or out, but not matter • Surroundings: all the matter around the system that absorbs or releases heat

  8. 3 TYPES OF SYSTEMS Identify the systems below:

  9. Why is E released during an exothermic reaction? • What is the source of E released during a reaction?

  10. Discuss the combustion of gasoline in terms of PE, KE, energy absorbed, energy released, net energy change, exo or endothermic, work. gasoline + oxygen  carbon dioxide + water + energy PE of bonds in reactants > PE of bonds in products KE of reactants < KE of products E absorbed to break bonds < E released when bonds are formed net energy change is release (exothermic) High KE of product gases does work on car parts to make the car move.

  11. Homework: • p. 282 #1,2,4,5 • p. 285 A-C • p. 291 #1,2,4,7-9

  12. Follow-up to 5.1 • What is the difference between thermal energy and temperature of a substance? • Compare the potential energies of the reactants to that of the products for the two types of chemical processes below: (a) endothermic process   (b) exothermic process 3. Label the following reactions as “endothermic,” “exothermic,” or “not enough information.” (a) O2 → O + O (b) O2 + O → O3 (c) H2 + Br2 → 2HBr (d)NaCl → Na+ + Cl− 4. Classify each of the following as an open system or a closed system. • a pot of boiling water • a sealed bottle of water • a helium balloon • a hot-air balloon

  13. How do we measure energy changes during a physical or chemical change? A Simple Calorimeter Key Assumptions: 1. Any thermal energy transfer b/w the calorimeter and the universe is negligible. 2. Any thermal energy absorbed by the calorimeter itself is negligible. 3. All dilute, aqueous solutions have d = 1.00 g/mL and c = 4.18 J/g°C.

  14. Calculations Involving Thermal Energy Transfer q=mcΔT q > 0 exo q < 0 endo Understanding the formula: q is thermal energy transferred, in J or kJ m is mass, in g ΔT = T2-T1 the temperature change, in °C c is the specific heat capacity, in J/(g°C) (Table 1 p. 292)

  15. Given Required Analysis Solution Paraphrase p. 295: Tutorial 1 Sample Problem 1: A student places 50.0 mL of liquid water at 21.00°C into a calorimeter. She places a sample of gold at 100.00°C into the calorimeter. The final temp of the water is 21.33°C. Calculate the quantity of thermal energy, q, absorbed by the water in the calorimeter.

  16. Sample Problem 2 Using the value of q from #1, calculate the specific heat capacity of the sample of gold if its mass is 6.77 g. Assume that the final temp of the gold sample was the same as the final temp of the water in the calorimeter.

  17. Sample Problem 3 A 50.0 mL sample of 1.0 mol/L HCl(aq) was mixed with 50.0 mL of 1.0 mol/L NaOH(aq) at 25°C in a calorimeter. After the solutions were mixed by stirring, the temperature was 31.9°C. Determine the quantity of thermal energy transferred by the reaction to the water and state whether the reaction was endothermic or exothermic.

  18. p. 297 #1-3

  19. Mini-Investigation: p. 297

  20. Demos less stable • KMnO4 and glycerol in fume hood • Ba(OH)2(s) + NH4SCN(s) system heat (qsurr) enthalpy change ΔHsys more stable surroundings

  21. Re-draw for an endothermic rxn!

  22. Enthalpy Change • Internal energy - sum of all KE and PE of all the components of the system • We can’t measure internal energy, so… • measure enthalpy change – the energy absorbed from or released to the surroundings during a reaction Recall: 2 types of energy Potential Energy includes: Protons and neutrons within nuclei Electrons in bonds Kinetic energy includes: Moving e- within atoms Vibrational, rotational, and translational motion

  23. ENTHALPY CHANGE (ΔH)aka heat of reaction, enthalpy of reaction, change in heat content ΔHsystem= - qsurroundings Calorimetry Calculations: qsurr = mcΔT ΔH = -q enthalpy per gram = ΔH/m molar enthalpy = ΔH/n where n = m/M ΔH > 0 , q < 0 Endothermic ΔH < 0 , q > 0 Exothermic ΔH° means @SATP: 100kPa, pure liquids/solids, 1 M, 25°C or 298K

  24. MOLAR ENTHALPY (ΔHX) ΔHx =ΔH n ΔH =nΔHx n = ΔH ΔHx ΔHX (molar enthalpy) energy released/absorbed by system per mole of reacting substance units of kJ/mol good for reference or comparison ΔH (enthalpy change) energy released/absorbed by system as R  P units of kJ depends on amount of substance reacting

  25. Mini-Lab p. 285 • Calculate the molar enthalpy of solution for the salt used in the expt.

  26. Molar Enthalpy Examples Ex 1: The molar enthalpy of vaporization of Freon-12 is 34.99 kJ/mol. If 100.0 g of Freon-12 (M = 120.91 g/mol) is vaporized, calculate the expected enthalpy change. Ex 2: 50.0 mL of 0.300 mol/L CuSO4(aq) is mixed with and equal volume of 0.600 mol/L NaOH(aq). The initial temperature of both solutions is 21.4°C. After mixing, the highest temperature reached is 24.6°C. Determine the enthalpy change of the reaction.

  27. Ex 3: A chemist wants to determine the molar enthalpy of neutralization for the following reaction. HCl(aq) + NaOH(aq)NaCl(aq) + H2O(l) The chemist uses a coffee cup calorimeter to neutralize 61.1 mL of 0.543 mol/L HCl(aq) with 42.6 mL of NaOH(aq). The initial temperature of both solutions is 17.8 °C. After neutralization, the highest recorded temperature is 21.6°C. Calculate the molar enthalpy of neutralization, in kJ/mol HCl. Ex 4: What mass of KCl must have dissolved if the temperature of 200.0 g of water increased by 5.5°C? (ΔHsol’n,KCl = 1.7 x 104 J/mol)

  28. Investigation 5.2.1 • Molar Enthalpy of a Chemical Change • p. 333 • Title • Purpose • Observations • Analyze and Evaluate (a-g) • Apply and Extend (h,i)

  29. Homework Practice Problems • p. 301 #1-4 Section Review • p. 306 #1,2,4

  30. REPRESENTING ENTHALPY CHANGES4 Methods- Exothermic= product side Endothermic= reactant side Equation must be balanced Watch +/- signs Molar Enthalpy!!! (kJ/mol) Standard molar enthalpies (ΔH°x ) are @ SATP 1. Include energy value in the thermochemical equation H2(g) + ½ O2H2O(l) + 285.8 kJ 2. Write the chemical equation and enthalpy change H2(g) + ½ O2H2O(l) ΔH= - 285.8 kJ 3. State the molar enthalpy for a specific reaction & substance ΔHForm= - 285.8 kJ/mol H2O(l) 4. Draw a chemical potential energy diagram

  31. REPRESENTING ENTHALPY CHANGES4 Methods- Potential Energy Diagram for the Formation of Water (kJ) • Remember: • Title • Label axis • Units • Reactants on LEFT • Products on Right 4. Draw a chemical potential energy diagram Changes to Ep as bonds are broken and formed Exothermic Reactions= Energy released= Ep=products lower than reactants Endothermic Reactants= Energy gained= Ep=products higher than reactants

  32. REPRESENTING ENTHALPY CHANGES4 Methods- What about the reverse reaction, the decomposition of water? 1. Include energy value in the thermochemical equation 2. Write the chemical equation and enthalpy change. 3. State the molar enthalpy for a specific reaction. 4. Draw a chemical potential energy diagram.

  33. REPRESENTING ENTHALPY CHANGES… Ex. What is the thermochemical equation for the following chemical equation and molar enthalpy of combustion for butane: 2 C4H10(g) + 13 O2 (g)  8 CO2(g) + 10 H2O (l) ΔHcomb= -2871 kJ/mol Remember that thermochemical equations contain ΔH values!

  34. Homework: p. 304 #1-4 p. 306 #3,5-7

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