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Principles of Reactivity: Energy and Chemical Reactions

Principles of Reactivity: Energy and Chemical Reactions. Chapter 6. Energy: Some Basics. From Physics: Force – a kind of push or pull on an object. Energy – the capacity to do work. Work – force applied over a distance w = F  d

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Principles of Reactivity: Energy and Chemical Reactions

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  1. Principles of Reactivity: Energy and Chemical Reactions Chapter 6 Chapter 6

  2. Energy: Some Basics From Physics: Force – a kind of push or pull on an object. Energy – the capacity to do work. Work – force applied over a distance w = Fd Heat – energy transferred from a warmer object to a cooler object. Chapter 6

  3. Energy: Some Basics Kinetic and Potential Energy Kinetic Energy (Thermal Energy) – energy due to motion. Chapter 6

  4. Energy: Some Basics Kinetic and Potential Energy Potential Energy (Stored Energy) – the energy an object possesses due to its position. • Potential energy can be converted into kinetic energy. Example: a ball of clay dropping off a building. Chapter 6

  5. Energy: Some Basics First Law of Thermodynamics “The total amount of energy in the universe is fixed.” Also referred to as the “Law of Conservation of Energy” Chapter 6

  6. Energy: Some Basics Temperature and Heat Temperature is a measure of heat energy • Heat is not the same as temperature. • The more thermal energy a substance has the greater its molecular motion (kinetic energy). • The total thermal energy in an object is the sum of the energies of all the “bodies” in the object. Chapter 6

  7. Energy: Some Basics Systems and Surroundings System – portion of the universe we wish to study. Surroundings – everything else. Universe = System + Surroundings Chapter 6

  8. Energy: Some Basics Directionality of Heat Heat energy always flows from the hot object to the cold object. - this flow continues until the two objects are at the same temperature (thermal equilibrium). Chapter 6

  9. Energy: Some Basics Directionality of Heat Exothermic – Heat is transferred from the system to the surroundings (object will feel “hot”). Endothermic – Heat is transferred to the system from the surroundings (object will fell “cold”). Chapter 6

  10. Energy: Some Basics Energy Units SI Unit for energy is the joule, J: A more traditional unit is the Calorie Calorie (cal) – amount of energy required to raise 1.0 g of water 1oC. 1cal = 4.184J Chapter 6

  11. Specific Heat Capacity The amount of heat transferred is dependant on three quantities: • Quantity of material • Size of temperature change • Identity of the material Chapter 6

  12. Specific Heat Capacity q = energy c = specific heat capacity DT = temperature change Chapter 6

  13. Specific Heat Capacity exothermic -DT -q endothermic +DT +q Chapter 6

  14. Specific Heat Capacity • Specific heat capacity can be either per gram (J/g(oC) or per mole (J/mol(oC). • The smaller a substances specific heat capacity, the better a thermal conductor it is. Chapter 6

  15. Energy and Changes of State Chapter 6

  16. Energy and Changes of State • In the previous slide there is a continuous, steady application of energy. • The sections that show increasing temperature are the result of the particular phase being warmed. q = cm(DT) • The “flat” sections occur when all the applied energy is used to change the phase of the substance. • Fusion – solid  liquid • Vaporization – liquid  gas Chapter 6

  17. Energy and Changes of State • The energy required to change the phase of a substance is unique and is described in a physical constant. • Solid  Liquid • Heat of Fusion (water, 333J/g) • Liquid  Gas • Heat of Vaporization (water, 2256J/g) • These constants can be used to determine the energy used in melting or vaporizing a substance. q = (Heat of Fusion)(mass of sample) q = (Heat of Vapor.)(mass of sample) Chapter 6

  18. Energy and Changes of State q = cm(DT) q = (Heat of Vapor.)(mass) Chapter 6

  19. First Law of Thermodynamics Internal Energy Internal Energy – sum of all kinetic and potential energy in an object. • It is very hard to determine an objects internal energy, but it is possible to determine the change in energy (DE). • Change in internal energy, DE = Efinal - Einitial • A positive DE means Efinal > Einitialor the system gained energy from the surroundings (endothermic) • A negative DE means Efinal < Einitialor the system lost energy to the surroundings (exothermic) Chapter 6

  20. First Law of Thermodynamics • Relating DE to Heat and Work • DE = q + w • q = heat w = work • Both heat energy and work can change a systems internal energy. Chapter 6

  21. First Law of Thermodynamics • State Functions • State function – a process that is determined by its initial and final conditions. Chapter 6

  22. First Law of Thermodynamics • State Functions • State function – a process that is determined by its initial and final conditions. • “A process that is not path dependant.” • Work (w) and heat (q) are not state functions. • Energy change (DE) is a state function. Chapter 6

  23. First Law of Thermodynamics • Enthalpy (H) - Heat transferred between the system and surroundings carried out under constant pressure. • DE = q + w • Most reactions occur under constant pressure, so • DE = q + (-P(DV)) • If volume is also constant, DV = 0 • DE = qp • So, Energy change is due to heat transfer, • DE = DH = qp Chapter 6

  24. Enthalpy Enthalpy Change (DH) – The heat evolved or absorbed in a reaction at constant pressure DH = Hfinal - Hinitial = qP Chapter 6

  25. Enthalpy • Enthalpy Change (DH) – The heat evolved or absorbed in a reaction at constant pressure • H and DH are state functions, depending only on the initial and final states. Chapter 6

  26. Enthalpies of Reaction 2 H2(g) + O2(g)  2 H2O(g) DH = -483.6 J Chapter 6

  27. Enthalpies of Reaction • For a reaction • Enthalpy is an extensive property (magnitude DH is directly proportional to amount): • CH4(g) + 2O2(g)  CO2(g) + 2H2O(g) DH = -802 kJ • 2CH4(g) + 4O2(g)  2CO2(g) + 4H2O(g) DH = -1604 kJ Chapter 6

  28. Enthalpies of Reaction • For a reaction • Enthalpy is an extensive property (magnitude DH is directly proportional to amount): • When we reverse a reaction, we change the sign of DH: • CO2(g) + 2H2O(g)  CH4(g) + 2O2(g) DH = +802 kJ • CH4(g) + 2O2(g)  CO2(g) + 2H2O(g) DH = -802 kJ Chapter 6

  29. Enthalpies of Reaction • For a reaction • Enthalpy is an extensive property (magnitude DH is directly proportional to amount): • When we reverse a reaction, we change the sign of DH: • Change in enthalpy depends on state: • CH4(g) + 2 O2(g)  CO2(g) + 2 H2O(g)DH = -802 kJ • CH4(g) + 2 O2(g)  CO2(g) + 2 H2O(l)DH = -890 kJ Chapter 6

  30. Enthalpies of Reaction 2 Mg(s) + O2(g)  2 MgO(s) DH = -1205 kJ a) Is this reaction endothermic or exothermic? Chapter 6

  31. Enthalpies of Reaction 2 Mg(s) + O2(g)  2 MgO(s) DH = -1205 kJ a) Is this reaction endothermic or exothermic? Exothermic, this is indicated by the negative DH. Chapter 6

  32. Enthalpies of Reaction • 2 Mg(s) + O2(g)  2 MgO(s) DH = -1205 kJ • Calculate the amount of heat transferred when 2.4g of Mg reacts at constant pressure. Chapter 6

  33. Enthalpies of Reaction • 2 Mg(s) + O2(g)  2 MgO(s) DH = -1205 kJ • Calculate the amount of heat transferred when 2.4g of Mg reacts at constant pressure. Chapter 6

  34. Enthalpies of Reaction • 2 Mg(s) + O2(g)  2 MgO(s) DH = -1205 kJ • Calculate the amount of heat transferred when 2.4g of Mg reacts at constant pressure. Chapter 6

  35. Enthalpies of Reaction • 2 Mg(s) + O2(g)  2 MgO(s) DH = -1205 kJ • Calculate the amount of heat transferred when 2.4g of Mg reacts at constant pressure. Chapter 6

  36. Enthalpies of Reaction 2 Mg(s) + O2(g)  2 MgO(s) DH = -1205 kJ c) How many grams of MgO are produced during an enthalpy change of 96.0 kJ? Chapter 6

  37. Enthalpies of Reaction 2 Mg(s) + O2(g)  2 MgO(s) DH = -1205 kJ c) How many grams of MgO are produced during an enthalpy change of 96.0 kJ? Chapter 6

  38. Enthalpies of Reaction 2 Mg(s) + O2(g)  2 MgO(s) DH = -1205 kJ c) How many grams of MgO are produced during an enthalpy change of 96.0 kJ? Chapter 6

  39. Enthalpies of Reaction 2 Mg(s) + O2(g)  2 MgO(s) DH = -1205 kJ c) How many grams of MgO are produced during an enthalpy change of 96.0 kJ? Chapter 6

  40. Enthalpies of Reaction 2 Mg(s) + O2(g)  2 MgO(s) DH = -1205 kJ d) How many kilojoules of heat are absorbed when 7.50g of MgO is decomposed into Mg and O2 at constant pressure? Chapter 6

  41. Enthalpies of Reaction 2 MgO(s)  2 Mg(s) + O2(g) DH = 1205 kJ d) How many kilojoules of heat are absorbed when 7.50g of MgO is decomposed into Mg and O2 at constant pressure? Chapter 6

  42. Enthalpies of Reaction 2 MgO(s)  2 Mg(s) + O2(g) DH = 1205 kJ d) How many kilojoules of heat are absorbed when 7.50g of MgO is decomposed into Mg and O2 at constant pressure? Chapter 6

  43. Enthalpies of Reaction 2 MgO(s)  2 Mg(s) + O2(g) DH = 1205 kJ d) How many kilojoules of heat are absorbed when 7.50g of MgO is decomposed into Mg and O2 at constant pressure? Chapter 6

  44. Calorimetry Constant-Pressure Calorimetry Chapter 6

  45. Calorimetry • Constant-Pressure Calorimetry • Atmospheric pressure is constant! • DH = qP • qsystem = -qsurroundings • The surroundings are composed of the water in the calorimeter and the calorimeter. • qsystem = -(qwater + qcalorimeter) Chapter 6

  46. Calorimetry • Constant-Pressure Calorimetry • Atmospheric pressure is constant! • DH = qP • qsystem = -qsurroundings • The surroundings are composed of the water in the calorimeter and the calorimeter. • For most calculations, the qcalorimeter can be ignored. • qsystem = - qwater • csystemmsystem DTsystem = - cwatermwater DTwater Chapter 6

  47. Calorimetry Bomb Calorimetry (Constant-Volume Calorimetry) Chapter 6

  48. Calorimetry • Bomb Calorimetry • (Constant-Volume Calorimetry) • Special calorimetry for combustion reactions • Substance of interest is placed in a “bomb” and filled to a high pressure of oxygen • The sealed bomb is ignited and the heat from the reaction is transferred to the water • This calculation must take into account the heat capacity of the calorimeter (this is grouped together with the heat capacity of water). • qrxn = -Ccalorimeter(DT) Chapter 6

  49. Calorimetry NH4NO3(s)  NH4+(aq) + NO3-(aq) DTwater = 16.9oC – 22.0oC = -5.1oC mwater = 60.0g cwater = 4.184J/goC msample = 4.25g qsample = -qwater qsample = -cwatermwater DTwater qsample = -(4.184J/goC)(60.0g)(-5.1oC) qsample = 1280.3J - Now calculate DH in kJ/mol Chapter 6

  50. Calorimetry NH4NO3(s)  NH4+(aq) + NO3-(aq) DTwater = 16.9oC – 22.0oC = -5.1oC mwater = 60.0g cwater = 4.184J/goC msample = 4.25g qsample = 1280.3J moles NH4NO3 = 4.25g/80.032g/mol = 0.0529 mol DH = qsample/moles DH = 1280.3J/0.0529mol DH = 24.2 kJ/mol Chapter 6

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