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2016 topic 5.1 measuring energy changes

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Introduction to enthalpy and calorimetry

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2016 topic 5.1 measuring energy changes

  1. 1. Lesson Exothermic and Endothermic Reactions and Calorimetry IB Chemistry Power Points Topic 05 Energetics www.pedagogics.ca
  2. 2. Great thanks to JONATHAN HOPTON & KNOCKHARDY PUBLISHING www.knockhardy.org.uk/sci.htm Much taken from ENTHALPY CHANGES
  3. 3. Background and Review First Law of Thermodynamics (Law of Energy Conservation) Energy can be neither created nor destroyed but it can be converted from one form to another Energy Changes in Chemical Reactions All chemical reactions are accompanied by some form of energy change Exothermic Energy is given out Endothermic Energy is absorbed Examples Exothermic combustion reactions neutralization (acid + base) Endothermic photosynthesis thermal decomposition of calcium carbonate
  4. 4. The heat content of a chemical system is called enthalpy (represented by H). Key Concept We cannot measure enthalpy directly, only the change in enthalpy ∆H i.e. the amount of heat released or absorbed when a chemical reaction occurs at constant pressure. ∆H = H(products) – H(reactants) ∆Ho (the STANDARD enthalpy of reaction) is the value measured when temperature is 298 K and pressure is 100.0 kPa.
  5. 5. If ∆H is negative, H(products) < H(reactants) There is an enthalpy decrease and heat is released to the surroundings. Enthalpy Diagram -Exothermic Change enthalpy
  6. 6. If ∆H is positive, H(products) > H(reactants) There is an enthalpy increase and heat is absorbed from the surroundings. Enthalpy Diagram - Endothermic Change enthalpy
  7. 7. Exothermic reactions release heat Example: Enthalpy change in a chemical reaction N2(g) + 3H2(g)  2 NH3(g) ∆H = -92.4 kJ/mol The coefficients in the balanced equation represent the number of moles of reactants and products.
  8. 8. N2(g) + 3H2(g)  2 NH3(g) ∆H = -92.4 kJ/mol State symbols are ESSENTIAL as changes of state involve changes in thermal energy. The enthalpy change is directly proportional to the number of moles of substance involved in the reaction. For the above equation, 92.4 kJ is released - for each mole of N2 reacted - for every 3 moles of H2 reacted - for every 2 moles of NH3 produced.
  9. 9. The reverse reaction 2 NH3(g)  N2(g) + 3H2(g) ∆H = +92.4 kJ/mol Note: the enthalpy change can be read directly from the enthalpy profile diagram.
  10. 10. Thermochemical Standard Conditions The ∆H value for a given reaction will depend on reaction conditions. Values for enthalpy changes are standardized : for the standard enthalpy ∆Ho -Temperature is 298 K - Pressure is 100 kPa - All solutions involved are 1 M concentration
  11. 11. Calorimetry – Part 1 Specific Heat Capacity The specific heat capacity of a substance is a physical property. It is defined as the amount of heat (Joules) required to change the temperature (oC or K) of a unit mass (g or kg) of substance by ONE degree. Specific heat capacity (SHC) is measured in J g-1 K-1 or kJ kg-1 K-1 (chemistry) J kg-1 K-1 (physics)
  12. 12. Calorimetry – continued Heat and temperature change Knowing the SHC is useful in thermal chemistry. Heat added or lost can be determined by measuring temperature change of a known substance (water). Q = mc∆T heat = mass x SHC x ∆Temp
  13. 13. Calorimetry – Part 2 Applications A calorimeter is used to measure the heat absorbed or released in a chemical (or other) process by measuring the temperature change of an insulated mass of water.
  14. 14. Sample Problem 1 When 3 g of sodium carbonate are added to 50 cm3 of 1.0 M HCl, the temperature rises from 22.0 °C to 28.5 °C. Calculate the heat required for this temperature change.
  15. 15. Sample problem 2: 50.0 cm 3 of a 1.00 mol dm -3 HCl solution is mixed with 25.0 cm 3 of 2.00 mol dm -3 NaOH. A neutralization reaction occurs. The initial temperature of both solutions was 26.7 o C. After stirring and accounting for heat loss, the highest temperature reached was 33.5 o C. Calculate the enthalpy change for this reaction. NaOH HCl both 26.7o . 26.7o  33.5o
  16. 16. After writing a balanced equation, the molar quantities and limiting reactant needs to be determined. Note that in this example there is exactly the right amount of each reactant. If one reactant is present in excess, the heat evolved will associated with the mole amount of limiting reactant. Sample problem 2: 50.0 cm 3 of a 1.00 mol dm -3 HCl solution is mixed with 25.0 cm 3 of 2.00 mol dm -3 NaOH. A neutralization reaction occurs. The initial temperature of both solutions was 26.7 o C. After stirring and accounting for heat loss, the highest temperature reached was 33.5 o C. Calculate the enthalpy change for this reaction.
  17. 17. Next step – determine how much heat was released. There are some assumptions in this calculation - Density of reaction mixture (to determine mass) - SHC of reaction mixture (to calculate Q) Sample problem 2: 50.0 cm 3 of a 1.00 mol dm -3 HCl solution is mixed with 25.0 cm 3 of 2.00 mol dm -3 NaOH. A neutralization reaction occurs. The initial temperature of both solutions was 26.7 o C. After stirring and accounting for heat loss, the highest temperature reached was 33.5 o C. Calculate the enthalpy change for this reaction.
  18. 18. Final step – calculate ΔH for the reaction Sample problem 2: 50.0 cm 3 of a 1.00 mol dm -3 HCl solution is mixed with 25.0 cm 3 of 2.00 mol dm -3 NaOH. A neutralization reaction occurs. The initial temperature of both solutions was 26.7 o C. After stirring and accounting for heat loss, the highest temperature reached was 33.5 o C. Calculate the enthalpy change for this reaction.
  19. 19. Sample problem 3: to determine the enthalpy of combustion for ethanol (see reaction), a calorimeter setup (below) was used. The burner was lit and allowed to heat the water for 60 s. The change in mass of the burner was 0.518 g and the temperature increase was measured to be 9.90 o C. What is the big assumption made with this type of experiment?
  20. 20. First step – calculate heat evolved using calorimetry Last step – determine ΔH for the reaction Sample problem 3: to determine the enthalpy of combustion for ethanol (see reaction), a calorimeter setup (below) was used. The burner was lit and allowed to heat the water for 60 s. The change in mass of the burner was 0.518 g and the temperature increase was measured to be 9.90 o C.
  21. 21. Sample problem 4: 100.0 cm 3 of 0.100 mol dm -3 copper II sulphate solution is placed in a styrofoam cup. 1.30 g of powdered zinc is added and a single replacement reaction occurs. The temperature of the solution over time is shown in the graph below. Determine the enthalpy value for this reaction. First step Make sure you understand the graph. Extrapolate to determine the change in temperature. The extrapolation is necessary to compensate for heat loss while the reaction is occurring. Why would powdered zinc be used?
  22. 22. Determine the limiting reactant Calculate Q Calculate the enthalpy for the reaction. Review Exercise 2 Sample problem 4: 100.0 cm 3 of 0.100 mol dm -3 copper II sulphate solution is placed in a styrofoam cup. 1.30 g of powdered zinc is added and a single replacement reaction occurs. The temperature of the solution over time is shown in the graph below. Determine the enthalpy value for this reaction.

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