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Calorimetry & Enthalpy Changes

Calorimetry & Enthalpy Changes. Learning Goals: I will be able to perform calorimetry and calculate the thermal energy transfers (q) I will learn to calculate enthalpy (H) I will understand the terms calorimetry , enthalpy and bond energy. Do you remember…. What is heat?

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Calorimetry & Enthalpy Changes

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  1. Calorimetry & Enthalpy Changes Learning Goals: I will be able to perform calorimetry and calculate the thermal energy transfers (q) I will learn to calculate enthalpy (H) I will understand the terms calorimetry, enthalpy and bond energy

  2. Do you remember…. • What is heat? • What is the symbol for heat? • What are the three types of energy changes?

  3. Measuring Energy Changes:Calorimetry Calculations

  4. Specific Heat Capacity • The amount of energy required to raise the temperature of 1g of a substance by 1°C or 1K • Symbol is ‘c’ • The units are J/g°C • A substance with a HIGH specific heat capacity requires a lot of energy to heat it • Eg. water (see Table 1, pg.301)

  5. Calorimetry & Thermal Energy Transfer • Calorimetry– the process of measuring thermal energy changes in a physical/chemical change using a calorimeter. • Any change in the thermal energy of a system is detected as a temperature change in water.

  6. Calorimeters 3 Types: Coffee cup calorimeter: • Constant volume calorimeter (open to the atmosphere) -reasonably accurate measurements of heat transfer • Measure initial and final temp of the solution.

  7. Calorimeters Laboratory calorimeter: • Can provide precise data for changes of state or chemical reactions that do not involve gases. Bomb Calorimeter: • Container inside which a fuel is burned • Much more accurate measurements of heat transfer. • Pressure changes depending on the reaction • Often described as ‘isolated’

  8. Assumptions made in Calorimetry • >  There is no/negligible heat transfer between the environment and the calorimeter. • >  The calorimeter doesn’t absorb energy (or we are saying it is negligible • >  All dilute, aqueous solutions are assumed to have the same density and specific heat capacity equal to water. pwater= 1.00g/mL cwater = 4.18 J/g°C

  9. Calorimetry Calculations • q=heat transferred in a reaction to the surroundings and... if q is + then it is endothermic if q is – then it is exothermic • m=mass (in grams) • c=specific heat capacity (in J/(g °C) see chart in textbook) • ΔT=the change in temperature. (measured in Celsius of Kelvin) ΔT=T2 ‐T1

  10. Calorimetrycont… • Therefore: HEAT LOST by the PROCESS = HEAT GAINED by the WATER OR HEAT GAINED by the PROCESS = HEAT LOST by the WATER

  11. Sample Problem #1 • A student places 125mL of water at 20°C into a calorimeter and adds a sample of silver at 125°C. The final temperature of the water is 22.5°C. Calculate the quantity of thermal energy absorbed by the water.

  12. Sample Problem 2 Many water heaters use the combustion of natural gas (assume methane) to heat the water in the tank. When 150.0 L of water at 10.0°C is heated to 65.0°C, how much heat flows into the water? m = 150L x 1kg/L = 150 kg = 150 000 g c = 4.18 J/g°C T = (65-10) = 55°C q = mcΔT = 150 000g x 4.18 x 55.0 = 34485000 q = 3.45 x 107 J

  13. Enthalpy • What is enthalpy (H)? • Total amount of thermal energy in a substance Molar enthalpy change (ΔHr) [r is reaction type] • Enthalpy change that occurs when 1 mol of a substance undergoes a physical, chemical or nuclear change (kJ/mol) • ****This value is given – in charts in the back of your textbook & pg307 ΔH= nΔHr

  14. Sample Problem #3 • What amount in g of ethylene glycol would vaporize while absorbing 200.0 kJ of heat? • (ΔHvap = 65.6 kJ/mol)

  15. Enthalpy Changeq … meet ΔH… • Enthalpy change(ΔH): energy change that occurs in a system during a reaction (kJ or J) • ΔH– endothermic is +ve; exothermic is -ve

  16. ΔH and q • Under constant pressure, the enthalpy change of a system is equal to the flow of thermal energy in and out of the system

  17. Changes in Kinetic and Potential Energy high potential energy high kinetic energy Energy ΔH q low kinetic energy low potential energy Reaction Progress Heat Transfer and Enthalpy Change It is not possible to measure all of these energies for a system Instead, we study the energy absorbed or released to the surroundings during a change in the system – the change in enthalpy, ΔH This example of exothermic change shows that the change in potential energy of the system (ΔH) equals the change in kinetic energy of the surroundings (q) ΔHsystem = ±lq surroundingsl

  18. Sample Problem #4 • In a calorimetry experiment, 3.50 g of potassium chloride is dissolved in 200.0mL of water at an initial internal temperature of 23.2°C. The final temp is 19.5°C. What is the molar enthalpy of potassium chloride?

  19. Bond Energies • Bond dissociation energy: the energy required to break a chemical bond (kJ/mol)

  20. You can break covalent bonds, if you supply enough energy...that energy is called bond dissociation energy • Table 1 on Page 307 shows several bond energies for single bonds (multiple bonds are on Pg. 308) • They are measured in kJ/mol, for example it takes 413 kJ to break 1 mol of C‐H bonds

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