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Exploring Energy: Laws, Measurement, and Practical Examples in Thermodynamics

This chapter delves into the fundamental concepts of energy, temperature, and heat, examining the measurement of energy changes through specific heat capacity, the First and Second Laws of Thermodynamics, and enthalpy changes. Practical examples illustrate the application of these concepts, such as calculating energy requirements for heating metals and water, and analyzing combustion reactions of fuels for modern use. The interplay of energy transformations in chemical reactions is also explored through Hess’s Law and entropy calculations, providing a comprehensive understanding of thermodynamic principles.

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Exploring Energy: Laws, Measurement, and Practical Examples in Thermodynamics

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  1. Chapter 10 – Energy • What is Energy? • Temperature & Heat • Measuring Energy Changes – Specific Heat Capacity • First Law of Thermodynamics • Enthalpy • Hess’ Law • Second Law of Thermodynamics – Entropy • Fuels for the Modern World

  2. Example 1 A brownie contains 320 Calories. Express this in calories and in Joules.

  3. Example 2 How much energy is required to heat 10.0 g of iron from 25.0°C to 100.0°C? The specific heat capacity of iron is 0.45 J/g°C.

  4. Example 3 A 19.6 g sample of an unknown metal was heated from 22.0°C to 53.8°C by the addition of 558.6 J of energy. Find the specific heat capacity of the metal. Refer to the table of specific heat capacities in your textbook to predict what metal this might be.

  5. Example 4 How much energy (in joules) is needed to heat a cup of water for coffee?  Assume the following: the cup holds 250 mL of water the density of water is 1.00 g/mL the water is initially at 19°C the ideal temperature for brewing coffee is 97°C Answer: 81.6 x 103 J

  6. Example 5 A system releases 255 calories of heat and does 428 calories of work. What is the change in internal energy of this system? Answer: -683 calories

  7. Example 6 The enthalpy of the reaction CH4 (g) + 2O2 (g) CO2 (g) + 2 H2O (g)is -890.4 kJ/mol. What is the enthalpy for the reaction below? 2 CO2 (g) + 4 H2O (g) 2 CH4 (g) + 4 O2 (g)

  8. Example 7 How much heat is released when 2.50 g of N2 reacts with excess H2 to form NH3 (g)? N2 (g)  +  3 H2 (g)    2 NH3 (g) DH° = -91.8 kJ per mole of N2 consumed. Answer: 8.20 kJ are released

  9. Example 8 Methane, CH4, is natural gas, a commonly used fuel. How much methane must be burned to release 100.0 kJ of energy? Assume the enthalpy change for the combustion of 1 mole of methane is -890.4 kJ.

  10. Example 9 Which provides more energy, burning 1.00 g of gasoline (C8H18) or 1.00 g of ethanol (C2H5OH)? C8H18(l) + 25/2 O2(g)  8 CO2(g) + 9 H2O(g) DH = -5468kJ per mole of gasoline burned C2H5OH(l) + 3 O2(g)  2 CO2(g) + 3 H2O(g) DH = -1367kJ per mole of ethanol burned

  11. Example 10 Given the following reactions: 2SO2 (g) + O2 (g)  2SO3 (g); DH1 = -196kJ 2S (s) + 3O2 (g)  2SO3 (g); DH2 = -790 kJ Calculate the heat of reaction for S (s) + O2 (g)  SO2 (g)

  12. Example 11 Given the following reactions: H2S (g) + 3/2 O2(g)  H2O(l) + SO2(g); DH1 = -563 kJ CS2 (l) + 3O2(g)  CO2(g) + 2SO2(g); DH2 = -1075 kJ Calculate DHrxn for this reaction: CS2(l) + 2 H2O(l)  2 H2S(g) + CO2(g)

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