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Energy

Energy. Chapter 16. 16.1 Energy. Energy is the ability to do work or produce heat. Heat is commonly measured in joules or calories. 1 calorie = 4.184 joules. Specific Heat.

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Energy

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  1. Energy Chapter 16

  2. 16.1 Energy • Energy is the ability to do work or produce heat. • Heat is commonly measured in joules or calories. • 1 calorie = 4.184 joules

  3. Specific Heat • The amount of heat produced required to raise the temperature of one gram of a substance by one degree Celsius is known as the specific heat. • Water has a high specific heat of 4.184 J / (g °C). • The equation to calculate heat is q = c x m x ΔT

  4. Example • A silver bar with a mass of 250.0 g is heated from 22.0 ° C to 68.5 ° C. How much heat does the silver bar absorb? Use Table-16-2.

  5. Practice • How much heat does a 23.0 g ice cube absorb as its temperature increases from -17.4 ° C to 0.0 ° C? Give the answer in both joules and calories. • A sample of an unknown metal has a mass of 120.7 g. As the sample cools from 90.5 ° C to 25.7 ° C, it releases 7020 J of energy. What is the specific heat of the sample? Identify the metal among those in Table 16-2 in your textbook.

  6. Practice • A 15.6 g sample of ethanol absorbs 868 J as it is heated. If the initial temperature of the ethanol is 21.5 ° C, what is the final temperature of the ethanol?

  7. Section 16.2 Heat in Chemical Reactions and Processes

  8. Calorimeter • The device used to measure heat changes is an insulated one called a calorimeter. • In the calorimeter, the temperature change of a known mass of water is used to determine the amount of energy gained or released. • Remember, the heat lost by one material is gained by the other.

  9. Example • A calorimeter contains 195 g of water at 20.4 ° C. A 37.8 g sample of an unknown metal is heated to 133 ° C and placed into the water in the calorimeter. Heat flows from the metal to the water until both reach a final temperature of 24.6 ° C. What is the specific heat of the metal?

  10. Practice • A 50.6 g sample of iron metal is heated and put into 104 g of water at 19.7 ° C in a calorimeter. If the final temperature of the iron sample and the water is 24.3 ° C, what was the temperature of the iron sample when it was placed in the water? • A 77.5 g sample of an unknown solid is heated to 62.5 ° C and placed into a calorimeter containing 93 g of water at 23.3 ° C. If the final temperature of the solid sample and the water is 26.2 ° C, what is the specific heat of the solid?

  11. Thermochemistry • Thermochemistry is the study of heat changes that accompany chemical reactions and phase changes. • The system is the reaction or process being studied and everything outside the system is called the surroundings. • The universe is defined as the system plus the surroundings.

  12. Enthalpy • The heat content of a system at constant is called the enthalpy (H) of the system. • The heat absorbed or released during a change in a system at constant pressure is the change in enthalpy (ΔH). • The enthalpy change for a rxn is called the enthalpy of reaction or heat of reaction. ΔHrxn = Hproducts - Hreactants

  13. Energy in Reactions • In an endothermic rxn, heat is absorbed and the ΔHrxn is positive. • In an exothermic rxn, heat is released and ΔHrxn is negative.

  14. 16.3 Thermochemical Equations • A thermochemical equation is a balanced chemical equation that includes the physical states of the reactants and products and the change in enthalpy. • For example, the combustion of ethanol is: C2H5OH (l) + 3O2 (g)  2CO2 (g) + 3H2O (l) ΔHcomb = -1367 kJ

  15. Enthalpy Change for Rxn The enthalpy change for the complete burning of one mole of a substance is the enthalpy (heat) of combustion (ΔHcomb) for that substance. Heat is also absorbed when materials change state. The heat required to vaporize one mole of a liquid is called molar enthalpy (heat) of vaporization (Δ Hvap). The heat required to melt one mole of a solid is its molar enthalpy (heat) of fusion (ΔHfus).

  16. Example The enthalpy of combustion for methanol (CH3OH) is -726 kJ/mol. How much heat is released when 82.1 g of methanol is burned?

  17. Practice Calculate the heat required for the following two processes and compare the results. A 100.0 g sample of solid ethanol melts at its melting point. ΔHfus = 4.94 kJ/mol A 100.0 g sample of liquid ethanol vaporizes at its boiling point. ΔHvap = 38.6 kJ/mol How much heat is evolved when 24.9 g of propanol (C3H7OH) is burned? ΔHcomb = -2010 kJ/mol

  18. 16.4 Calculating Enthalpy Change The ΔH for a chemical reaction can be determined by Hess’s law. This equation states that two or more thermochemical equations can be added to produce a final reaction, and the enthalpy change for the final reaction equals the sum of the enthalpy changes for the individual rxn.

  19. Example Use thermochemical equations a and b to determine ΔH for the oxidation of ethanol (C2H5OH) to form acetaldehyde (C2H4O) and water. 2C2H5OH (l) + O2 (g)  2C2H4O (g) + 2H2O (l) a. 2C2H4O (g) + 5O2 (g)  4CO2 (g) + 4H2O (l) ΔH = -2385 kJ b. C2H5OH (l) + 3O2 (g)  2CO2(g) + 3H2O (l) ΔH = -1367 kJ

  20. Practice 10. Use reactions a and b to determine ΔH for this single-displacement rxn. Cl2 (g) + 2HBr (g)  2HCl (g) + Br2 (g) a. H2 (g) + Cl2 (g) 2 HCl (g) ΔH = -185 kJ b. H2 (g) + Br2 (g)  2HBr (g) ΔH= -73 kJ

  21. Practice • Use rxns a, b and c to determine ΔH for the rxn of carbon monoxide and hydrogen to form methanol (CH3OH). CO (g) + 2H2 (g)  CH3OH (l) • 2CO (g) + O2 (g)  2CO2 (g) ΔH = -566 kJ • 2H2 (g) + O2 (g)  2H2O (l) ΔH = -572 kJ • 2CH3OH (l) + 3O2 (g) 2CO2 (g) + 4H2O (l) ΔH = -1452 kJ (Hint, First find ΔH for the rxn 2CO (g) + 4H2 (g)  2CH3OH (l), then divide this result by 2 to obtain your final answer.)

  22. Standard Heat of Formation • The standard state of a substance is the normal state of the substance at 298 K and 1 atm. • The change in enthalpy that accompanies the formation of one mole of a compound in its standard state from its constituent elements in their standard sates is called the standard enthalpy (heat) of formation (ΔH°f).

  23. 16.5 Reaction Spontaneity • Entropy (S) is a measure of the disorder or randomness of the particles that make up a system. • Spontaneous processes always result in an increase in the entropy of the universe. • The change in entropy of a system is given by the equation ΔSsystem = Sproducts - Sreactants

  24. Predicting S Values • Whether Δssystem is positive or negative can be predicted in some cases by examining the rxn or process. • There are several factors that affect the change in entropy of a system. • Changes in state • Dissolving of a gas in a solvent • Change in the number of gaseous particles • Dissolving of a solid or liquid to form a solution • Change in temperature

  25. Changes of State • Entropy increases when a solid changes to a liquid and when a liquid changes to a gas because the particles are allowed to move more freely.

  26. Dissolving of a Gas in a Solvent • When a gas is dissolved in a liquid or solid solvent, the motion and randomness of the particles are limited and the entropy of the gas decreases. Entropy Goes Down

  27. Change in the Number of Gaseous Particles • When the number of gaseous particles increases, the entropy of system usually increases because more random arrangements are possible.

  28. Dissolving of a Solid or Liquid to Form a Solution • When solute particles become dispersed in a solvent, the disorder of the particles and the entropy of the system usually increases.

  29. Change in Temperature • A temperature increase results in increased disorder of the particles and an increase in entropy.

  30. Free Energy • For a rxn or process occurring at constant temperature and pressure, the energy that is available to do work is the free energy (G). • Free energy is related to enthalpy and entropy by the following equation: ΔGsystem = Δhsystem-TΔSsystem

  31. Free Energy Equation • In the equation T is the Kelvin temperature. • If ΔGsystem is negative, the reaction or process is spontaneous. • If ΔGsystem is positive, the reaction or process is nonspontaneous.

  32. Example • For a chemical rxn, ΔHsystem = -81 kJ and ΔSsystem = -215 J/K. Is the rxn spontaneous?

  33. Practice • Calculate ΔGsystem for each process, and state if the process is spontaneous or nonspontaneous. • ΔHsystem = 147 kJ, T = 422 K, ΔSsystem = -67 J/K • ΔHsystem = -43 kJ, T = 21°C, ΔSsystem = -118 J/K • ΔHsystem = 227 kJ, T = 574 K, ΔSsystem = 349 J/K

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