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Thermochemistry #1

Thermochemistry #1. Heat Capacity and Specific Heat (Heat Transfer). Chapter 17 Thermodynamics 17.1 The Flow of Energy - Heat & Work. Thermochemistry: study of energy changes that occur during chemical reactions & changes of state. Energy. The ability to do work or transfer heat.

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Thermochemistry #1

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  1. Thermochemistry #1 Heat Capacity and Specific Heat (Heat Transfer)

  2. Chapter 17 Thermodynamics17.1 The Flow of Energy - Heat & Work Thermochemistry: study of energy changes that occur during chemical reactions & changes of state

  3. Energy • The ability to do work or transfer heat. • Work: Energy used to cause an object that has mass to move. • Heat: Energy used to cause the temperature of an object to rise.

  4. The Nature of Energy (cont.) • The law of conservation of energy states that in any chemical reaction or physical process, energy can be converted from one form to another, but it is neither created nor destroyed. • the first law of thermodynamics.

  5. The Nature of Energy • Two forms of energy exist, potential and kinetic. • Potential energy is due to composition or position. • Kinetic energy is energy of motion.

  6. Measuring Heat A calorie is defined as the amount of energy required to raise the temperature of one gram of water one degree Celsius. 1 calorie = amount of heat needed to raise 1 g of water 1oC

  7. Food is measured in Calories, or 1000 calories (kilocalorie). • 1 Calorie = 1 kilocalorie = 1000 calories

  8. Heat • Heat is energy that is in the process of flowing from a warmer object to a cooler object. • q is used to symbolize heat.

  9. System and Surroundings • The system includes the molecules we want to study (here, the hydrogen and oxygen molecules). • The surroundings are everything else (here, the cylinder and piston).

  10. Exchange of Heat between System and Surroundings • When heat is absorbed by the system from the surroundings, the process is endothermic. Endothermic: system gains heat (+q)

  11. Exchange of Heat between System and Surroundings • When heat is released by the system to the surroundings, the process is exothermic. Exothermic: system loses heat (-q)

  12. Heat Transfer • qsystem = -qsurroundings • The relationship between the algebraic signs of q in a transfer of heat = opposite but equal!

  13. Endothermic: Heat absorbing reaction • Reactants have less energy than products. • Heat is considered a reactant. • Exothermic: Heat releasing reaction • Products have less energy than reactants. • Heat is considered a product.

  14. Practice 17.1 • All exothermic reactions studied under constant pressure are characterized by q <,=, > 0? • The heat energy involved in the vaporization of a liquid is characterized by q <,=, > 0? • Heat energy added to a system from the surroundings is characterized by q <,=, > 0?

  15. Practice 17.1 • All exothermic reactions studied under constant pressure are characterized by q < 0. Q is negative when heat is leaving the system • The heat energy involved in the vaporization of a liquid is characterized by q > 0. Q is positive because heat is required to get liquid to gas. Endothermic • Heat energy added to a system from the surroundings is characterized by q > 0. Endothermic reaction: heat is added to system.

  16. Thermal Properties • The physical properties of a substance that concern its ability to absorb heat without changing chemically are called its thermal properties. • Three examples are heat capacity, molar heat capacity, specific heat capacity, which is usually just called specific heat.

  17. Heat Transfer: Heat Capacity & Specific Heat • Heat capacity: amount of heat needed to raise a substance 1oC(J/oC). Mass is not a factor. • q= C • ΔT • Specific heat (C): amount of heat needed to raise 1g of a substance 1oC (J/g.oC). Mass is a factor. • C = q/(m • ΔT)

  18. Specific Heat • Some objects require more heat than others to raise their temperature. • High Specific Heat Capacity is when a substance takes a long period of time to heat up, and cool down. • Example: Large bodies of water. They help to moderate the temperature of the land around it.

  19. Specific Heat (cont.) Calculating heat absorbed and released. Dependent on mass. Measured per gram unit • q = c × m × ΔT • q = heat absorbed or released • c = specific heat of substance (sometimes seen as s) • m = mass of substance in grams • ΔT = change in temperature in Celsius

  20. Practice 17.2 The specific heat for water is 4.184 J/g.oC. How many joules of heat energy must you provide 245 grams of water in order to increase its temperature from 25.0oC to 100.0oC?

  21. Practice 17.2 The specific heat for water is 4.184 J/g.oC. How many joules of heat energy must you provide 245 grams of water in order to increase its temperature from 25.0oC to 100.0oC?

  22. Molar Heat Capacity • The molar heat capacity of a substance is the amount of energy required to raise one mole of the substance by one degree. • The standard unit is joules/mol•K. • Specific heat differs from molar heat capacity in that it is measured per gram instead of per mole.

  23. Molar Heat Practice 1 How much heat is lost when 36.0 g platinum (molar heat capacity for platinum = 25.64 J/mol•oC) at 35.0oC is cooled to 27.4oC?

  24. Molar Heat Practice 1 How much heat is lost when 36.0 g platinum (molar heat capacity for platinum = 25.64 J/mol•oC) at 35.0oC is cooled to 27.4oC?

  25. Heat Capacity • When heat is transferred to an object, the temperature of the object increases. • When heat is removed from an object, the temperature of the object decreases. • The relationship between the heat ( q ) that is transferred and the change in temperature ( ΔT ) is • q = C ΔT = C ( Tf - Ti) • Unit = J oC-1 or J K-1

  26. Calorimeter • A calorimeter is a device in which a chemical reaction or physical process takes place. • The calorimeter is well-insulated so that, ideally, no heat enters or leaves the calorimeter from the surroundings. • Any heat liberated by the reaction or process is picked up by the calorimeter and other substances in the calorimeter.

  27. Constant Pressure Calorimetry By carrying out a reaction in aqueous solution in a simple calorimeter, one can indirectly measure the heat change for the system by measuring the heat change for the water in the calorimeter.

  28. Constant Pressure Calorimetry Because the specific heat for water is well known (4.184 J/mol-K), we can measure the heat change of the reaction with this equation: q = m  c  T

  29. qsystem = -qsurroundings • In chemistry the thermodynamic sign convention is things entering the system a positive in value, and things leaving the system are negative in value.

  30. Practice 17.4 • Central processing chips in computers generate a tremendous amount of heat – enough to damage themselves if not cooled. Aluminum “heat sinks” are often attached to the chips to carry away excess heat. • Suppose that a heat sink at 71.3 °C is dropped into a styrofoam cup containing 100.0 g of water at 25.0 °C. The temp of the water raises to 27.4 °C. What is the heat capacity of the sink?

  31. To solve: • Heat capacity is related to temperature change: C=q/Δt • q(heat sink) = -q(H2O) • In this problem: As the heat sink loses water we expect it to be +, as the water gains heat from sink we expect it be - • To calculate q(H2O) we use the specific heat of water. Once we find q(H2O) we change the sign to get q(heat sink)

  32. Practice 17.5 A piece of metal weighing 59.047 g was heated to 100.0 °C and then put it into 100.0 mL of water (initially at 23.7 °C). The metal and water were allowed to come to an equilibrium temperature, determined to be 27.8 °C. Assuming no heat lost to the environment, calculate the specific heat of the metal.

  33. -qmetal = qwater -[(mass) (Δt) (Cp)] = (mass) (Δt) (Cp) -[(59.04 g)(-72.2°C)(x)] = (100.0g)(4.1°C)(4.184 J/g °C) x = 0.40 J/g°C

  34. Molar Heat Practice 2 • Compound A is burned in a bomb calorimeter that contains 2.50 liters of water. If the combustion of 0.175 moles of this compound causes the temperature of the water to rise 45.00 °C, what is the molar heat capacity of combustion of compound A? • The heat capacity of water is 4.184 J/g°C

  35. q = mCΔT q = (2.50 x 103 g H2O)(4.184 J/g°C)(45.00 °C) q = 471,000 J Because 471 kJ of energy are given of when 0.175 moles of compound A burn.

  36. Practice 17.6  Suppose a piece of iron with a mass of 21.5 g at a temp of 100.0 °C is dropped into an insulated container of water. The mass of the water is 132.0 g and its temperature before adding the iron is 20.0 °C. What will be the final temp of the system? Specific heat of iron is 0.449 kJ/kg K.

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