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Thermochemistry

Thermochemistry. Thermochemistry is the study of heat changes that occur during chemical reactions. Heat (q) - energy that is transferred from one object to another because of a temperature difference between them. Heat always flows from a warmer object to a cooler object.

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Thermochemistry

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  1. Thermochemistry

  2. Thermochemistry is the study of heat changes that occur during chemical reactions. • Heat (q) - energy that is transferred from one object to another because of a temperature difference between them. • Heat always flows from a warmer object to a cooler object. • Heat cannot be measured directly. • Temperature – the average kinetic energy of the particles in a sample of matter • Energy - the capacity for doing work or supplying heat.

  3. Energy in Reactions • Endothermic process - system absorbs heat from the surroundings. (Energy must be added to the reaction for the reaction to take place) • Exothermic process - system releases heat to the surroundings.

  4. Ex. Endothermic or exothermic process? -evaporating alcohol -leaves burning -boiling water -water cooling -melting ice -freezing water

  5. To calculate the heat energy requiredfor a temperature change, use the following formula: Q = (m)(C)(DT) Q = heat energy absorbed or released (J) • + q = endothermic (heat absorbed) • - q = exothermic (heat released) m = mass (g)C = specific heat of the object (J/goC) • The amount of heat it takes to raise the temperature of 1 g of a substance 1oC. • Water has a uniquely high specific heat compared to other substances. DT= change in temperature of object (oC)

  6. Example • How much energy is required to heat an iron nail with a mass of 7.0 g from 25oC until it becomes red hot at 752oC? C Fe= 0.46J/goC T = 752-25 = 727oC q = mC T q = 7.0g(0.46J/goC) 727oC q = 2300 J

  7. Phase Change • Heat energy can be used to not only change the temperature of matter, but also its phase. • The energy goes into separating or organizing the molecules into a new state • The amount of heat energy necessary to cause a phase change can be calculated using the formula: Q = mLf (solid/liquid) or Q = mLv (liquid/gas)

  8. Q = heat required for phase change (J) • m = mass (kg) • Lf = latent heat of fusion (J/kg) • Energy required to change one kilogram of a material from a solid to a liquid • Water Lf = 334,000 (J/kg) • Lv = latent heat of vaporization (J/kg) • Energy required to change one kilogram of a material from a liquid to a gas • Water Lv = 2,260,000 (J/kg)

  9. Example • Find the amount of heat needed for melting 1.3 kg ice at 0 ºC?

  10. Notice you calculate each change separately and then add them all up. Q = mCT Q = m Lv Q = mCT Q = m Lf Q = mCT

  11. Enthalpy ( H) • The amount of heat energy absorbed or lost by a system • Exothermic reactions have -  H • Endothermic reactions have +  H •  H is also called the heat of reaction.

  12. Calculating Heat Changes • Bond energy is the amount of energy required to break or make a bond • Breaking a bond is endothermic (+ H) • Forming a bond is exothermic (-  H) •  H = Hproducts – Hreactants

  13. Calculate the energy of the reaction using heats of formation. H2(g) + Br2(g)  2 HBr(g)  H HBr = -36.29 kJ/mol All elements and diatomic molecules have an enthalpy of zero. H = Hproducts – Hreactants  H = [2 x -36.29 kJ/mol] – [0] H = -72.58 kJ/mol Exothermic Example 1

  14. Example 2 • Calculate the energy of the reaction using heats of formation. CH4(g) + 2 O2(g)  CO2(g) + 2 H2O(l) •  H CH4 = -74.9 kJ/mol •  H CO2 = -393.5 kJ/mol •  H H2O = -285.8 kJ/mol •  H = Hproducts – Hreactants •  H = [-393.5 kJ/mol + (2 x -285.8 kJ/mol)] – [-74.9 kJ/mol+ (2 x 0)] •  H = -890.2 kJ/mol • Exothermic

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