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Thermochemistry Some Like It Hot!!!!!
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The Flow of Energy ► Thermochemistry – concerned with heat changes that occur during chemical reactions ► Energy - capacity for doing work or supplying heat
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Is there a difference between heat and temperature? ► Temperature – Average kinetic energy of the particles in matter ► Heat - represented by “q”, is energy that transfers from one object to another, because of a temperature difference between them. only changes can be detected! flows from warmer cooler object
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Exothermic and Endothermic Processes ► Essentially all chemical reactions and changes in physical state involve either: release of heat, or absorption of heat ► In studying heat changes, think of defining these two parts: the system - the part of the universe on which you focus your attention the surroundings - includes everything else in the universe ► Law of Conservation of Energy – Energy is neither created nor destroyed
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Exothermic vs Endothermic ► Endothermic – Heat flowing into a system from its surroundings Surroundings get cooler but the system gets warmer q has a positive value ► Exothermic - Heat flowing out of a system into it’s surroundings System gets colder but the surroundings get warmer q has a negative value
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► Endothermic ► Exothermic System Heat System increases in temperature, Surroundings decrease in temperature System decreases in temperature, surroundings increase in temperature
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Endothermic and Exothermic Reaction Review ► Every reaction has an energy change associated with it ► Exothermic reactions release energy, usually in the form of heat. ► Endothermic reactions absorb energy ► Energy is stored in bonds between atoms
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Measuring Heat Flow ► Measured with 2 common units Calorie – quantity of heat needed to raise the temperature of 1 g of pure water 1 o C. Joule - SI unit of heat and energy ► How are the units similar? 1 J = 0.2390 calories 1 Calorie = 4.184 J
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Heat Capacity and Specific Heat ► Heat Capacity – the amount of heat needed to increase the temperature of an object exactly 1°C Dependant on 2 things: ► Mass of the object ► Chemical Composition ► Specific Heat – The amount of heat it takes to raise the temperature of 1 g of the substance 1°C Written as J/(g x °C)
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The higher the specific heat the more energy needed to change the temperature Note the tremendous difference in Specific Heat. Water’s value is VERY HIGH.
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Calculating the Quantity of Heat Energy ► To calculate, use the formula: q = mass (in grams) x T x C ► heat is abbreviated as “q” ► T = change in temperature T = Final Temp. – Initial Temp ► C = Specific Heat Units are either: J/(g o C) or cal/(g o C)
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Measuring and Expressing Enthalpy Changes ► Enthalpy (H) - The heat content of a system at constant pressure ► Calorimetry - the measurement of the heat into or out of a system for chemical and physical processes. Based on the fact that the heat released = the heat absorbed ► The device used to measure the absorption or release of heat in chemical or physical processes is called a “Calorimeter”
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► For systems at constant pressure, the “heat content” is the same as a property called Enthalpy (H) of the system A foam cup calorimeter – here, two cups are nestled together for better insulation
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Calorimetry ► Changes in enthalpy = H ► q = H These terms will be used interchangeably in this textbook ► Thus, q = H = m x C x T 4 H is negative for an exothermic reaction 4 H is positive for an endothermic reaction
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Thermochemical Equations ► In a chemical equation, the heat change for the reaction can be written as either a reactant or a product
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C + O 2 → CO 2 Energy ReactantsProducts C + O 2 CO 2 395kJ given off + 395 kJ
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Exothermic ► The products are lower in energy than the reactants ► Thus, energy is released. ► ΔH = -395 kJ The negative sign does not mean negative energy, but instead that energy is lost.
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CaCO 3 → CaO + CO 2 Energy ReactantsProducts CaCO 3 CaO + CO 2 176 kJ absorbed CaCO 3 + 176 kJ → CaO + CO 2
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Endothermic ► The products are higher in energy than the reactants ► Thus, energy is absorbed. ► ΔH = +176 kJ The positive sign means energy is absorbed
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Heat in Change of State ► Heat of Fusion What happens if you place an ice cube on a table in warm room? ► Heat from the room travels to the cooler ice cube causing the ice to melt ► The gain of heat causes a change in state instead of a change in temperature ► The heat required to change the water from a solid to a liquid is called the molar heat of fusion
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Heat of Fusion Molar Heat of Fusion ( H fus. ) = the heat absorbed by one mole of a substance in melting from a solid to a liquid q = mol x H fus. (no temperature change) Values given in Table 17.3, page 522 Example: How much heat is required to melt 5 moles of ice? q = mol x H fus. q = 5 moles H 2 O (s) x 6 kJ/mole q = 30 kJ
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Another Example How many grams of ice at 0°C will melt if 2.25 kJ of heat are added? q = mol x H fus. 2.25 kJ x (1 mol ice/6.01 kJ) x (18 g Ice/1 mol ice) 6.7 g of Ice
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Heat in Change of State ► Molar Heat of Vaporization Similar to Molar Heat of Fusion, however the state of matter change is from a liquid to a vapor The amount of heat necessary to vaporize one mole of a given liquid
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Heat of Vaporization Molar Heat of Vaporization ( H vap. ) = the amount of heat necessary to vaporize one mole of a given liquid. q = mol x H vap. (no temperature change) ► Table 17.3, page 522 ► How much heat is transferred when 5 moles of mercury (Hg) is vaporized? q = mol x H vap. q = 5 moles X 59.1 kJ /mol q = 295.5 kJ
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Another Example ► Example: How much heat (in kJ) is absorbed when 24.8 g H 2 O(l) at 100°C and 101.3 kPa is converted to steam at 100°C? 24.8 g H 2 O x (1 mol H 2 O/18 g H 2 O) x (40.7 kJ/1mol H 2 O) ΔH = 56.1 kJ
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The solid temperature is rising from -20 to 0 o C (use q = mol x ΔT x C) The solid is melting at 0 o C; no temperature change (use q = mol x ΔH fus. ) The liquid temperature is rising from 0 to 100 o C (use q = mol x ΔT x C) The liquid is boiling at 100 o C; no temperature change (use q = mol x ΔH vap. ) The gas temperature is rising from 100 to 120 o C (use q = mol x ΔT x C) The Heat Curve for Water, going from -20 to 120 o C, similar to the picture on page 523 120
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How to determine ΔH Fus. and ΔH Vap ► This is for water only: ΔH Fus. and ΔH Vap. are in kJ/mol ΔH Fus. = (.334 kJ/g) x ( 18g H 2 0/mole) = 6.01 kJ/mol 334 J/g can be found in your reference tables for the heat of fusion of water. You must convert it to.334 kJ because ΔH Fus. is in kJ/mol. ΔH Vap = (2.26 kJ/g) x (18 g H 2 0/mol) = 40.7 kJ/mol 2260 J/g can be found in your reference tables for the heat of vaporization of water. You must convert it to 2.26 kJ/g because ΔH Fus. is in kJ/mol.
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Entropy ► Defined as a measure of the disorder of a system Systems tend to go from a state of order (low entropy)to a state of maximum disorder (high entropy) Increasing Entropy Solids Liquid Gas Ionic Compounds Ions in solution
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- Page 570 Entropy of the gas is greater than the solid or liquid Entropy is increased when a substance is divided into parts Entropy increases when there are more product molecules than reactant molecules Entropy increases when temperature increases
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