Exothermic and Endothermic Reactions Modeling Exothermic and Endothermic Reactions (1) Qualitative Models Endothermic Reaction: (Endo: “In”); Energy is absorbed Note: Energy of the Product is greater than the Energy of the Reactant Energy Exothermic Reaction : (Exo: “Out”); Energy is released Note: Energy of the Reactant is greater than the Energy of the Product Energy You Do Qualitative Models: Your Turn Choose a design of your choice to illustrate the Endothermic and Exothermic Reactions Endothermic Reaction: (Endo: “In”); Energy is absorbed Note: Energy of the Product is greater than the Energy of the Reactant Exothermic Reaction : (Exo: “Out”); Energy is released Note: Energy of the Reactant is greater than the Energy of the Product (2) Quantitative Models – 1: Guided Practice Let each arrow be a bond and the numeric value on it indicates its Bond Energy in kJ/mol) Endothermic Reaction: Energy is Absorbed 12 Energy 18 18 12 12 20 12 12 12 Exothermic Reaction: Energy is Released Energy 12 12 You Do (2) Quantitative Models -1: Your Turn Choose a design of your choice to illustrate the Endothermic and Exothermic Reactions Enthalpy SC2g Develop a model to illustrate the release or absorption of energy (endothermic or exothermic) from a chemical reaction system depends upon the changes in total bond energy. SC5a Plan and carry out an investigation to calculate the amount of heat absorbed or released by chemical or physical processes. SC5b Construct an explanation using a heating curve as evidence of the effects of energy and intermolecular forces on phase changes. 1. Enthalpy is the scientific term for Heat of Reactions. 2. We cannot determine the absolute value of Enthalpy. 3. We can only determine the changes in Enthalpy during a chemical reaction. The Change in Enthalpy is given by the notation, ∆H. Enthalpy change is determined by the formula: Enthalpy Change = Standard Heat of Formation of Products – Standard Heat of Formation of Reactants Heat of formation of various compounds are available on the Internet. The enthalpy of formation for an element in its elemental state will always be 0 because it takes no energy to form a naturally-occurring compound. Heat of formation is a net sum of the bond energies present in a compound determined by thermochemical methods (bomb calorimeter.) So, scientists determine bond energies and then calculate the heat of formation of compounds. So, heat of formation is based strictly on bond energies. The change in enthalpy can be positive or negative depending on whether the reaction is endothermic or exothermic. In Exothermic Reactions, the Products have less energy or heat content than the Reactants, as energy is released. In Endothermic Reactions, Products have more energy or heat content than the Reactants. In chemistry, IUPAC changed the definition of standard temperature and pressure (STP) in 1982:[1] Until 1982, STP was defined as a temperature of 273.15 K (0 °C, 32 °F) and an absolute pressure of exactly 1 atm (101.325 kPa). Since 1982, STP is defined as a temperature of 273.15 K (0 °C, 32 °F) and an absolute pressure of exactly 105 Pa (100 kPa, 1 bar). 5. Enthalpy is a State Function not a path function. This means that regardless of the pathway for the chemical reaction, so long as the state of the chemical prepared or manufactured is the same, the Enthalpy Change will be the same. Here, “State” refers to the State of Matter: Solid, Liquid, Gas, Plasma, and Bode-Einstein Condensate Some Standard Heat of Formation of Values M H ø Compound Formula f g mol-1 kJ mol-1 Carbon monoxide CO(g) 28.0 -110.5 Carbon dioxide CO2(g) 44.0 -393.5 Methane CH4(g) 16.0 -74.8 Ethane CH3CH3(g) 30.1 -84.7 Propane CH3CH2CH3(g) 44.1 -104.5 Butane CH3(CH2)2CH3(g) 58.1 -126.5 The enthalpy of Pentane CH3(CH2)3CH3(l) 72.2 -173.2 formation for Hexane CH3(CH2)4CH3(l) 86.2 -198.6 an element in its Heptane CH3(CH2)5CH3(l) 100.2 -224.0 state of natural Octane CH3(CH2)6CH3(l) 114.2 -250.0 occurrence will Nonane CH (CH ) CH 128.3 -274.0 3 2 7 3(l) always be 0 because Decane CH (CH ) CH 142.3 -300.9 3 2 8 3(l) it takes no energy to Undecane CH3(CH2)9CH3(l) 156.3 -327.2 Dodecane CH (CH ) CH 170.3 -350.9 form a naturally- 3 2 12 3(l) occurring element. 2-Methylpropane (CH3)2CHCH3(g) 58.1 -134.5 2-Methylbutane (CH3)2CHCH2CH3(l) 72.2 -178.9 2-Methylpentane (CH3)2CH(CH2)3CH3(l) 86.2 -204.6 2-Methylhexane (CH3)2CH(CH2)4CH3(l) 100.2 -229.5 2-Methylheptane (CH3)2CH(CH2)5CH3(l) 114.2 -255.0 2,2-Dimethylpropane C(CH3)4(g) 72.2 -189.8 Mathematical Models of Endothermic and Exothermic Reactions: Guided Practice (1) How is heat of formation calculated Given Bond energy of H-H bond is 432 kJ/mol Bond Energy of Cl-Cl bond is 239 kJ/mol Calculate the heat of formation of Hydrogen chloride The balanced chemical equation is H2(g) + Cl2(g) → 2HCl(g) Mathematical Models of Endothermic and Exothermic Reactions Collaborative Practice -(-1260) -2361.0 -1714.8 +1260 -241 -2815.8 kJ/mol Independent Practice Independent Practice Determine the Enthalpy for the following reaction and label it as exothermic and endothermic. Determine the Enthalpy for the following reactions and label them each as exothermic and endothermic. (Refer to the chart for heat of formation values) C2H2(g) +2H2(g) → C2H6(g) Heat of formation of ethane is –84.7 kJ mol-1 Heat of formation of acetylene is +228.3 kJ mol-1 -99 CO(g) Independent Practice Standard heat of formation of OF2(g) is 24.7 kJ/mol Standard heat of formation of SF4 -763.16 kJ/mol Standard heat of formation of SO2 -296.84 kJ/mol Standard heat of formation of hydrogen fluoride -321.09 kJ/mol Standard heat of formation of liquid water -285.8 kJ/mol Heat of formation of sulfur gas 9.8 kJ mol-1 Heat of formation of sulfur dioxide gas -296.8 kJ/mol From the following data, prove that Enthalpy is a State Function .