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Ch.16 Reaction Energy

Ch.16 Reaction Energy. Mrs. Geisler Chem. 2 2014. Sec. thermochemistry. Study of the transfers of energy as heat that accompany chemical reactions and physical changes. Calorimeter – energy absorbed or released as heat in a chemical or physical change is measured in this apparatus.

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Ch.16 Reaction Energy

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  1. Ch.16 Reaction Energy Mrs. Geisler Chem. 2 2014

  2. Sec. thermochemistry • Study of the transfers of energy as heat that accompany chemical reactions and physical changes. • Calorimeter – energy absorbed or released as heat in a chemical or physical change is measured in this apparatus.

  3. Temperature – measure of the average kinetic energy of the particles in a sample of matter. • The greater the KE the higher the temperature is. • K = 273.15 + C • Joule (J) – SI unit of heat (and other forms of energy) J = Nxm = kgxm2/s2

  4. Heat – energy transferred between samples of matter because of a difference in their temperatures. • Ex. Feeling the metal part of the desk, feels cold because metal is a good conductor and heat is leaving your hand.

  5. Specific heat – amount of energy required to raise the temperature of one gram of a substance by one Celsius degree or one kelvin. Symbol = cp **make sure your units equal!! Table 1 pg. 533 (shows specific heat of common substances) Water (l) = 4.18 J/gK Water(s) = 2.06 J/gK

  6. Specific heat equation Cp = q m x ΔT q = heat m = mass T = temperature (T final – T initial)

  7. Enthalpy of Reaction • ΔH = energy absorbed as heat during a chemical reaction at constant pressure. • H = quantity called enthalpy • Enthalpy change = amount of energy absorbed by a system as heat during a process at constant pressure. ΔH = Hproducts - Hreactants

  8. Enthalpy of reaction = quantity of energy transferred as heat during a chemical reaction. Difference between the stored energy of the reactants and the products. “heat of reaction”

  9. Example Hydrogen and oxygen ignited, water will form and energy will be released explosively. Energy released comes from the reactants as they form products. Exothermic reaction = energy is released. Endothermic reaction = energy is absorbed. Energy contained in product must be less than the reactants. **coefficients represent moles in enthalpy reactions. 2H2(g) + O2(g)  2H2O(g) + 483.6 kJ thermochemical equation

  10. Thermochemical equation – an equation that includes the quantity of energy released or absorbed as heat during the reaction as written. 2H2(g) + O2(g)  2H2O(g) + 483.6 kJ H2(g) + ½ O2(g)  H2O(g) + 241.8 kJ **producing half as much water would require one-half as many moles of reactants and would release only one-half as much energy.

  11. Endothermic reaction Reverse 2H2O(g) + 483.6 kJ  2H2(g) + O2(g) Thermochemical equation written by designating the value of ΔH. Positive when endothermic, negative when exothermic. 2H2(g) + O2(g)  2H2O(g) ΔH = -483.6 kJ 2H2O(g)  2H2(g) + O2(g) ΔH = +483.6 kJ

  12. Keep in mind the following when using thermochemical equations: • Coefficients represents moles. • Physical state is important to note. • Change in enthalpy is directly proportional to the number of moles of the substances. • Value of the enthalpy change is usually not significantly influenced by changing temperature.

  13. Enthalpy of Formation Molar enthalpy of formation = enthalpy change that occurs when one mole of a compound is formed from its elements in their standard state at 25°C and 1 atm. Standard state is represented as 0, that is added to the enthalpy symbol ΔH 0 ΔH 0f = f stands for formation

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