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Energy Changes in Chemical Reactions

Energy Changes in Chemical Reactions. Thermochemistry. Thermodynamics and Thermochemistry. Thermodynamics = the study of heat and its transformations. Thermochemistry = the part of thermodynamics that deals with changes in heat that take place during chemical reactions as in most

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Energy Changes in Chemical Reactions

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  1. Energy Changes in Chemical Reactions Thermochemistry

  2. Thermodynamics and Thermochemistry Thermodynamics = the study of heat and its transformations. Thermochemistry = the part of thermodynamics that deals with changes in heat that take place during chemical reactions as in most chemical reactions, energy is absorbed or released.

  3. What is Internal Energy? Internal Energy (E)- the combined kinetic and potential energy of all particles in a system. The total internal energy of a system is not usually known. Energy changes (DE) can be measured.

  4. Examples of Potential and Kinetic that Contribute to Internal Energy • Potential Energy- • Intermolecular forces of attraction or repulsion • Intramolecular forces of attraction or repulsion • Kinetic Energy • Rotational nature of molecules • Vibrational nature of molecules interactions between particles of matter down to the subatomic level (p, n, and e-) but we are generally concerned about interactions at the molecular level.

  5. System What is the system? The object or substance undergoing a change of physical state or reaction. What we are studying. Surroundings are everything around the system that can exchange energy with the system.

  6. Types of Systems Open system can exchange mass and energy (usually in the form of heat) with the surroundings. Closed system can exchange energy but not mass. Isolated system – no transfer of energy or mass.

  7. State Functions State functions are path independent. They depend only on present state and are independent of history of the system. The change in a state function does not depend on how the process is carried out.

  8. State Functions

  9. Total Internal Energy (E) in a System Etotal = mgh + ½ mv2 OR Potential Energy + Kinetic Energy Internal energy of the system = sum of kinetic and potential energies making up a substance down to the subatomic level. We often don’t know the total energy of a system. Chemists aren’t so interested in total energy in a system- they’re more interested in the energy changes that happen during chemical reactions.

  10. First Law of Thermodynamics The total energy in the universe is constant (Law of Conservation of Energy). “You can’t get something for nothing” There are essentially two ways to change the energy of a system- heat(q) and work(w). DE = q + w

  11. Energy Change DE Energy change in a system can be determined: DE = Energy can be transferred as heat + work DE = Ef – Ei= q + w Where work = force X distance OR pressure X volume change w = p DV Note: Expansion volume = -pDV. Work is being done by the system on the surroundings.

  12. Work in Chemical Reactions Often times in chemical reactions, there isn’t a volume change or a force applied over a distance. Therefore, for the chemist, heat changes are the primary concern when considering the energy changes that exist in a chemical reaction. DE = q (for many, not all, chemical reactions)

  13. Focus on Heat Heat = thermal energy (motion of atoms) Something is “hot” because it has thermal energy and its atoms are moving rapidly. The more energy, the faster they move. The total thermal energy of a substance is the sum of the individual thermal energies of the atoms and molecules making up a substance.

  14. Chemists are Interested in Heat Transfer (q) during chemical reactions Heat is a form of energy transfer. Heat transfers from hot object to cooler object. q>0 Heat is transferred from surroundings to system. ENDOTHERMIC q<0 Heat is transferred from system to surroundings. EXOTHERMIC

  15. Remember the Trick to Temperature A measure of heat and relates to the average kinetic energy of atoms in a sample. The higher the temperature, the greater the thermal motion. Thermal energy depends upon sample size. Heat ≠ Temperature Thermometers display changes in thermal energy.

  16. Exothermic Process By convention in an exothermic process, heat is transferred FROM the system TO the surroundings. Heat is a product. 2H2 (g) + O2 (g) → 2H2O (l) + energy Combustion reactions are exothermic! Chemical (potential) energy is converted to heat (kinetic energy).

  17. Endothermic Reaction Heat is transferred TO the system FROM the surroundings. Heat is a reactant. Energy + 2HgO (s) → 2Hg(l) + O2 (g) Thermal energy (kinetic) is converted to chemical energy (potential).

  18. Specific Heat and Heat Capacity Specific heat = the heat needed to raise the temperature of one gram of a substance by 1.0 0 C. Units J/g0C. Intensive property- not dependent on how much is present. Heat Capacity = the measure of an overall effect of heat transfer on the temperature of an object. Units J/g . Extensive property (depends on mass).

  19. Specific Heats of Various Substances

  20. Side Note: Specific Heat of Water The specific heat of water is one of the highest known specific heats. That’s why temperature fluctuations are smaller near large bodies of water like Puget Sound than someplace like the desert in Arizona.

  21. Calorimetry The laboratory technique used to measure the heat released or absorbed during a chemical or physical change. In a calorimeter, there is no volume change in the system, therefore, DE = q q = mcDT

  22. Calorimetry  heat absorbed = mcDTliquid + mcDTcalorimeter Calorimetry depends on the assumption that all the heat involved changes the contents of the calorimeter and its contents no heat is gained or lost through the environment. For many chemical reactions, a styrofoam cup is a convenient calorimeter because it has a very low heat capacity (for most purposes negligible) and it is a good insulator.

  23. Sample Calculation- Specific Heat A piece of iron with a mass of 72.4 grams is heated to 100.oC and plunged into 100.g of water that is initially at 10.0oC. Calculate the final temperature that is reached assuming no heat loss to the surroundings. Cwater = 4.18 J/goC Ciron =0.449 J/goC NOTE: the heat gained by the cooler body + the heat lost by the warmer body = 0 Answer: Tfinal= 16.5oC

  24. Sample Calorimetry Problem A 1.5886 g sample of glucose was ignited in a bomb calorimeter. The temperature increased by 3.682oC. The heat capacity of the calorimeter was 3.52 kJ/oC, and the calorimeter contained 1.000 kg of water. Find the total heat released and the molar heat of reaction. Remember: Heat energy goes into heating up the water and the calorimeter itself. Answer: = -28.35 kJ -3211 kJ/mol

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