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Thermody-nizzle-amics

Thermody-nizzle-amics. A (Josh)^2 Production. Heating Shindig. Amount of energy needed to change a given substance a given temperature depends on; The mass you are heating (how much) It’s specific heat (how well it can absorb energy)

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Thermody-nizzle-amics

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  1. Thermody-nizzle-amics A (Josh)^2 Production

  2. Heating Shindig • Amount of energy needed to change a given substance a given temperature depends on; • The mass you are heating (how much) • It’s specific heat (how well it can absorb energy) • Some random stuff: Specific heat changes from phase to phase; Water has high specific heat (4.186 J/g*K); Metals have low specific heat • Q = MCT

  3. Shifting Phaz0rz • Phase change occurs at a constant temperature • How much energy is required to change a substance from one phase or another depends on: • The mass of the substance present • A constant called Heat of fusion (S <->L) or heat of vaporization (L<->G) • Q = KM

  4. Sample Heating Curve of Californium

  5. Chemical Systems Chemical systems simply refer to the actual chemicals (the system) and it’s surroundings System Surroundings

  6. Enthalpy is Hot (sometimes) • Enthalpy refers to how much energy is stored within a chemical system • When a reaction occurs, there is a change in enthalpy, H • A positive change in enthalpy, H > 0, means that energy is entering the system. Such a process is known as endothermic. • A negative change in enthalpy, H < 0, means that energy is leaving the system. This is an exothermic process.

  7. But wait… Q: If the energy in the system is going down during an exothermic process, why does it seem exothermic processes make things hot? A: Cuz I said so… and because you’re in the surroundings gaining the energy lost from the system.

  8. How to Find Change in Enthalpy • Bond Energies • Heat of Formation • Hess’ Law

  9. Bond Energies • Given a table of bond energies, find the total bond energy of the products and of the reactants. Remember coefficients • H = ∑Bond EnergyReactants - ∑Bond EnergyProducts • Note: this is the only one where it’s reactants minus products

  10. Heat of Formation • This one’s even easier. First find the heat of formation of the products and reactants (It’ll be given, unless…) • All elemental substances (like O2, Fe, etc.) have H°F = 0 J • To find enthalpy; H °rxn = ∑ H°FProducts - ∑ H°FReactants

  11. My Homedawg Hess Hess’ Law: If a reaction is the sum of two or more other reactions, the enthalpy change for the overall process is the sum of the enthalpy changes for the constituent reactions. Sweet.

  12. Let’s make methane!?! • Use these reactions to find H°Fof CH4 • C + O2 CO2 H = -393.5 kJ • H2 + ½ O2  H2O H = -285.8 kJ • CH4 + 2O2  CO2 + 2H2O H = -890.3 kJ

  13. Stop defying Entropy; go combust right now! • Entropy is a measure of disorder in a system. • It is measured by how much energy a substance has compared to its absolute temperature. S = q/T (J/(mol * K)) • Chaotic things have greater entropy; gas has more entropy than liquid, and liquid more than solid. Stone’s hair has lots of entropy. • Potentially useful: S rxn = ∑ SProducts - ∑ SReactants

  14. Hippopotami can cause increases in entropy

  15. How the Universe Works • Everything moves towards maximum entropy and minimum enthalpy • Basically, everything wants to spread its matter and energy across the universe • 2nd Law of Thermodynamics: The Entropy of the Universe is always increasing. (don’t worry about the others)

  16. The work of a True Hero, J. Willard Gibbs. S Universe = S System + S Surroundings S Universe = S System - H System /T -TS Universe = H System - TS System -TS Universe is given as G, change in Gibbs free energy. Since S Universe is positive if G is negative, if G < 0, than a reaction is spontaneous (IE product favoring) Summary: G = H System - TS System

  17. Possible Reaction Types • Exothermic and increasing entropy; always spontaneous • Endothermic and increasing entropy; spontaneous at higher temperatures • Exothermic and decreasing in entropy; spontaneous at lower temperatures • Endothermic and decreasing in entropy; never spontaneous

  18. Gibbs and Q and K Sorry, no proof this time G = G° + RT lnQ R = 8.314 J/ (Mol * K) If reaction is at equilibrium, G = 0 so G° = -RT lnK

  19. Applications

  20. Throwing Hot Blocks of Metal into Water or Jello or maybe Ethanol • Remember energy is conserved and the system will reach equilibrium at a common temperature • Basic form M1C1(Tf – Ti1) = M2C2(Ti2 – Tf)

  21. Calorimeters • Calorimeters measure reaction enthalpy by recording change in surroundings. • Calorimeters will often have their own specific heat, but it factors in their mass so Q = C T for calorimeters. It is assumed they don’t melt. • Enthalpy is found by measuring the change in temperature of the calorimeter, calculating its increase in energy, and comparing that to how much stuff was reacted.

  22. Chipotle Burritos Have Lots of Calories. Gana can eat three of them. Gana is a true hero just like J. Willard Gibbs.

  23. Melting/Freezing and Evaporation/Condensation Points All phase changes are increase in entropy and endothermic or vice versa. The reaction has no favored side when G is zero, therefore that temperature is the melting/freezing etc. point. From the earlier equation Tphase change = Hphase change / Sphase change

  24. FIN

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