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THERMODYNAMICS

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• a system: Some portion of the universe that you wish to study a glass of water plagioclase the mantle Changes in a system are associated with the transfer of energy lift an object: stored chemical  potential drop an object: potential  kinetic pump up a bicycle tire: chemical  mechanical  heat (friction + compression) add an acid to a base: chemical  heat
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1. 1. Thermodynamics a system: Some portion of the universe that you wish to study the surroundings: The adjacent part of the universe outside the system Changes in a system are associated with the transfer of energyNatural systems tend toward states of minimum energy
2. 2. Energy StatesUnstable:Stable:Metastable: in low-energy perch Figure 5-1. Stability states. Winter (2001) An Introduction to Igneous and Metamorphic Petrology. Prentice Hall.
3. 3. A review of basic thermodynamics: A refresher The ball represents mass exchange The arrow represents energy exchange
4. 4. The First Law of Thermodynamics• Heat and work are equivalent• Energy is conserved in any transformation• The change of energy of a system is independent of the path taken Energy can be neither created nor detroyed ∆E = q - w or dE = dq - dw dE = dq - P dV E = internal energy P = pressure q = heat V = volume w = work
5. 5. EnthalpydE = dq - P dVH = E + PVdH = dqH = enthalpyThe change in the enthalpy of a system (∆H) during a reversiblechange in state at constant pressure is equal to the heat absorbedby the system during that change in state.The enthalpy of formation of compounds and their ions andmolecules in aqueous solution is the heat absorbed or given off bychemical reactions in which the compounds, ions, and moleculesform from the elements in the standard state (25°C, 1 atm)
6. 6. Heats of Reaction ∆H = Σ nH (products) - Σ nH (reactants) n = molar coefficient of each reactant/productWhen ∆H is positive, the reaction is endothermic (heat flowsfrom the surroundings to the system);When ∆H is negative, the reaction is exothermic (heat flowsfrom the system to the surroundings
7. 7. Heats of Reaction ∆H = Σ nH (products) - Σ nH (reactants)For example, evaporation: H2O(l) H2O(g)∆H = H(H2O(g)) - H(H2O(l))∆H = (-57.80) - (-68.32) = 10.52 kcalThe reaction is endothermic (i.e., sweating is a mechanism for cooling the body)
8. 8. Heat CapacityWhen heat is added to a solid, liquid, or gas, the temperatureof the substance increases: dq = C dT dq = dH dH = C dT, at constant pressure (important in geochemistry) C = heat capacity T = temperatureHeat capacities vary with temperature…
9. 9. The Second Law of Thermodynamics• It is impossible to construct a machine that is able to convey heatby a cyclical process from one reservoir at a lower temperature toanother at a higher temperature unless outside work is done(i.e, air conditioning is never free)• Heat cannot be entirely extracted from a body and turned into work(i.e., an engine can never run 100% efficiently) — a certain fractionof the enthalpy of a system is consumed by an increase in entropy• Every system left to itself will, on average, change toward acondition of maximum randomness — entropy of a systemincreases spontaneously and energy must be spent to reverse thistendency
10. 10. The entropy of the universe always increases or“You can’t shovel manure into the rear end of a horse and expect to get hay out of its mouth”
11. 11. Entropy ∆S = Σ nS (products) - Σ nS (reactants)For example: H2O(l) H2O(g)∆S = S(H2O(l)) - S(H2O(g))∆S = 45.10 - 16.71 = 28.39 cal/degWhen ∆S is positive, entropy of the system increases withthe change of state;When ∆S is negative, entropy decreases
12. 12. The fundamental equation of thermodynamicsThe ratio of heat gained or lost to temperature will alwaysbe the same, regardless of path, for a reversible reaction dE = T dS - P dV When dE = 0, T dS = P dV Look familiar? dP ∆S dS = dq/T = P dV/T = dT ∆V dS = dq/T — reversible process dS > dq/T — irreversible process