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Thermochemistry

Thermochemistry

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Thermochemistry

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  1. Thermochemistry Chapter 17 Pages 511-524

  2. Phase Diagrams -a graph of pressure vs. temp that shows the conditions under which the phases of a substance exist Triple point – indicates the T and P conditions at which the solid, liquid, and vapor of the substance coexist at eq Critical point – indicates the critical T and critical P Critical T - T above which the substance can’t exist in the liquid state Critical P – lowest P at which the substance can exist as a liquid at the critical T

  3. Phase Diagram Critical Pressure Critical Point Solid Liquid Normal Freezing Point Pressure (atm) 1.00 Normal Boiling Point Gas Triple Point Critical Temperature Temperature C

  4. Thermochemistry • Study of the transfers of energy as heat that accompany chemical reactions and physical changes • All chemical reactions are accompanied by a change in energy • Usually absorb (endothermic) or release (exothermic) energy as heat

  5. Calorimeter -instrument for measuring the energy absorbed or released as heat in a chemical or physical change -data collected are temp changes because heat cannot be measured directly; but temp, which is affected by the transfer of energy as heat, is directly measurable

  6. Temperature • The measure of the average kinetic energy of the particles in a sample of matter • The greater the kinetic energy, the higher the temp, and the hotter it feels • Measured in °Celsius or Kelvins

  7. Joule (J) -the SI unit of heat as well as all other forms of energy -measures the amount of energy transferred as heat A joule is relatively small so the kilojoule, kJ, is commonly used

  8. Heat -the energy transferred between samples of matter because of a difference in their temperatures -energy moves spontaneously from higher to lower temp; when the temps are equal, energy is no longer transferred

  9. At the boiling point energy is being transferred to the system, but until all the water is vaporized the temperature will not increase. The transfer of heat may not affect the temperature. This is particularly true during phase changes. In this example the temperature of ice at the melting point will not increase until all the ice has melted. The same is true at the boiling point.

  10. Heat Capacity -the quantity of energy transferred as heat depends on the nature of the material changing temp, the mass changing temp, and the size of the temp change -a quantity called specific heat can be used to compare heat absorption capacities for different materials

  11. Specific Heat Definition: the amount of energy required to raise the temperature of one gram of substance by one Celsius degree or one Kelvin Finding specific heat: Energy in Joules q cp = m T Change in temp in ºC or K Specific heat is measured in J/g K (and is under constant pressure) Mass in grams

  12. You will frequently see this equation written q (or Q) = cpmT which is the way it appears on the TAKS formula chart. As written, the equation is being solved for energy in J rather than specific heat. Change in temperature can be either Celsius or Kelvin as the magnitude of the change is equal for either.

  13. Heat of Reaction -the quantity of energy released or absorbed as heat during a chemical reaction -can be thought of as the difference between the stored energy of the reactants and the products Example: in the formation of water from its elements energy is released and the reaction is exothermic 2H2(g) + O2(g)  2H2O(g) + 483.6 kJ These reactions follow the law of conservation of energy.

  14. Thermochemical equation 2H2(g) + O2(g)  2H2O(g) + 483.6 kJ This is an example of a thermochemical equation, an equation that includes the quantity of energy released or absorbed as heat during the reaction as written. In any thermochemical equation the coefficients are treated as moles. The quantity of energy released as heat depends on the amounts of reactants and products. Producing twice as much water would require twice as many moles of reactants and would release 2 x 483.6 kJ of energy. Producing ½ as much would require ½ as many moles of reactants and would release ½ as much energy. Fractional coefficients are sometimes used.

  15. Enthalpy ΔH ΔH – the energy absorbed or released as heat during a chemical reaction at constant pressure H is the symbol for a quantity called enthalpy. Enthalpy has traditionally been defined as the heat content of a system at constant pressure. ΔH is read as “change in enthalpy” An enthalpy change is the amount of energy absorbed or lost by a system as heat during a process at constant pressure. ΔH = Hproducts - Hreactants

  16. Exothermic Reactions -reactants have a higher energy than the products -∆H is always negative (system loses energy) -energy is included in the products and is measured in kJ/mol 2H2(g) + O2(g) 2H2O(g) + 483.6 kJ/mol -for this equation ∆H = -483.6 kJ/mol

  17. Exothermic Reaction Pathway The energy of the products is less than that of the reactants; therefore energy in the form of heat is released and ∆H is negative.

  18. Endothermic Reactions -products have a higher energy than reactants -∆H is always positive (the system gains energy) -energy is included in the reactants 2H2O(g) + 483.6 kJ  2H2(g) + O2(g) -for this reactions ∆H = 483.6 kJ/mol The physical states of reactants and products must always be included in thermochemical equations because they influence the overall amount of energy exchanged. EX: the energy needed for the decomposition of water would be greater than 483.6 kJ if we started with ice because extra energy would be needed to melt the ice and to change the liquid into a vapor.

  19. Endothermic Reaction Pathway The reactants have less energy than the products meaning the reactants have absorbed energy and ∆H is positive.

  20. Calculating Heats of Reaction ΔH = Hproducts - Hreactants This equation literally means that the change in enthalpy for a reaction is equal to the sum of bonds formed in the products minus the sum of the bonds broken in the reactants. Table A-14, page 902, in your textbook lists the heats of formation for many common substances. States, signs, and number of moles must be carefully addressed when using this equation.