Chapter 6 Thermochemistry.notebook Chapter 6: Thermochemistry Jan 29 1:37 PM 6.1 The Nature of Energy Thermodynamics: The study of energy and its interconversions Energy: the capacity to do work or to produce heat. The law of conservation of energy, also called the first law of thermodynamics states that energy can be converted from one form to another, but can be neither created nor destroyed. Another way of stating the first law is that THE ENERGY OF THE UNIVERSE IS CONSTANT. Potential energy: energy due to position or composition Kinetic energy: energy due to motion of an object Heat: Involves transfer of energy between two objects. System: that part of the universe on which attention is to be focused Surroundings: everything in the universe surrounding a thermodynamic system Dec 9 10:00 AM 1
Chemical Energy Exothermic Reaction CH 4 + O 2 CO 2 + H 2 O + energy Alkali Metals +H 2 O Aluminum + Bromine In an exothermic process, the bonds in the product molecules are stronger (on average) than those in the reactant molecules. The net result is that the quantity of energy ΔE is transferred to the surroundings as heat when reactants are converted to products. Dec 9 10:00 AM Endothermic Reaction N 2 + O 2 + energy 2NO 2 For an endothermic process, energy flows into the system from the surroundings as heat to increase the potential energy of the system. In an endothermic process, the products have higher potential energy (weaker bonds on average) than the reactants. Barium hydroxide octahydrate + ammonium nitrate Dec 9 10:00 AM 2
The SIGN of the energy change must be viewed from the point of view of the SYSTEM. Diagram of Exothermic and Endothermic Process ΔE = means the system loses energy ΔE = + means the system gains energy. Jan 18 11:21 PM 6.2 Enthalpy and Calorimetry Solve problems relating to both enthalpy and calorimetry. Enthalpy (H). At constant pressure the enthalpy equals the energy flow as heat. where ΔH is the H final minus H initial and q is heat, or ΔH = H products H reactants If ΔH > 0, the reaction is endothermic. (Heat is absorbed by the system.) If ΔH < 0, the reaction is exothermic. (Heat is given off by the system.) Example 6.2A Enthalpy When solid potassium hydroxide pellets are added to water, the following reaction takes place: NaOH(s) NaOH(aq) For this reaction at constant pressure, ΔH = 43 kj/mol. Answer the following questions regarding the addition of 14g of NaOH to water: a. Does the beaker get warmer or colder? b. Is the reaction exo or endothermic? c. What is the enthalpy change for the dissolution? Dec 9 10:59 AM 3
Calorimetry is the experimental technique used to determine the heat exchange/flow (q) associated with a reaction. At constant pressure, q = ΔH At constant volume, q = ΔE In both cases, however, heat gain or loss is being determined. The amount of heat exchanged in a reaction depends upon: The net temperature change during the reaction. The amount of substance. The more you have, the more heat can be exchanged. The heat capacity (C p ) of a substance. A coffee cup calorimeter made of two styrofoam cups C p = heat absorbed Increase in temperature x mass = J/g C Some substances (such as water) can absorb more heat than others for a given temperature change. Dec 10 8:10 AM There are three ways of expressing heat capacity. Heat capacity (as above) = J/ C Specific heat capacity = heat capacity per gram of substance (J/g C or J/g K ) Molar heat capacity = heat capacity per mole of substance (J/mol C or J/mol K ) Remember you must use dimensional analysis to solve calorimetry problems! Example 6.2B Specific Heat Capacity Look at Table 6.1. Based on the values for specific heat capacity, which conducts heat better, water or aluminum? Why is this important in cooking? Dec 10 8:10 AM 4
Constant Pressure Calorimetry using the "coffee cup calorimeter". ΔH = heat absorbed = specific heat capacity x mass x change in temperature ΔH = C p x m x ΔT See Examples p.244 246 Example 6.2C Constant Pressure Calorimeter Recall from example 6.2A that the enthalpy change (ΔH) per mole of NaOH is 43 kj/mol when NaOH(s) NaOH(aq) If 10.0 g of solid NaOH is added to 1.00 L of water (specific heat capacity 4.18J/ C g) at 25.0 C in a constant pressure calorimeter, what will be the final temperature of the solution? (Assume the density of the final solution is 1.05 g/ml) Dec 10 8:10 AM Heating/Cooling Curves Freezing is the phase change as a substance changes from a liquid to a solid. Melting is the phase change as a substance changes from a solid to a liquid. Condensation is the phase change as a substance changes from a gas to a liquid. Vaporization is the phase change as a substance changes from a liquid to a gas. The heat of fusion is the amount of energy given off when a substance freezes (or the amount of energy the substances requires to liquify). The heat of vaporization is the amount of energy given off when a gas liquifies (or the amount of energy needed to vaporize the liquid). Dec 13 9:34 AM 5
The following diagram shows the uptake of heat by 1 kg of water, as it passes from ice at 50 ºC to steam at temperatures above 100 ºC, affects the temperature of the sample. A: Rise in temperature as ice absorbs heat. Q = mcpδt Cp (for solid) = 0.44 cal/g c B: Absorption of latent heat of fusion (Lf). Q = mδhfus ΔHfus (of water) = 80 cal/g C: Rise in temperature as liquid water absorbs heat. Q = mcpδt Cp (for liquid) = 1 cal/g c D: Water boils and absorbs latent heat of vaporization (Lv). Q = mδhvap ΔHvap (of water) = 540 cal/g E: Steam absorbs heat and thus increases its temperature. Q = mcpδt Cp (for gas) = 0.48 cal/g c The above is an example of a heating curve. One could reverse the process, and obtain a cooling curve. The flat portions of such curves indicate the phase changes. Feb 26 10:47 PM 6.3 Hess's Law Use Hess s Law to calculate enthalpy changes for a variety of reactions. Since enthalpy changes are state functions, then it does not matter if ΔH for a reaction is calculated in one step or a series of steps. This is called Hess s Law. N 2 (g) + O 2 (g) 2NO(g) 2NO(g) + O 2 (g) 2NO 2 (g) Net Reaction: N 2 (g) + 2O 2 (g) 2NO 2 (g) ΔH 2 = 180 kj ΔH 3 = 112 kj ΔH 1 = ΔH 2 + ΔH 3 = 68 kj The sum of the two steps gives the net, or overall, enthalpy of the reaction. Dec 13 9:34 AM 6
Characteristics of Enthalpy Changes 1. We can reverse the entire equation. By doing this, the products become reactants and vice versa. If your reverse the equation, you must multiply ΔH by 1. (An exothermic reaction becomes endothermic and vice versa.) 2. We can multiply the entire equation by a factor such as 3, 2, ½, or 1 / 3. If you multiply an equation by 2, you must multiply ΔH by 2. 3. We can do both #1 and #2. The most important thing to keep in mind is that WHEN YOU MANIPULATE AN EQUATION, YOU MUST MANIPULATE THE ΔH VALUE IN EXACTLY THE SAME WAY! Example 6.3A Hess s Law Given the following reactions and ΔH values, a. 2N 2 O(g) O 2 (g) + 2N 2 (g) ΔH a = 164 kj b. 2NH 3 (g) + 3N 2 O(g) 4N 2 (g) + 3H 2 O(l) ΔH b = 1012 kj Calculate ΔH for: 4NH 3 (g) + 3O 2 (g) 2N 2 (g) + 6H 2 O(l). Jan 18 11:58 PM Example 6.3B Practice with Hess s Law Given the following reactions and ΔH values, a. B 2 O 3 (s) + 3H 2 O(g) B 2 H 6 (g) + 3O 2 (g) ΔH = +2035 kj b. 2H 2 O(l) 2H 2 O (g) ΔH = +88 kj c. H 2 (g) + ½O 2 (g) H 2 O (l) ΔH = 286 kj d. 2B(s) + 3H 2 (g) B 2 H 6 (g) ΔH = +36 kj Calculate ΔH for: 2B(s) + 3 / 2 O 2 (g) B 2 O 3 (s) Dec 15 8:26 AM 7
6.4 Standard Enthalpies of Formation Use your knowledge of standard states and standard enthalpies of formation to calculate ΔH for a variety of reactions. Standard Enthalpy of Formation (ΔH f ) is the change in enthalpy that accompanies the formation of one mole of a compound from its elements with all substances in their standard states. 1. ΔH f is always given per mole of compound formed. 2. ΔH f involves formation of a compound from its elements with the substances in their standard states. 3. Standard state conditions are: For an element: It is the form, which the element exists in at 25 C and 1 atm. For a compound: For a gas it is a pressure of exactly 1 atmosphere For a substance in solution, it is a concentration of exactly 1 M. For a pure solid or liquid, it is the pure solid or liquid. 4. ΔH f for an element in its standard state, such as Ba(s) or N 2 (g), equals 0. Dec 15 8:26 AM Example 6.4A Standard Enthalpies of Formation By consulting Appendix 4 of your textbook, and from your knowledge of standard states, list the standard enthalpy of formation for each of the following substances. a. Al 2 O 3 (s) b. Ti(s) c. P 4 (g) d. SO 4 2 (aq) e. F 2 (g) Jan 19 12:10 AM 8
The key to calculating standard enthalpy changes in reactions is to remember products minus reactants. ΔH reaction = n p ΔH f (products) n r ΔH f (reactants). This reads the sum of ΔH f for n moles of each of the products minus the sum of the ΔH f for n moles of each of the reactants. Pathway for the Combustion of Methane Dec 15 8:58 AM Schematic Diagram of Energy Changes Dec 15 9:21 PM 9
Remember!!! Just as in Hess s law problems, when you multiply the substance by an integer coefficient in a balanced equation, you must multiply the ΔH f value by that integer as well! Example 6.4B Calculating Standard Enthalpies of Formation Using the data in Appendix 4 in your textbook, calculate the ΔH for the following reaction 2C 3 H 6 (g) + 9O 2 (g) 6CO 2 (g) + 6H 2 O(l). Feb 2 12:27 PM Example 6.4C Practice with Standard Enthalpies The "thermite" reaction is discussed in Example 6.10 in your textbook. It is one in which molten iron is made from the reaction of aluminum powder and iron oxide. A variation on that reaction was described in the October 1984 Journal of Chemical Education. The reaction is, 2Al(s) + Cr 2 O 3 (s) Al 2 O 3 (s) + 2Cr(s) a. Calculate ΔH for this reaction. b. Which reaction yields more energy per gram of metal formed, thermite or this one? Dec 15 8:58 AM 10