Thermochemistry Textbook PDF

Chapter 5 Thermochemistry  Section 1: The Nature of Chemical Energy Energy – capacity to do work or to produce heat Thermodynamics is the study of energy and its transformations. We will look at a type of thermodynamics called thermochemistry, or the relations...

Chapter 5 Thermochemistry  Section 1: The Nature of Chemical Energy Energy – capacity to do work or to produce heat Thermodynamics is the study of energy and its transformations. We will look at a type of thermodynamics called thermochemistry, or the relationships b/w chemical rxns and energy changes that involve heat. Types of energy : Radiant energy - solar energy Thermal energy - energy associated with random motion of atoms and molecules. Potential energy - energy available by virtue of an object’s position. All forms of energy can be converted from one form to another. Law of Conservation of Energy – energy cannot be created or destroyed, only converted. Potential energy – energy due to position or composition Kinetic energy – energy of motion  Section 2: First Law of Thermodynamics Almost all chemical rxns absorb or produce energy, generally in the form of heat. It is important to understand the difference b/w thermal energy and heat. Heat is the transfer of thermal energy b/w 2 bodies that are at different temperatures. The first law of thermodynamics states that energy cannot be created or destroyed but converted to different forms. To analyze energy changes associated with chemical rxns, we must first define the system. It is the specific part of the universe that is of interest to us. The surroundings are the rest of the universe outside the system. Open system, closed system, isolated system Exothermic process - gives off heat (transfers thermal energy to the surroundings) Endothermic process - heat has to be applied to the system  Internal energy of a system is the sum of all the kinetic and potential energies of the components of the system. ΔE = q + w ΔE = change in system’s internal energy, q = heat, w = work Will have # and a sign, either + or – to indicate the direction of energy flow ΔE is negative if exothermic and energy leaves, work done on surroundings ΔE is positive if endothermic and energy enters, work done on system  Example Calculate ΔE for a system undergoing an endothermic process in which 15.6kJ of heat flows and where 1.4kJ of work is done on the system. ΔE = q + w ΔE = 15.6 + 1.4 = 17.0kJ – positive because endothermic and also b/c work is done on the system  Section 3: Enthalpy Enthalpy – H, a way to measure the flow of heat for a system. Will use ΔH, or change in enthalpy: ΔH = Hproducts – Hreactants If products have more enthalpy, heat will be absorbed, ΔH will be positive, endothermic. If reactants have more enthalpy, heat is released, ΔH is negative, exothermic.  Section 4: Enthalpies of Rxn The enthalpy change that accompanies a rxn is called the enthalpy of rxn or the heat of rxn. 2H2(g) + O2(g) → 2H2O(g) △H = -483.6kJ The negative sign tells us it is exothermic. Enthalpy is extensive property. The enthalpy change for a rxn is equal in magnitude but opposite in sign to the reverse rxn. The enthalpy change for a rxn depends on the states of the reactants and products.  Example When 1 mole of methane (CH4) is burned at constant pressure, 890kJ of energy is released as heat. Calculate ΔH for a process in which a 5.8g sample of methane is burned at constant pressure. 5.8gCH4 x 1mol = 0.36molCH4 16.0g 0.36molCH4 x –890kJ = -320kJ mol CH4  Section 5: Calorimetry Science of measuring heat, based on observing temp change when system absorbs or discharges energy as heat. Calorimeter is instrument used experimentally to measure heat change. The specific heat of a substance is the amount of heat required to raise the temp of one gram of the substance by 1 degree C. Units are J/g°C Molar heat capacity is the heat capacity of one mole of a substance. 2 cups nested inside each other Inner cup holds soln in which rxn occurs Measure temps initially and then after rxn starts If soln heats up, heat is released from rxn, so exothermic If soln cools down, heat is absorbed for rxn, so endothermic  Example of Calorimetry We mix 50.0mL of 1.0M HCl at 25C with 50.0mL of 1.0M NaOH at 25C in a calorimeter. After mixing by stirring, temp raises to 31.9C. Calculate the change in enthalpy. H+ + OH- → H2O H = spec. heat cap. x mass x change in temp (final – initial) H = C m ΔT H = (4.18)(1.00 x 102)(6.9) m →d = m/v 1.00 = m/1.00 x 102 H = - 2.9 x 103J  Another Calorimetry Problem 45mL of 1.0M sulfuric acid is mixed with 45mL of sodium hydroxide in a calorimeter. Initially, the soln is at 23.3C and then the temp increases to 44.7C. Calculate the enthalpy of the soln in kJ/mol. 1.) q = (90.)(4.18)(21.1) 2.) q = -7900J = -7.9kJ 3.) 1.0M = mol = 0.045mol.045L 4.) -7.9kJ/.045mol = -180kJ/mol  Section 6: Hess’s Law When going from a particular set of reactants to a particular set of products, the change in enthalpy is the same whether the rxn takes place in one step or in a series of steps. N2(g) + 2O2(g) → 2NO2(g) ΔH = 68kJ Happens in 2 distinct steps: N2 + O2 →2NO ΔH = 180kJ 2NO + O2 → 2NO2 ΔH = -112kJ N2 + 2O2 → 2NO2 ΔH = 68kJ  Characteristics of Enthalpy 1.If rxn is reversed, sign of ΔH is reversed. 2.The magnitude of ΔH is directly proportional to the quantities of reactants and products in the rxn. If the coefficients in a balanced equation are multiplied by an integer, the value for ΔH is multiplied by the same integer.  Hess Example Calculate the enthalpy for this rxn. 2C(s) + H2(g) → C2H2(g) H = ??? kJ Given the following chemical equations: C2H2(g) + 5/2 O2(g) → 2 CO2(g) + H2O(l) H = -1299.5kJ C(s) + O2(g) → CO2(g) H = -393.5kJ H2(g) + ½ O2(g) → H2O(l) H = -285.8kJ  Enthalpy Example #2 Diborane is a highly reactive boron hydride. Calculate ΔH for the synthesis of diborane from its elements. 2B(s) + 3H2(g) → B2H6(g) Rxns: a) 2B(s) + 3/2O2(g) → B2O3(s) ΔH = -1273kJ b) B2H6(g) + 3O2(g) → B2O3(s) + 3H2O(g) ΔH = -2035kJ c) H2(g) + 1/2O2 (g) → H2O(l) ΔH = -286kJ d) H2O(l) → H2O(g) ΔH = 44kJ Need to combine rxns a, b, c, d to get all parts of main rxn; we need B, H2, and B2H6  Example #2 Cont. a)2B(s) + 3/2O2(g) → B2O3(s) ΔH=-1273kJ b) B2O3(s) + 3H2O(g) → B2H6(s) + 3O2 ΔH=2035kJ 2B(s) + B2O3(s) + 3/2O2(g) + 3H2O(g) → B2O3(s) + B2H6(s) + 3O2(g) ΔH=762kJ cancel species that appear on both sides of arrow **2B(s) + 3H2O(g) → B2H6(s) + 3/2O2(g) ΔH=762kJ Example continued… Need to get rid of H2O and O2 and gain H2, so need rxn c and d 2B(s) + 3H2O(g) → B2H6(s) + 3/2O2(g) ΔH=762kJ c) 3[H2(g) + 1/2O2(g) → H2O(l)] ΔH=3(-286kJ) 2B(s) + 3H2(g) + 3/2O2(g) + 3H2O(g) → B2H6(s) + 3/2O2(g) + 3H2O(l) ΔH=-96kJ can cancel 3/2O2(g) but not H2O b/c in different phases 2B(s) + 3H2(g) + 3H2O(g) → B2H6(s) + 3H2O(l) ΔH=-96kJ Example continued… 2B(s) + 3H2(g) + 3H2O(g) → B2H6(s) + 3H2O(l) ΔH=-96kJ d) 3[H2O(l) → H2O(g)] ΔH=3(44kJ) 2B(s) + 3H2(g) + 3H2O(g) + 3H2O(l) → B2H6(s) + 3H2O(l) + 3H2O(g) ΔH=+36kJ 2B(s) + 3H2(g) → B2H6(s) ΔH=+36kJ  Another Hess’s Law Example 2O3(g) → 3O2(g) ΔH = -427kJ O2(g) → 2O(g) ΔH = +495kJ NO(g) + O3(g) → NO2(g) + O2(g) ΔH = -199kJ Calculate ΔH for the rxn NO(g) + O(g) → NO2(g)  Section 7: Enthalpies of Formation If substances are at standard states, they are 1atm and often 298K, rather than 273K as with STP. Use degree sign to indicate at standard state We have a chart with the △H° for the elements and cmpds and we can use them to calculate the standard enthalpy of rxn, △H°rxn. We use Hess’s Law to find it. We can also use the rxn to find it.  Using Enthalpies to calculate the enthalpies of rxn △H° = Σn△Hf(products) - Σn△Hf(reactants) Calculate the standard enthalpy of formation for the combustion of 1 mol of methane. CH4 + 2O2 → CO2 + 2H2O △H° = [CO2 + 2H20] - [CH4 + 2O2] △H° = [1(-393.5) + 2(-285.8)] - [1(-74.81) + 2(0) △H° = -890.29kJ = -890.3kJ Calculate the standard enthalpy of formation for the formation of water.  Section 8: Bond Enthalpies The energy changes that accompany chemical rxns are closely related to the changes associated with forming and breaking chemical bonds. Bond enthalpy is the enthalpy change for the breaking of a particular bond in one mole of a gaseous substance.  Section 9: Foods and Fuels Most chemical rxns used for the production of heat are combustion rxns. The energy released when one g of any substance is combusted is the fuel value. The fuel value of any food or fuel can be measured by calorimetry. Most of the energy our bodies need comes from carbs and fats. The body uses the chemical energy from foods to maintain body temp to contract muscles and to construct and repair tissues. Any excess is stored as fat. For fuels, we have fossil fuels: natural gas, petroleum, coal. They bonds break and release heat that we can use as well...to heat our homes and make electricity.