Chapter 5 Thermochemistry Energy -Very much a chemistry topic Every chemical change has an accompanying change of. Combustion of fossil fuels The discharging a battery Metabolism of foods If we are to understand chemistry, we must understand the energy changes that accompany chemical reactions. Thermochemistry The study of energy and its transformations is known as. It studies the relationships between content of a system. In this chapter we will examine an aspect of thermodynamics that involves the relationships between chemical reactions and energy changes involving heat. This portion of thermodynamics is called. 5.1 The Nature of Energy The capacity to do work or to transfer heat.. A function. independent of the path, or how you get from point A to B. Work is energy used to cause an object. Work = is energy used to cause the temperature of an object to increase. The energy transferred as a result of difference. Kinetic and Potential Energy Kinetic Energy = Energy of E k = Mass Velocity Potential Energy = Energy of E p = Mass Gravitational constant = 9.8 m/s2 Height of the object relative to some reference height Potential Energy in Chemistry is an important kind of force for large objects, but not in chemistry. Gravitational force is negligible in extremely small particles such as and. More important are the forces that arise from. Potential Energy the interactions between charged particles. 1
Electrostatic Potential Energy kq1q2 E el = d Q1 and Q2 represent the on the two interacting objects, typically the charge on an electron 1.60 x 10-19 Coulombs d is the distance separating them K is a proportionality constant, 8.99 x 10 9 J-m/C 2 Joule-meter per Coulomb squared Electrostatic Potential Energy When Q1 and Q2 have the same sign (both positive, or negative), the two charges repel one another, pushing them apart, and Eel is. When they have opposite signs, they attract one another, pulling them toward each other, and Eel is. The lower the energy of a system, the more it is. Thus, the more strongly interact, the more stable the system. Chemistry and Energy The chemical energy of substances is due to the stored in the arrangements of the atoms of the substance. The energy possessed by a substance because of its temperature is associated with of the molecules. Units of Energy The SI unit for energy is the. in honor of James Joule, a British scientist who investigated work and heat. Joule = A mass of kg moving at a speed of 1 m/s possesses a kinetic energy of 1 joule Remember, E k = ½ mv 2 A small unit, so (kj) is often used when discussing the energies associated with chemical reactions. Calories A more familiar unit to us. The amount of energy needed to raise the temperature of 1 grams of water from 14.5 C to 15.5 C. 1 calorie = J (exactly) 1 Calorie = 1000 calories = 1 kcal. The Universe is divided into two halves. the and the. The system is the part you are concerned with. The surroundings are the rest. 2
reactions release energy to the surroundings. reactions absorb energy from the surroundings. Direction Every energy measurement has three parts. A unit ( Joules of calories). A number how many. and a sign to tell direction. negative - positive- Same rules for heat and work Heat given off is. Heat absorbed is. Work done by system on surroundings is. Work done on system by surroundings is. - The study of energy and the changes it undergoes. You try it A bowler lifts a 5.4 kg bowling ball from ground level to a height of 1.6 meters and then drops the ball back to the ground. What happens to the potential energy of the bowling ball as it is raised from the ground? What quantity of work, in J, is used to raise the ball? After the ball is dropped, it gains kinetic energy. If we assume that all the work done in part b (above) has been converted to kinetic energy by the time the ball strikes the ground, what is the speed of the ball at the instant just before it hits the ground? You try it What is the kinetic energy, in J, of An Argon atom moving with a speed of 650 m/s? Hint: 1 amu = 1.66 x 10-27 kg A mole of argon atoms moving with a speed of 650 m/s? 3
5.2 First Law of Thermodynamics The energy of the universe is it may be transferred between the and its. Law of of energy. Internal Energy The sum of all the and energy of all its components. Represented by E Generally don t know the actual numerical value We can measure the change in internal energy, Delta E, Remember the three parts of E, A number A unit (together these 2 give the magnitude of change) A sign that gives the direction Signs of E A positive value of E results when Efinal >E initial, indicating the system has gained energy from its surroundings. A negative value of E is obtained when E final < E initial, indicating the system has lost energy to its surroundings. Relating E to heat and work q = heat w = work E= q + w First Law of Thermodynamics in algebraic terms It takes point of view to decide signs. When heat is added to a system or work is done on a system, its internal energy. Sign Conventions for q, w and E For q For w For E State Functions Reminder, Internal Energy is a state function A property of a system that is determined by specifying the system's (in terms of temperature, pressure, location, etc.) 4
The value of a state function depends only on the of the system, not on the the system took to reach that state. E only depends on the and states of the system, not on how the change occurs. Analogy Suppose you are traveling between Chicago and Denver. Chicago is 596 ft above sea level; Denver is 5280 ft above sea level. No matter what route you take, the altitude change will be 4684 ft. The distance you travel, however, will depend on your route. is a state function is not a state function E= q + w E= q and w Thus E depends only on the and states of the system, regardless of how the transfers of energy occur in terms of heat and work. 5.3 Enthalpy The chemical and physical changes that occur around us essentially occur at constant. Most commonly, the only kind of work produced by chemical and physical changes open to the atmosphere is the work associated with change in the of the system. Pressure-Volume Work The work involved in the expansion or compression of gases if called P-V work. When pressure is constant, as in our example, the sign and magnitude of the pressurevolume work is given by w = The negative sign is necessary to conform to the sign conventions given earlier. w = - P V When the volume, V is a positive quantity and w is a negative quantity. Energy leaves the system as work, indicating that work is done the system the surroundings. When the volume, V is a negative quantity and w is a positive quantity. Energy enters the system as work, indicating that work is done the system the surroundings. Enthalpy, denoted H From a Greek word meaning to warm Accounts for heat flow in processes occurring at when no forms of work are performed other than P-V work. 5
H = (its definition) Is a because internal energy, pressure, and volume are all state functions. Enthalpy When a change occurs at constant pressure, the change in enthalpy is given by the following relationship: H = the heat at constant pressure q p can be calculated from E = = q p = = H Enthalpy The change in enthalpy equals the heat at constant pressure. Because q p is something we can either measure or readily calculate and because so many physical and chemical changes occur at constant pressure, enthalpy is a more useful function than internal energy. For most reactions the difference in H and E are so small because P V is small. Think about this Suppose we confine 1 gram of butane and sufficient oxygen to completely combust it in a cylinder capped with a piston. The cylinder is perfectly insulating, so no heat can escape to the surroundings. A spark initiates combustion of the butane, which forms carbon dioxide and water vapor. If we use this apparatus to measure the enthalpy change in the reaction, would the piston rise, fall, or stay the same? Explain. 5.4 Enthalpies of Reaction Because H = Hfinal Hinitial, the enthalpy change for a chemical reaction given by the enthalpy of the products minus the enthalpy of the reactants. H = the enthalpy change that accompanies a reaction Sometimes called Hrxn Thermochemical equations Balanced chemical equations that show the associated enthalpy change 2 H2(g) + O2 (g) 2 H20 (g) H = -483.6 kj Would this be exothermic or endothermic? This reaction can be shown on an energy diagram: 6
Guidelines when using thermochemical equations and energy diagrams: 1. Enthalpy is an property H is directly proportional to the amount of reactant consumed in the process. 2. The enthalpy change for a reaction is l in magnitude, but in sign, to the H for the reverse reaction. 3. The enthalpy change for a reaction depends on the of the reactants and products. Important to note the states in your equations. You Try It How much heat is released when 4.5 g methane gas is burned in a constant pressure system? H = -890 kj per mole of methane. Hydrogen peroxide can decompose to water and oxygen by the following reaction: 2H2O2 (l) 2H2O (l) + O2(g) H= -196 kj Calculate the value of q when 5 grams of peroxide decomposes at constant pressure. Calorimetry Measurement of. Use a. The temperature change experienced by an object when it absorbs a certain amount of heat. It is a unique value for all substances. the amount of heat required to raise its temperature by 1 K. The greater the heat capacity, the greater the heat required to produce a given increase in temperature. Molar Heat Capacity The heat capacity of a substance. Cmolar Specific heat capacity (specific heat) the heat capacity of one gram of a substance Give formulas: Specific Heat Values Examples: N2(g) = 1.04 J/g-K Al(s) = 0.90 J/g-K Fe(s) = 0.45 J/g-K H2O (l) = 4.18 J/g-K What do you notice about the value of water compared to others? Why would this be important to us? 7
You Try It How much heat is needed to warm 250 grams of water from 22 C to near its boiling point, 98 C? What is the molar heat capacity of water? Constant-Pressure Calorimetry Uses a coffee-cup calorimeter Not sealed so reaction occurs at constant pressure = For an reaction, heat is lost by the reaction and gained by the solution, so the temperature of the solution rises. Opposite is true for an endothermic reaction. Measuring H Using a Coffee-Cup Calorimeter When a student mixes 50 ml of 1.0 M HCl and 50 ml of 1.0 M NaOH in a coffee-cup calorimeter, the temperature of the resultant solution increases from 21.0 C to 27.5 C. Calculate the enthalpy change for the reaction in kj/mol, assuming that the calorimeter loses only a negligible quantity of heat, that the total volume of the solution is 100 ml, that its density is 1.0 g/ml, and that its specific heat is 4.18 J/g-K. More Practice When 50 ml of 0.1 M silver nitrate and 50 ml of 0.1 M hydrochloric acid are mixed in a constant-pressure calorimeter, the temperature of the mixture increases from 22.20 C to 23.11 C. Calculate H for this reaction in kj/mol silver nitrate, assuming that the combined solution has a mass of 100 g and a specific heat of 4.18 J/g-C 8
Bomb Calorimetry Constant volume calorimeter is called a bomb calorimeter. Material is put in a container with pure oxygen. Wires are used to start the combustion. The container is put into a container of water. The heat capacity of the calorimeter is known and tested. Since V = 0, P V = 0, E = q You Try It Methylhydrazine (CH6N2) is commonly used as a liquid rocket fuel. The combustion of methylhydrazine with oxygen produced N2(g), CO2(g) and H2O (l). Write the balanced equation. When 4 grams of methylhydrazine is combusted in a bomb calorimeter, the temperature of the calorimeter increases from 25.00 C to 39.50 C. In a separate experiment the heat capacity of the calorimeter is measured to be 7.794 kj/c. What is the heat of reaction for the combustion of a mole of methylhydrazine? Properties properties not related to the amount of substance. density, specific heat, temperature. property - does depend on the amount of stuff. Heat capacity, mass, heat from a reaction. 5.6 Hess s Law Enthalpy is a state function. It is. We can add equations to come up with the desired final product, and add the Two rules If the reaction is reversed the sign of H is changed If the reaction is multiplied or divided, so is H H Using Hess s Law to Calculate H The following information is known: C(s) + O2(g) CO2(g) H1 = -393.5 kj CO(g) + ½ O2 (g) CO2 (g) H2 = -283.0 kj Using these data, calculate the enthalpy for: C(s) + ½ O2(g) CO(g) 9
More Practice with Hess s Law Calculate H for the reaction 2C(s) + H2(g) C2H2(g) Given the following chemical equations and their respective H. C2H2(g) + 5/2O2 2CO2(g) + H2O (l) H = - 1299.6 kj C(s) + O2(g) CO2(g) H = -393.5 kj H2(g) + ½ O2(g) H2O(l) H = -285.8 kj You Try It Calculate H for the reaction NO(g) + O(g) NO2 (g) Given the following information: NO(g) + O3(g) NO2(g) + O2(g) O3(g) 3/2 O2(g) O2(g) 2 O (g) H = -198.9 kj H = -142.3 kj H = 495.8 kj Remember H is a state function, so for a particular set of reactants and products, whether the reaction takes place in one step or in a series of steps. H is the same 5.7 Enthalpies of Formation The enthalpy change associated with the of a compound from its constituent elements. Labeled The enthalpy change for a reaction at standard conditions (25ºC, 1 atm, 1 M solutions). These are standard state conditions. Symbol When using Hess s Law, work by adding the equations up to make it look like the answer. The other parts will cancel out. 10
Standard Enthalpies of Formation Hess s Law is much more useful if you know lots of reactions. Made a table of standard heats of formation. The amount of heat needed to for of a compound from its elements in their Standard states are 1 atm, 1M and 25ºC For an element it is 0 There is a table in Appendix C (pg 1123) Think about this For which of the following reactions at 25 C would the enthalpy change represent the standard enthalpy of formation? For those where it does not, what changes would need to be made in the reaction? 2Na(s) + ½ O2(g) Na2O(s) 2K(l) + Cl2(g) 2KCl (s) C6H12O6(s) 6C(diamond) + 6H2(g) + 3O2(g) One More Try Write the equation corresponding to the standard enthalpy of formation of liquid carbon tetrachloride. Using Enthalpies of Formation to Calculate Enthalpies of Reactions Using Appendix C, we can use Hess s Law to calculate the standard enthalpy change for any reaction for which we know the standard enthalpy of formation values for all reactants and products. Practice Consider the combustion of propane gas, C3H8 Use Hess s law to calculate the standard enthalpy change for this reaction. Breaking it down 11
We obtain the general result that the standard enthalpy change of a reaction is the sum of the standard enthalpies of formation of the products minus the standard enthalpies of formation of the reactants. Let s try it Calculate the standard enthalpy change for the combustion of 1 mole of benzene, C6H6 (l). Calculate the enthalpy change for the combustion of 1 mol of ethanol. The standard enthalpy change for the decomposition of calcium carbonate is 178.1 kj. From the values for the standard enthalpies of formation of CaO(s) and CO2(g), calculate the standard enthalpy of formation of CaCO3(s) Practice Given the following standard enthalpy change, use the standard enthalpies of formation to calculate the standard enthalpy of formation of CuO(s) CuO(s) + H2(g) Cu(s) + H2O(l) H = -129.7 kj 12