Internal Energy (U) of a system is the total energy contained within the system, partly as kinetic energy and partly as potential energy

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1 אנרגיה פנימית Internal Energy Internal Energy (U) of a system is the total energy contained within the system, partly as kinetic energy and partly as potential energy Kinetic energy involves three types of molecular motion: Translation העתקה Rotation סיבוב Vibration ויבראציה Collectively, these are sometimes called thermal energy 1 אנרגיה פנימית Internal Energy Internal Energy (U) of a system is the total energy contained within the system, partly as kinetic energy and partly as potential energy Potential energy involves intramolecular interactions (i.e., bonds): and intermolecular interactions: Intramolecular forces כוחות תוך-מולקולריים Intermolecular forces כוחות בין-מולקולריים Collectively, these are sometimes called chemical energy 2 1

2 Thermochemistry Thermochemistry is the study of energy changes that occur during chemical reactions Universe Surroundings System Surroundings Surroundings Focus is on heat and matter transfer between the מערכת system and the סביבה surroundings 3 Thermochemistry There are three types of systems in thermochemistry: Matter Matter Matter Matter Matter Matter Energy Energy Energy Energy Energy Energy CLOSED מערכת סגורה OPEN מערכת פתוחה ISOLATED מערכת מבודדת 4 2

3 Energy Transfer Mechanisms Energy can be transferred between the system and its surroundings as: Heat: the reaction in the system changes the temperature of the surroundings Work: the reaction in the system causes work to be done (i.e., force is moving through a distance) Different types of work: expansion/compression, electrical, etc. There is no such thing as negative energy nor positive energy ; the sign of work (or heat) signifies the direction of energy flow. 5 חום (q) Heat Heat is energy transfer resulting from thermal differences between the system and surroundings Heat flows spontaneously from higher T lower T Heat flow stops at thermal equilibrium 6 3

4 Heat Transfer Illustrated 7 Heat Transfer Mechanism Illustrated Inelastic molecular collisions are responsible for heat transfer: 2. transfer energy to less energetic molecules. 1. More energetic molecules Heat transfer demo 8 4

5 עבודה (w) Work Work is an energy transfer between a system and its surroundings Compression: work done ON the system System gains energy (+w) compression expansion Expansion: work done BY the system System loses energy (-w) 9 Pressure-Volume Work For now we will consider only pressure-volume work. w = PΔV 10 5

6 פונקציות מצב State Functions The state of a system: its exact condition at a fixed instant. State is determined by the kinds and amounts of matter present, the structure of this matter at the molecular level, and the prevailing pressure and temperature. A state function is a function whose value depend only the present state of a system, and does not depend on how the state was reached (i.e., does not depend on the history of the system). 11 First Law of Thermodynamics Law of Conservation of Energy in a physical or chemical change, energy can be exchanged between a system and its surroundings, but no energy can be created חוק שי מור ה אנרגיה destroyed. or The internal energy change of a system is simply the difference between its final and initial states: ΔU = U final U initial If energy change occurs only as heat (q) and/or work (w), then: ΔU = q + w 12 6

7 First Law: Sign Convention Think from the point of view of the system: Energy entering a system carries a positive sign: heat absorbed by the system (q > 0), or work done on the system (w > 0) Energy leaving a system carries a negative sign heat released by the system (q < 0) work done by the system (w < 0) 13 חום תגובה ) rxn Heats of Reaction (q q rxn is the quantity of heat exchanged between a reaction system and its surroundings. An exothermic reaction releases heat In isolated systems, system T. The system goes from higher to lower energy; q rxn <0. An endothermic reaction absorbs heat In isolated systems, system T. The system goes from lower to higher energy; q rxn >

8 Conceptualizing an Exothermic Reaction Surroundings are at 25 C 25 C 2. Typical situation: some heat is released to the surroundings, some heat is absorbed by the solution C 35.4 C 1. Hypothetical situation: all heat is instantly released to the surroundings. Heat = q rxn 3. In an isolated system, all heat is absorbed by the solution. Maximum temperature rise. 15 For a system where the reaction is carried out at constant pressure, ΔU = q P PΔV or ΔU + PΔV = q P Most of the thermal energy is released as heat. Some work is done to expand the system against the surroundings (push back the atmosphere). Internal Energy Change at Constant Pressure 16 8

9 Enthalpy and Enthalpy Change Enthalpy is defined as: H = U + PV Enthalpy change is thus: ΔH = ΔU + Δ(PV) For a process carried out under a constant pressure: ΔH = ΔU + PΔV ΔH = q p For a process carried out under a constant pressure and in which the volume does not change ( no work is done): ΔH = ΔU The evolved H 2 pushes back the atmosphere; work is done at constant pressure. Mg + 2 HCl MgCl 2 + H 2 Most reactions occur at constant pressure, so for most reactions, the heat evolved equals the enthalpy change. 17 Properties of Enthalpy Enthalpy is an extensive property. It depends on how much of the substance is present. Since U, P, and V are all state functions, enthalpy H must be a state function also. Enthalpy changes are unique for each reaction. Two logs on a fire give off twice as much heat as does one log. Enthalpy change depends only on the initial and final states. In a chemical reaction we call the initial state the and the final state the. 18 9

10 Enthalpy Diagrams Values of ΔH are measured experimentally. ΔH < 0 exothermic reactions. ΔH > 0 endothermic reactions. A decrease in enthalpy during the reaction; ΔH < 0. An increase in enthalpy during the reaction; ΔH > Reversing a Reaction ΔH changes sign when a process is reversed. Therefore, a cyclic process has the value ΔH = 0. Same magnitude; different signs

11 Given the equation (a) H 2 (g) + I 2 (s) 2 HI(g) ΔH = kj calculate ΔH for the reaction (b) HI(g) ½H 2 (g) + ½ I 2 (s). The complete combustion of liquid octane, C 8 H 18, to produce gaseous carbon dioxide and liquid water at 25 C and at a constant pressure gives off 47.9 kj of heat per gram of octane. Write a chemical equation to represent this information. 21 ΔH in Stoichiometric Calculations For problem-solving, heat evolved (exothermic reaction) can be thought of as a product. Heat absorbed (endothermic reaction) can be thought of as a reactant. We can generate conversion factors involving ΔH. For example, the reaction: H 2 (g) + Cl 2 (g) 2 HCl(g) ΔH = kj can be used to write: kj kj 1 mol H 2 1 mol Cl kj 2 mol HCl 22 11

12 What is the enthalpy change associated with the formation of 5.67 mol HCl(g) in this reaction? H 2 (g) + Cl 2 (g) 2 HCl(g) ΔH = kj 23 Hess s Law of Constant Heat Summation The heat of a reaction is constant, regardless of the number of steps in the process ΔH overall = sum (ΔH s of individual reactions) When it is necessary to reverse a chemical equation, change the sign of ΔH for that reaction When multiplying equation coefficients, multiply values of ΔH for that reaction 24 12

13 Hess s Law: An Enthalpy Diagram We can find ΔH (a) by subtracting ΔH (b) from ΔH (c) 25 Calculate the enthalpy change for reaction (a) given the data in equations (b), (c), and (d). (a) 2 C(graphite) + 2 H 2 (g) C 2 H 4 (g) ΔH =? (b) C(graphite) + O 2 (g) CO 2 (g) ΔH = kj (c) C 2 H 4 (g) + 3 O 2 2 CO 2 (g) + 2 H 2 O(l) ΔH = kj (d) H 2 (g) + ½ O 2 H 2 O(l) ΔH = kj 26 13

14 Calculating Enthalpy Changes We would like to calculate the change in enthalpy in a reaction like that: ΔH = H products H reactants Problem: we can only measure enthalpy changes (i.e., we can t obtain absolute values of enthalpy of compounds). ΔH = H products H reactants What could be a common reference? = H products H ref. H reactants + H ref. = (H products H ref. ) (H reactants H ref. ) 1. ΔH of a certain reaction that forms the products from reference molecules 2. ΔH of a certain reaction that forms the reactants from the same reference molecules 27 Standard Enthalpies of Formation The common reference: the elements from which both reactants and products are made! We define the standard state of a substance as the state of the pure substance at 1 atm pressure and the temperature of interest (usually 25 C). The standard enthalpy change (ΔH ) for a reaction is the enthalpy change in which reactants and products are in their standard states. The standard enthalpy of formation (ΔH f ) of a substance is the enthalpy change of forming 1 mol of a substance from its אנתלפי ית היצ י רה states. component elements in their standard ΔH f of a pure element =

15 Standard Enthalpies of Formation at 25 o C 29 Using standard enthalpies of formation to calculate ΔH 0 of a reaction Example: an exothermic reaction: elements ΔH f0 = 0 enthalpy 1. ΔH f0 (reactants) is known reactants 3. ΔH 0 can be calculated! 2. ΔH f0 (products) is known products 30 15

16 ΔH Calculations Based on Standard Enthalpies of Formation 0 0 [ ( )] [ 0 rxn = cp ΔH f products cr ΔH f ( reactants) ] products reactants Summation over all products Stoichiometric coefficients of products Summation over all reactants Stoichiometric coefficients of reactants In other words 1. Add all of the values for ΔH f of the products (multiplied by the corresponding stoichiometric coefficients). 2. Add all of the values for ΔH f of the reactants (multiplied by the corresponding stoichiometric coefficients). 3. Subtract #2 from #1 31 Synthesis gas is a mixture of carbon monoxide and hydrogen that is used to synthesize a variety of organic compounds. One reaction for producing synthesis gas is 3 CH 4 (g) + 2 H 2 O(l) + CO 2 (g) 4 CO(g) + 8 H 2 (g) ΔH =? Use standard enthalpies of formation to calculate the standard enthalpy change for this reaction

17 The combustion of isopropyl alcohol, common rubbing alcohol, is represented by the equation 2 (CH 3 ) 2 CHOH(l) + 9 O 2 (g) 6 CO 2 (g) + 8 H 2 O(l) ΔH = 4011 kj Use this equation and data from Table 6.2 to establish the standard enthalpy of formation for isopropyl alcohol. Without performing a calculation, determine which of these two substances should yield the greater quantity of heat per mole upon complete combustion: ethane, C 2 H 6 (g), or ethanol, CH 3 CH 2 OH(l). 33 Important Exhothermic Reactions Combustion of fossil fuels: Coal: C(s) + O 2 (g) CO 2 (g) High-grade coal yields 30 kj/gram of coal Natural gas: CH O 2 (g) CO 2 (g) + 2 H 2 O(l) ΔH 0 = kj/mol 55.5 kj/g Petroleum: C 8 H 18 (l) + 25 / 2 O 2 (g) 8 CO 2 (g) + 9 H 2 O(l) ΔH 0 = 5450 kj/mol 52.4 kj/g Food: Fuels for the body Carbohydrates (starches and sugars), fats, and proteins During digestion, carbohydrates are converted into the simple sugar glucose (C 6 H 12 O 6 ) C 6 H 12 O 6 (s) + 6 O 2 (g) 6 CO 2 (g) + 6 H 2 O (l) ΔH 0 = 2803 kj/mol 15.6 kj/g 34 17

18 Summary of Concepts Thermochemistry concerns energy changes in physical processes or chemical reactions Thermochemistry includes the notion of a system and its surroundings; the concepts of kinetic energy, potential energy, and internal energy; and the distinction between two types of energy exchanges: heat (q) and work (w) Internal energy (U) is a function of state Enthalpy (H) is a function based on internal energy, but modified for use with constant-pressure processes 35 Summary of Concepts The first law of thermodynamics relates the heat and work exchanged between a system and its surroundings to changes in the internal energy of a system The concepts of standard state, a standard enthalpy change, and a standard enthalpy of formation are important in thermochemical calculations Some practical applications of thermochemistry deal with the heats of combustion of fossil fuels and the energy content of foods 36 18