Thermochemistry

1. Thermochemistry -a study of energy

2. Look at the world around us. The world we have built depends on our use of energy. The cars we drive, the electricity that powers our lights and machines, the dynamite we use in construction, and the way we heat our homes, all depend on our efficient use of energy. The way we use this energy is therefore very important to us. • Two aspects of energy use should be considered: • a) Sources of energy we use.
b) The technologies we have created to use the sources of energy efficiently.

3. Nitroglycerin C3H5(NO3)3(l) gives off 1500 KJ of energy per mole when it decomposes. • Canadian Minute • It is the speed of the decomposition reaction which makes nitroglycerin such a violent explosive. Unlike burning, which can only travel as fast as the flame front can move through the material, high explosives are decomposed almost instantaneously by a supersonic shock wave passing through the material. This instantaneous destruction of all the molecules in the sample is called a detonation, and the rapid expansion of hot gases that results is what causes the destructive blast. In fact, 4 moles of nitroglycerin produces 35 moles of hot gases.

4. Nitroglycerin gives off 1500 KJ of energy per mole when it decomposes. • Alfred Nobel found a way to stabilize nitroglycerin when he invented dynamite.

5. Energy changes accompany every chemical and physical process from the decomposition of nitroglycerin to the vaporization of water. We will explore the energy changes that accompany various processes.

6. Energy is defined as the ability to do work (we are really talking about the movement that occurs against a restraining force and is equal to the force x the distance over which the motion occurs) or produce heat. • Energy comes in many forms- solar, nuclear, and electrical are just a few examples. In chemistry, heat energy is often what we are interested in. • In previous science classes you have most likely learned that the individual atoms and molecules that make up all substances are in motion. This motion is called kinetic energy. The energy associated with this motion is heat, or thermal energy.

7. Heat can also be defined as the flow of energy. This energy transfer is always from the hotter substance to the colder object. If you hold an ice cube in your hand, your hand feels colder. Your hand is giving heat (or losing heat) to the ice cube. The ice cube does not transfer “coldness” to your hand.

8. In addition to kinetic energy there is another basic form of energy- potential energy. Potential energy is energy of position (stored energy).

9. We will be concerned with one type of potential energy in this unit – chemicalbond energy. • Chemical bond energy comes from the attractive forces between molecules, and between atoms. This form of energy plays an important role in chemical reactions.

10. The standard unit for energy is the joule (J). Another standard unit for energy is the calorie. Calories (with a capital C) are used to measure the energy in food; 1 Calorie (capital C) = 1000 calories (lowercase c). . One joule is the amount of energy used by a 100 W light bulb in 0.1 s. . One joule is approximately equal to the amount of energy you expend when you bring a cheeseburger to your mouth. . An ordinary paper match releases a little over 1,000 J when completely burned. Temperature describes the amount of motion that the molecules or atoms in a material have. Fast movement represents high temperature and slow movement represents low temperature. Temperature then is a measure of the average kinetic energy of the particles that make up the substance. Kelvin = oC + 273.15 oC = K – 273.15

11. Thermodynamics and Thermochemistry • Thermodynamics is the study of heat and how it can be converted into other energy forms. In thermodynamics, the system is a specific part, often the system is a chemical reaction. The surrounding are typically everything outside the test tube where the chemical reaction is taking place.

12. An open system can exchange both energy and matter with its surroundings. An active volcano is an example of an open system. • A closed system can exchange energy but not matter. A sealed greenhouse is an example of a closed system. • An isolated system can exchange neither energy nor matter. A sealed thermos bottle is an isolated system.

13. The First Law of Thermodynamics also known as Law of Conservation of Energy The first law states that energy cannot be created or destroyed. In other words, when a system gains or loses energy from the surroundings, the total energy will be constant. . When octane (C3H8, the main component of gasoline) is burned, chemical bond energy (potential energy) is converted into mechanical energy (pistons moving in the car); kinetic energy) and heat. . When we eat, our bodies convert the chemical energy of the food into movement of our muscles; heat is also produced as a product of this conversion.

14. This concept can be expressed as: ΔE = q + w The E represents the internal energy. This includes but is not limited to all kinetic energy and potential energy possessed by all components of a system. ΔE is the change in internal energy of a system as a result of heat flow and work. Heat flow is indicated by q, and the work done by the system is w. . q is a positive value if heat is absorbed by the system. . q is a negative value if heat is lost by the system. . w is a positive value if work is done on the system. . w is a negative value if work is done by the system. "In this house we obey the laws of thermodynamics!" (after Lisa constructs a perpetual motion machine whose energy increases with time) — Homer Simpson

15. The Second Law of Thermodynamics states that heat energy always travels spontaneously from a warmer body (the body with the higher temperature) to a colder body. • You have known since you were a child that if you touch a hot stove you will get burned. The stove is warmer than your body and therefore transfer heat to your hand!

16. If you go outside in the very cold without gloves you can get frostbite. The heat is moving away from your body into the surroundings. When your body can no longer replace the heat as quickly as it is being lost, your skin freezes and you get frostbite.

17. Endothermic and Exothermic Reactions All chemical reactions (and physical changes and nuclear changes as well) involve energy changes. Exothermic reactions are reactions in which there is a net release of energy. When energy is released, an energy term will appear on the product side of the equation. Fe2O3(s) + 2 Al (s) Al2O3(s) + 2Fe(l) + 847.6 kJ

18. Endothermic reactions are reactions that require a net input of energy. This is indicated by writing the energy term on the reactant side of the equation. 2S03(g) + 198 kJ  2 SO2(g) + O2(g)

19. Measuring Heat: Calorimetry An instrument called a calorimeter measures the heat which is exchanged during a reaction.

20. A calorimeter is a thermally insulated container where a reaction system is contained and the energy exchange between the system and its surroundings can be measured. • If the reaction is exothermic, energy is given off causing the temperature of the water in the calorimeter to rise. • If the reaction is endothermic, energy will be required. This causes the temperature of the water to drop. • By noting the temperature change of the water in the calorimeter, we can calculate how much energy was exchanged.

21. The specific heat constant (c) of a substance is defined as the energy required to raise the temperature of one gram of a substance by one degree Celsius. • Every substance has a characteristic specific heat constant. The specific heat constant of water is 4.19 J/goC. • The First Law of Thermodynamics says that the heat energy of the system is equal to the negative heat energy of the surroundings. • q surroundings = -q system • The heat, q entering or exiting can be determined by using the heat equation for the surroundings.

22. q(heat)=(mass) (specific heat)(final temperature – initial temperature) • q = mcΔt • Water has a relatively high specific heat which allows it to absorb and release large quantities of heat. This is why temperature changes are often more moderate in areas near large lakes and other large bodies of water. There is a reason why people buy cottages other than boating!

23. Q= mcΔT Where Q = amount of heat transferred, in kJ m = mass of the sample in grams ΔT = change in temperature, in oC Note: We indicate an increase in temperature as a positive change and a decrease in temperature as a negative change. If the temperature started at 50 oC and ended at 40 oC, ΔT would have a value of – 10oC. The Calculations

24. A sample of 100 grams of water is placed in a calorimeter at a temperature of 20.0 oC. Ammonium nitrate is dissolved in the water lowering the temperature to 10.0 oC. Calculate the heat change for the water. Specific heat of water = 4.19 J/goC. • Q = mcΔT = 100 g 4.19 J/goC -10oC = - 4190 J The qwater is negative indicating that the heat flows from the surroundings toward the solution. The solution is an endothermic solution ΔHsolution = +4190 J or + 4.190 kJ

25. A 1000.0 g mass of water had a starting temperature of 50oC. It lost 33,600 J of heat over a 5 minute period. What was the temperature of the water at the end of this period? • Q = mcΔT so • ΔT = Q mc = 33,600 J 1000.0 g 4.184 J/goC = 8 oC The temperature decreased 8oC so the fianl temperature was 50 – 8 = 42oC.

26. There are energy changes during a physical change. They are changes of state or phase. During the heating of a solid, liquid or a gas, there is a temperature change. During the state change itself, the temperature remains constant. These heats are hidden (latent). • The latent heat of fusion is the energy change that occurs during the phase change between solid and liquid. The latent heat of vaporization is the energy change that occurs during the phase change between a liquid and a gas. • Molar heat of fusion is the amount of energy required to melt one mole of substance. • Molar heat of vaporization is the amount of energy required to vaporize one mole of substance. • Fusion (melting) the substance absorbs heat, ΔHfus is positive • Freezing the substance releases energy, ΔHsolid is negative • ΔHfus = -ΔHsolid • Vaporization the substance absorbs energy, ΔHvap is positive • Condensation the substance releases energy, ΔH cond is negative • ΔHvap = -ΔHcond

27. To calculate the energy for the phase change you need multiply the number of moles (n) by the molar heat for the phase change: • q = nΔHphase

28. Potential Energy Diagram of a Physical Change

29. Examples of Energy Diagrams 1. Potential Energy Profile for a Phase Change NH4NO3(s)  NH4NO3 (aq) ΔH = +25 kJ

30. NaOH(S)NaOH(aq) ΔH = -44.2 kJ

31. An equation that indicates that the heat change is a thermochemical equation. • NH4NO3(s) + 25 kJ  NH4NO3(aq) • NaOH(s)NaOH(aq) + 44.2 kJ • Another way to show heat change is to use the ΔH notation, the heat change for the system is shown outside the equation. A positive ΔH represents an endothermic process (system increases in energy). A negative ΔH represents an exothermic process (system decreases in energy). • Enthalpy or ΔH is written in terms of the heat change of the system.

32. Enthalpy, H is a state function used to describe the heat changes that occur in a reaction under constant pressure. • When a reaction is allowed to take place in an open container, a quantity of heat proportional to the quantity of matter present, will be released or absorbed. • This flow of heat is the enthalpy change, ΔH. The units for ΔH are kJ or kJ/mol. • Reactions that release heat are termed exothermic. They have negative values of ΔH. • Reactions that absorb heat are termed endothermic. They have positive values of ΔH. • Generally processes that feel cold (like an ice pack) are endothermic and processes that feel hot (like a fire) are exothermic. Enthalpy

33. Enthalpy Changes (Energy changes during phase, chemical and nuclear change) • When a phase, chemical or nuclear change takes place not only does the kinetic energy of the system change but so does the potential energies of the molecules. • Chemists therefore consider the total energy of the system when calculating these energy changes. Chemists define the total internal energy of a system at constant pressure as Enthalpy . • Enthalpy is given the symbol H. • Chemists have no method of measuring the total energy of a system easily, but they can measure the changes that occur in this energy. They do this by examining what happens to the surroundings of the system. These changes are termed Δ H or enthalpy change. • SATP The ΔH of a reaction changes with varying conditions of temperature and pressure, so chemists define a set of conditions called standard atmospheric temperature and pressure (SATP). Under these conditions they measure the value of H. The standard atmospheric temperature is 25 oC and the standard pressure is 100 KPa.

34. Comparison of Enthalpy change • The enthalpy change for phase changes are typically smaller than those for chemical changes. Enthalpy changes for nuclear changes are typically very large compared to chemical changes. See below. • Phase change ( dissolving ) NH4NO3 (s)NH4+ (aq)+NO3 - (aq) H=+ 27 kJ Chemical change ( decomposition ) • 2 Fe2O3 (s)4 Fe (s)+3 O2 (g) H=1625 kJ • Representing (communicating) Enthalpy Changes in Reactions Chemists can represent the change in enthalpy using several methods: 1. Using a graph of Enthalpy change vs the course of the reaction. 
2. Using the balanced equation and enthalpy as a term in the equation. 
3. Using the balanced equation and enthalpy written separately as a "H".

35. 1. Graphical Graph of Endothermic Change for the decomposition of water H H2(g)- ½ O2 (g) ΔH = +285.8 KJ H2O (l) Course of Reaction

36. 2) Using a balanced equation and energy as a term. There is a direct relationship between the amount of substances that reacts and forms in an equation. The factor that determines the exact relationship is the equation factor or mole ratio (balance in equation). For example in the reaction below: 2H2 (g) + 1 O2 (g) ------> 2 H2O (l) If one mole of hydrogen reacts with excess oxygen, one mole of water forms. If two moles of oxygen reacts with excess hydrogen, 4 moles of water forms. What relationship exists between the amount of reactant or product substance and the change in enthalpy. (ΔH )? Complete the dry lab assignment to determine the answer to this question. The heat absorbed and produced in a chemical reaction also varies directly as the amount of substance that reacts and the exact amount is determined by the heat change for the reaction (ΔH) . If you double the amount of substance reacted then you will double the heat change. Examine the following reaction between Hydrogen and oxygen to form water 2H2 (g) + 1 O2 (g) -----> 2H2O (l) If one mole of oxygen when reacted releases 572 kJ of energy , then 2 moles will release 1046 kJ Because the relationship is a direct one, energy can be written as a term in the equation. If the reaction is exothermic the energy is written the products. If the reaction is endothermic the energy is written in the reactants.

37. 3) " ΔH" notation The third method used by chemists to represent enthalpy change is called " H" notation. In this case the enthalpy is written separately from the reaction. If the reaction is exothermic the ΔH is negative. If the reaction is endothermic the ΔH is positive. Using the equations above: Exothermic: (-H value) 2Al (s) + 3Cl2 (g) ------> 2 AlCl3 (s) 
 Δ H = - 1408 kJb) Endothermic: (+H value) 2 H2O ------> 2 H2 + 1 O2 (g) 
 Δ H = + 572 kJ What would happen if we cut the # of moles of reactant and product in half in each of the above reactions? Enthalpy change would be cut in half. Complete the assignment “Communicating Enthalpies”

38. The heat absorbed and produced in a chemical reaction also varies directly as the amount of substance that reacts and the exact amount is determined by the heat change for the reaction (H) . If you double the amount of substance reacted then you will double the heat change. • Examine the following reaction between Hydrogen and oxygen to form water • 2H2 (g) + 1 O2 (g) -----> 2H2O (l) • If one mole of oxygen when reacted releases 572 kJ of energy , then 2 moles will release 1046 kJ • Because the relationship is a direct one, energy can be written as a term in the equation. • If the reaction is exothermic the energy is written the products. 
If the reaction is endothermic the energy is written in the reactants.

39. Example: a) Exothermic: written in the products. 2Al (s) + 3Cl2 (g) ----> 2 AlCl3 (s) + 1408 kJ • If we double the amount of the reactants and products, the heat doubles 4Al (s) + 6Cl2 (g) ----> 4 AlCl3 (s) + 2816 kJ b) Endothermic: written in the reactants. H2O + 286 kJ ------> H2 + 1/2 O2 (g) If we double the amount of the reactants and products we double the heat. 2 H2O + 572 kJ ------> 2 H2 + 1 O2 (g)

42. Energy Profile for a Chemical Change To keep comparisons fair the chemist will measure heat changes under standard conditions (SATP 25 ºC and 101kPa). The reactants and products are at SATP. To show that the ΔH was measured at SATP we use the symbol ΔHº. CaO(s) + H2O(l) → Ca(OH)2(s) + 65.2kJ (exothermic) Another way to show heat change is to use the ΔH notation, the heat change for the system is shown outside the equation. A negative ΔH represents an exothermic process (system decreases in energy). CaO(s) + H2O(l) → Ca(OH)2(s)  ΔHº = -65.2kJ/mol

43. This can be represented by the energy profile:

44. 2NaHCO3 + 129kJ → Na2CO3(s) +H2O(g) + CO2(g)  (endothermic)  • 2NaHCO3 → Na2CO3(s) + H2O(g) + CO2 (g)  ΔHº = +129kJ • A positive ΔH represents an endothermic process (system increases in energy). This reaction can be represented by the energy profile:

45. The endothermic reaction shown below indicates that 92.2 kJ are absorbed when 2 moles of NH3 decompose to form 1 mole of N2 and 3 moles of H2. 2NH3 (g)  N2 (g) + 3H2 (g) ΔH = + 92.2 kJ or 92.2 kJ + 2NH3 (g)  N2 (g) + 3H2 (g) Question : In the reaction above, 100 g of NH3 are allowed to react to produce N2 and H2, how many kJ of heat will be absorbed? Answer: The reaction tells us that for every 2 moles of NH3 consumed, 92.9 kJ will be absorbed. This information will be used to calculate the conversion from moles of NH3 to kJ. 100 g NH3 x1 mol NH3x92.2 kJ = 271 kJ 17 g NH3 2 mol NH3 When a reaction is reversed, the sign of ΔH is changed: N2 (g) + 3H2 (g)  2NH3 (g) ΔH = -92.2 kJ Enthalpy Calculations

46. Hess’s Law states that the enthalpies of reactions may be added when these reactions are added. Substances appearing on the same side are added, while those on opposite sides are subtracted. Typically, some reactions will need to be reversed and multiplied through by a number, so that when they are combined the desired equation will result. Hess’s Law

47. In math we can add equations because we are adding the same to each side of the equation. Common terms on the opposite sides of the equations can be cancelled out before we add the equations (b and c terms can be cancelled) a + b = c + d a + c = b + d 2a    =      2d 

48. The heat of reaction #3 could be found using a colorimeter or it can be found indirectly by calculating the sum of Reaction #1 and Reaction #2. The additive property of chemical reactions is called Hess’s Law.

49. Find the enthalpy of the equation: • C2H2 (g) + 5