Thermochemistry

1. Thermochemistry

2. Energy • Thermochemistry is the study of energy changes and exchanges in chemical systems. • Energy is basically the ability of a system to supply heat or to perform work. • You already know the 2 principal classes of energy. • Kinetic • Potential

3. Kinetic Energy • Kinetic Energy, the energy of motion. • When atoms or molecules move, their mass and speed give them energy:

4. Potential Energy • Potential Energy, internal or stored energy. • It may be stored because of position, or it may be stored internally in chemical substances. • A ball or boulder on a hill has potential energy because of its position in space: if it gets pushed, it will roll down the hill, converting potential energy into kinetic energy (and other energies).

5. Chemical Energy • Chemical Energy is actually both potential and kinetic energy. • It is the energy a chemical substance has based on the positions and motions of its atoms and electrons. • When a chemical rxn takes place, new chemical substances are produced which have different energies as they have different positions and motions of electrons, etc.

6. Conservation of Energy & the First Law • You learned the Law of Conservation of Energy, which isalso called the First Law of Thermodynamics: • Energy can’t be created nor destroyed; it can only be converted from one form of energy to another and transferred from one object to another.

7. Conservation of Energy & the First Law • To follow energy changes, we have defined a system and the surroundings. • The contents of a rxn vessel constitute the system. • Everything else is the surroundings. • Note: if a rxn takes place in a solvent, like water, then the solvent is usually classified as part of the surroundings.

8. Conservation of Energy & the First Law • The surroundings either supply energy to the system or absorb energy released by the system. (and of course we can state the same for the system) • This means that the energy change for the system equals the negative of the energy change of the surroundings: • As we are following the rxn (system), when we say energy or E, we typically mean the Esys

9. The First Law of Thermodynamics • The first law is stated in 2 ways: • The energy of the Universe is constant. • The energy of an isolated system (there is NO energy transfer with any surroundings) is constant. • So there is no free lunch in the Universe! • If someone tries to sell you a product that makes energy out of nothing, don’t buy it!

10. Units of Energy • Scientists use the energy unit joule, J. • The J is an abbreviated unit, here’s the units that you would get from the above equation: KE = (kg)(m2)/s2 • We still use the old unit calorie, cal, where: 1 cal = 4.184J (exact) • If you look at your food label and it says 140 Cal, this is a food calorie, Cal, where 1 Cal = 1000 cal. • The typical male is supposed to eat 2000 Cal/day. How many J is this?

11. Internal Energy • The internal energy of a chemical system is the sum of all of the kinetic and potential energies of all of the particles in the system. E = KE + PE • As stated earlier, it is the ability to produce heat and work.

12. Internal Energy • Unfortunately, it is basically impossible to measure the actual internal energy of a system. • What we can measure is the change in energy of a system as it undergoes a chemical rxn. ΔE = Heat + Work = q + w • If the system is open to atmospheric pressure (as many rxns are), then q is called qp • If the system is closed and has a constant volume, then q is called qv • We will typically do problems which are at constant pressure, so qp is used.

13. Internal Energy • Work in a chemical system is: w = -PΔV • The negative sign reflects the standard terminology that work produced is a negative quantity. • For work produced, the volume increases so ΔVis positive.

14. Internal Energy • Work produced by a chemical system is typically small so the following assumption is made: • Although this is an approximation, it is generally (but not always) within 1%.

15. Energy and Enthalpy • Chemists also have defined another energy term, enthalpy, H, or the enthalpy change,ΔH. • ΔH is the heat energy change, or the enthalpy change of a constant pressure system. • SoΔH = qp • ΔH is also just called q • You will use both terms!

16. Enthalpy andΔH • Enthalpy, H, andΔH are state functions as are E andΔE. • What does this mean? • State functions are independent of path, or how the system arrives at its final state. • What’s another state function that you know?

17. Enthalpy andΔH • We can find the change in enthalpy,ΔH, for a chemical reaction or for a chemical process:

18. Enthalpy andΔH • IfΔH is positive, then heat energy was absorbed by the system. This is an endothermic process or rxn. • IfΔH is negative, then heat energy was released by the system. This is an exothermic process or rxn. • Since we want to findΔH for a rxn or process, we need to know the individualΔH values for all the products and reactants. (more on this later)

19. Meaning ofΔH • Let’s look at a rxn: • Look at the ways we can showΔH. • What does thisΔH value mean? • As the units are kJ/mol rxn, it is a conversion factor, which can take us between mol of a reactant or product and heat energy required or released! • What if we write the reverse reaction?

20. Stoichiometry andΔH • For the above rxn, 2043 kJ of heat energy is released for every mol of propane which is burned. • So how much heat energy would be released if 3.5 moles of propane were burned? • How much heat energy would be released if 25.0 g of carbon dioxide was produced?

21. Calorimetry and Enthalpy Changes • One common experimental method to find the q orΔH for a rxn is to conduct the rxn inside a calorimeter. • Calorimeters may be constant pressure or constant volume. • In the lab, you will use a “coffee cup” calorimeter, which is constant pressure. • Another common type is a “bomb” calorimeter, which is constant volume.

24. Calorimetry and Enthalpy Changes • Whatever type of calorimeter is used, the temperature change of the system, or the surrounding water reservoir, is measured. • This gives usΔT, whereΔT = Tf - Ti • But how can we get fromΔT toΔH?

25. Calorimetry and Enthalpy Changes • A property called the heat capacity of a chemical (or a mixture) lets us make this conversion. • Heat capacity is a measure of a substance’s (or a mixture’s) ability to store heat. • The higher the heat capacity of an object, the more heat energy it can store without its temperature changing.

26. Calorimetry and Enthalpy Changes • We have tables of heat capacities, given in 2 forms: • Specific Heat Capacity, s or cp, which is defined as the amount of heat necessary to raise exactly 1 g of a substance by exactly 1°C. The units are J/g•°C. • Molar Heat Capacity, C or Cm, which is the amount of heat necessary to raise exactly 1 mol of a substance by exactly 1°C. The units are J/mol•°C.

28. Calorimetry and Enthalpy Changes • So the heat capacity, the amount of a substance and theΔT for a rxn can let us calculate theΔH for a rxn,ΔHrxn • What kind of substances have high heat capacities? • Water has one of the highest heat capacities, much higher than most common substances. It’s specific heat capacity is 4.184 J/g•°C. • Water can store a lot of heat energy, and this is crucial for life on our planet. • As our planet is a liquid water-based planet, water is a heat sink or heat reservoir. • So our oceans and large lakes moderate temperature fluctuations on our planet, keeping it from getting too hot during the day and too cold at night.

29. Calorimetry and Enthalpy Changes • Here’s the equations and some problems!

30. Hess’s Law: Adding Rxns Together • If we want to find theΔH for a rxn, and we haveΔH values for other rxns, sometimes we can use Hess’s Law to calculate the desiredΔH from the givenΔH values. • Hess’s Law: the overall enthalpy change is equal to the sum of the individual rxns which make up the rxn. • For example, what if we want to find theΔHrxnfor the following rxn?

31. Hess’s Law: Adding Rxns Together • Can you see how to manipulate Rxn 2) and 3) in order to get Rxn 1)? • Do you add them, add the reverse of one, multiply or divide them by some whole number, etc? • In this case, addition of Rxn 2) and 2 times Rxn 3) gives you Rxn 1) • So what’s theΔHrxnfor Rxn 1)? It’s the sum of theΔHrxnfor Rxn 2) and 2 x 3).

32. Standard Heats of Formation andΔHrxn • Earlier, you learned the following: ΔH°rxn =ΔH°products -ΔH°reactants • This means that we need to know theΔH values for all of the reactants and products in order to calculateΔHrxn • Where can we find these values?

33. Standard Heats of Formation andΔHrxn • There are Tables and Books of Tables which list the thermodynamic values forΔH for thousands of compounds. • To avoid confusion, these values are listed as ΔHf°, or the Standard State Enthalpy Changes. • What’s Standard State? • Standard State is defined as 1 atm pressure (now it’s 1 bar); 1 M for all solutions; and at a specified temperature, which is usually 25°C.

34. Standard Heats of Formation,ΔH°f • But what do theseΔH°f values mean? • They are the standard heats of formation for a substance. • They are the enthalpy change when exactly 1 mol of the substance is made from its elements under standard state conditions. • We can write equations which show exactly what we mean.

35. ΔH° Tables

36. Standard Heats of Formation,ΔH°f • How would you make propane, C3H8, from the elements? • You make it from the most stable elemental form of the element. • Here’s the heat of formation equation for propane: • Note that graphite is the stable elemental form of carbon.

37. Standard Heats of Formation,ΔH°f • What’s the heat of formation equation for liquid water? • Note in the above that we are ALLOWED (actually we are MANDATED) to have fractions in the rxn equation! • Why? Because of the definition of aΔH°f : exactly 1 mol of the substance is made!

38. Standard Heats of Formation,ΔH°f • Now we can useΔH°rxn =ΔH°products -ΔH°reactants • There really is a more mathematically correct way to write this equation:

39. Standard Heats of Formation,ΔH°f • So we just look up theΔH°f for all of the products and reactants and add and subtract them together. • Example: Using Tables ofΔH°f values, find theΔH°rxn for the combustion of propane: • What is interesting is that theΔH°f for the stable elemental state of an element is 0. • In problems, you will not usually be given theΔH°f value for stable elemental forms, you are expected to know that it is zero! (Remember this!)

40. Heat of Combustion,ΔH°comb • You noticed that in Heat of Formation equations, you could legally have fractions in the rxn equation. • There is another common case where it is legal to have fractions in the rxn equation: Heat of Combustion Rxn Equations. • The Heat of Combustion,ΔH°comb (or justΔH°c) is defined is the heat energy change when exactly 1 mol of a substance is combusted with oxygen gas. • We are allowed to have fractions in combustion rxn equations if we are calculating heats of combustion.

41. Heat of Combustion,ΔH°comb • Write the combustion rxn equation for benzene, C6H6, and calculate itsΔH°comb. • Of course, combustions rxns are very important to us as they heat our homes, power our cars, and light our homes (as most electricity is produced by burning fossil fuels). • We can useΔH°combvalues for the reactants and products in a chemical reaction to findΔH°rxn • However, as combustion is for breaking apart a compound, while heats of formation were for making a compound, we have to change the signs of theΔH°combvalues before we use them in products - reactants.

43. Enthalpy Changes for Phase Changes • There are 6 changes of state that a chemical may undergo: • Vaporization • Condensation • Sublimation • Melting or Fusing • Freezing • Deposition

44. Enthalpy Changes for Phase Changes • There is aΔH associated with all of these phase changes. • ΔHvapis the enthalpy change associated with the vaporization process or the heat of vaporization;ΔHfis the heat of fusion; andΔHsubis the heat of sublimation. • As condensation if the reverse of vaporization, we don’t have a special term for the heat of condensation, it’s just –ΔHvap • 3 of the phase changes are exothermic: which?

45. Enthalpy Changes for Phase Changes • We can also combine 2 phase changes! • When we sublime something, it goes from the solid state to the gas state directly. • But asΔH is independent of path (it’s a State Function), we could first melt the solid and then vaporize it to the gas. TheΔH value for the 2 paths would be the same, so long as the temperature was held constant. • So for constant T,ΔHsub=ΔHvap+ΔHf