1. Chapter 6: Thermochemistry

2. Energy Energy - the capacity to do work
Work is the result of a force acting over a distance
Types of Energy
Kinetic
Potential
Law of conservation of energy - energy can not be created or destroyed, but it can be transformed

3. Kinetic energy
Kinetic energy - Energy associated with the motion of an object
Mechanical energy – the energy of moving macroscale objects
Thermal energy (heat) – the energy of moving microscale objects

4. Potential energy
Potential energy – energy associated with position or composition, stored energy
Has the potential to be transformed into kinetic energy
Chemical energy – relative positions of electrons and nuclei in atoms/molecules represents potential energy

5. Energy: System vs surroundings
System - A component or set of components of interest
You define the system
Surroundings – everything that a system can exchange energy with, everything touching/connected to it
That which “surrounds” the system

6. Units of energy Joule (J) - basic unit of energy in the metric system
1J =1kg •m2/s2
1kJ = 1000 J
calorie (cal) – 1cal = the energy required to raise the temp of 1g of H2O by 1ºC
4.184 J = 1 cal
1000 cal = 1kcal = 1Cal
Calorie (Cal) – unit associated with “food” calories
Calorie = kcal = food calorie

7. First law of thermodynamics
Thermodynamics – the study of energy and its interconversions
First law of thermodynamics – the total energy of the universe is constant
Law of conservation of energy is built into the definition
Perpetual motion machine (a device that puts out constant energy) cannot exist

8. Internal energy
Internal energy (E) – the sum of the kinetic and potential energies of all particles that compose a system
E is a state function which is only dependent on the state of the system, not how it got there

9. Energy of a system in a reaction
Where does lost energy go or where does gained energy come from? ………. the surroundings!

10. Practice
A chemical reaction occurs!!! The reactants have a total energy of 75 kJ, while the products have a total energy of 45 kJ.
What is the change in internal energy of this system (the compounds involved in the chemical reaction are the system)?
Did the system lose or gain energy?
Did the surroundings lose or gain energy?

11. Heat and work
Heat and work can be exchanged between the surroundings and the system
q = heat, the flow of energy caused by a temperature difference
w = work

12. Heat flow
Thermal equilibrium – heat is distributed to molecules in contact with them until the thermal energy of the system and surroundings is the same temperature

13. Quantifying heat
Substances can transfer heat, but have different capacities to do so
Heat capacity (C) – the quantity of heat required to change a substances temperature by 1 °C

14. Problems
A thief breaks into a safe at a bank to steal a block of gold in the shape of a coin from Super Mario Bros. The safe keeps the gold at 14.0 °C. Once the thief grabs the gold coin in his hand, the coin eventually reaches body temperature at 37.0°C. How much heat was transferred to the gold coin which has a mass of 11.25g?

15. Thermal energy transfer
When two substances of different temperatures reach thermal equilibrium with each other, the heat from the hotter substance is transferred to the colder substance
The hotter substance cools down, as the colder substance warms up, with their final temperatures being the same
𝑞 𝑠𝑦𝑡𝑒𝑚 =− 𝑞 𝑠𝑢𝑟𝑟𝑜𝑢𝑛𝑑𝑖𝑛𝑔
𝑞 𝑚𝑒𝑡𝑎𝑙 =− 𝑞 𝑤𝑎𝑡𝑒𝑟
𝑚 𝑚𝑒𝑡𝑎𝑙 𝐶 𝑚𝑒𝑡𝑎𝑙 Δ 𝑇 𝑚𝑒𝑡𝑎𝑙 =− 𝑚 𝑤𝑎𝑡𝑒𝑟 𝐶 𝑤𝑎𝑡𝑒𝑟 Δ 𝑇 𝑤𝑎𝑡𝑒𝑟

16. Thermal energy transfer problems
With the masses and initial temperatures of the metal and water, what will the final temperature be when they both reach thermal equilibrium?

17. Distribute Kmetal and Kwater
Collect Tfinal terms on the leftside, with other terms going to right side
Factor out Tfinal
Isolate Tfinal

18. Practice
A 32.5 g block of aluminum initially at 45.8 °C is submerged into g of water initially at 15.4 °C. What is the final temperature once the two substances reach thermal equilibrium?

19. Pressure-volume work – when a force causes a volume change against an external pressure
If the system does work to increase the volume against an external pressure, the ΔV is positive
Pressure-volume work

20. Using the equation Units of pressure = atm Units of volume = L
Pressure x volume = L x atm
You can convert units of L x atm to units of J
1 L x atm = J

21. ΔE for chemical reactions
If a system is at constant volume, then there is no work done
If a chemical reaction does no work….
qv = the heat at constant volume

22. Measuring ΔErxn
Calorimetry- measurement of the external energy exchanged between a reaction(system) and the surrounding by monitoring the temperature change
Bomb calorimeter is a device used in calorimetry
The chemical reaction(system) transfers heat to the bomb (surrounding)

23. Enthalpy (H)
Enthalpy – the sum of the internal energy of a system and the product of its pressure and volume
For a process at a constant pressure…..
ΔH is the heat absorbed or evolved by the system
Typically ΔE and ΔH are the same

24. Endothermic vs Exothermic
If the ΔH of a chemical reaction is positive, the system absorbs heat from the surroundings, this is an endothermic reaction
If the ΔH of a chemical reaction is negative, the system releases heat to the surroundings, this is an exothermic reaction

25. How does a chemical reaction give off heat???
Chemical “potential” energy is stored in the bonds
Chemical energy arises from the electrostatic forces between protons and electron that compose atoms and molecules
Breaking bonds absorbs energy, while making bonds releases energy
Breaking stronger bonds absorbs more energy than breaking weaker bonds

26. ΔH and thermochemical equations
ΔHrxn – the enthalpy of reaction or heat of reaction, the heat transferred in a reaction
C3H8(g) + 5O2(g)  3CO2(g) + 4H2O(g) ΔHrxn = kJ
When the reaction occurs with the corresponding moles from the equation, 2044kJ of heat is released from the system (exothermic reaction)
1 mol C3H8: -2044kJ or 5 mol O2: -2044kJ

27. Practice
An LP gas tank in a home barbeque contains 13.2 kg of propane, C3H8. Calculate the heat (in kJ) associated with the complete combustion of all the propane in the tank.
C3H8(g) + 5O2(g)  3CO2(g) + 4H2O(g) ΔHrxn = kJ

28. Measuring ΔHrxn 𝑞 soln = 𝑚 soln × 𝐶 s,soln ×Δ𝑇 − 𝑞 soln = 𝑞 𝑟𝑥𝑛
Can calculate ΔHrxn using a coffee-cup calorimeter
By recording the change in temperature after a reaction occurs, depending on the mass and solution heat capacity, you can calculate ΔHrxn
Coffee-cup calorimetry occurs at constant pressure
𝑞 soln = 𝑚 soln × 𝐶 s,soln ×Δ𝑇
− 𝑞 soln = 𝑞 𝑟𝑥𝑛
Δ 𝐻 𝑟𝑥𝑛 = 𝑞 𝑝 = 𝑞 𝑟𝑥𝑛

29. Practice
Mg(s) + 2 HCl(aq)  MgCl2(aq) + H2(g) In an experiment to determine the enthalpy change for this reaction, g Mg metal is combined with enough HCl to make mL of solution in a coffee cup calorimeter. The HCl is sufficiently concentrated so that the Mg completely reacts. The temperature of the solution rises from 25.6 °C to 32.8 °C as a result from the reaction. The density of the solution is 1.00 g/mL and the Cs,soln = 4.18 J/(g·°C). What is the mass of the solution? What is the qsoln? What is the qrxn? How many moles of Mg are initially present? What is the ΔHrxn in kJ/mol?

30. Chemical equations and ΔHrxn
Rules for chemical equations and ΔHrxn
1. If a chemical equation is multiplied by some factor, then ΔHrxn is also multiplied by the same factor
2. If a chemical equation is reversed, then ΔHrxn changes sign

31. Chemical equations and ΔHrxn
3. If a chemical equation can be expressed by the sum of a series of steps, then ΔHrxn for the overall equation is the sum of the heats of reaction for each step
Anything on both the product and reactant side cancels out
This is known as Hess’s law

32. Practice Calculate ΔHrxn for the following reaction
Fe2O3(s) + 3CO(g)  2Fe(s) + 3CO2(g)
Use the following reactions
2Fe(s) + 3/2O2(g)  Fe2O3(s) ΔH = kJ
CO(g) +1/2O2(g)  CO2(g) ΔH = kJ

33. Standard enthalpies of formation ΔH°f
ΔH°f is the change in enthalpy when 1 mol of product forms from its constituent elements
H2(g) + ½O2(g)  H2O(g) ΔH°f = kJ/mol
ΔH°rxn can be calculated using ΔH°f of each reactant and product if all compounds are in their standard states

34. Standard states For a gas: when a pure gas is at a pressure of 1 atm
Standard states are defined below
For a gas: when a pure gas is at a pressure of 1 atm
For a liquid/solid: A pure substance in its most stable form at 1 atm and 25° C
For a substance in solution: a substance in a solution at a concentration of 1M

36. Calculating the ΔH°rxn
Standard enthalpy change of a reaction ΔH°rxn – can be determined through the following equation
np = moles of product
nr = moles of reactant

37. Gathering ΔH°f values 4NH3(g) + 5O2(g)  4NO(g) + 6H2O(g)
Identify the ΔHf° of the products and reactants using charts
ΔHf° of NH3(g) = kJ/mol
ΔHf° of O2(g) = 0 kJ/mol
ΔHf° of NO(g) = 91.3 kJ/mol
ΔHf° of H2O(g) = kJ/mol

38. Setting up equation for ΔH°rxn
4NH3(g) + 5O2(g)  4NO(g) + 6H2O(g)
Take the ΔHf° of a product and multiply by the number of moles
Then do this for each product
Then take the sum of those values
4(ΔHf°NO)+ 6(ΔHf°H2O) =
4(91.3kJ/mol)+ 6( kJ/mol) = kJ (products)
Do the same for the reactants
4(ΔHf°NH3)+ 5(ΔHf°O2) =
4(-45.9 kJ/mol)+ 5(0 kJ/mol) = kJ (reactants)

39. Finishing the calculation for ΔH°rxn
Finally: products – reactants
( kJ) – ( kJ) = kJ = ΔH°rxn

40. Problem
What is the ΔH°rxn for the combustion of C3H8(g)?

41. Chapter 6 done.