1. Energetics and Thermochemistry
Topic 5
Energetics and Thermochemistry

3. Remember from 9th grade…
Work- force acting over a distance (in the same direction)
Heat- involves the transfer of energy between two objects due to a temperature difference
Temperature- property that reflects the random motion of particles in a substance
Law of Conservation of Energy- energy can be converted from one form to another, but it can be neither created nor destroyed
Kinetic Energy- energy due to motion (1/2mv2)
Potential Energy- energy due to position or composition (mgh)

4. Energy is the ability to do work or transfer heat.
Energy used to cause an object that has mass to move is called work.
Energy used to cause the temperature of an object to rise is called heat.
This chapter is about thermodynamics, which is the study of energy transformations, and thermochemistry, which applies the field to chemical reactions, specifically.

5. Kinetic Energy
Kinetic energy is energy an object possesses by virtue of its motion:

6. What is Energy?
“the ability to do work” or “the capacity to do work or produce heat”
Come up with your own definition
Different forms + conversions
Did you eat breakfast today?
Did you drive your car to school today?
Wind turbines

7. Definitions: Work
Energy used to move an object over some distance is work:
w = F  d
where w is work, F is the force, and d is the distance over which the force is exerted.

8. Heat Energy can also be transferred as heat.
Heat flows from warmer objects to cooler objects.

9. Conversion of Energy Energy can be converted from one type to another.
The cyclist has potential energy as she sits on top of the hill.
As she coasts down the hill, her potential energy is converted to kinetic energy until the bottom, where the energy is converted to kinetic energy.

10. Units of Energy The SI unit of energy is the joule (J):
An older, non-SI unit is still in widespread use, the calorie (cal):
1 cal = J
(Note: this is not the same as the calorie of foods; the food calorie is 1 kcal!)

11. What is a system? System- Surroundings- Open System- Closed System-
Isolated system- no exchange of energy or matter is possible

12. Definitions: System and Surroundings
The system includes the molecules we want to study (here, the hydrogen and oxygen molecules).
The surroundings are everything else (here, the cylinder and piston).

13. First Law of Thermodynamics
Energy is neither created nor destroyed.
In other words, the total energy of the universe is a constant; if the system loses energy, it must be gained by the surroundings, and vice versa.

14. What is Enthalpy? The heat content of a system (H)
Total energy remains the same, but heat can be transferred between a system and the surroundings
Changes in enthalpy=
Heat added to the system ΔH =
Heat released from the system ΔH =

15. Exchange of Heat between System and Surroundings
When heat is absorbed by the system from the surroundings, the process is endothermic.

16. Exchange of Heat between System and Surroundings
When heat is released by the system into the surroundings, the process is exothermic.

18. Exothermic v. Endothermic
Enthalpy is stored in chemical bonds and intermolecular forces as potential energy
Most chemical reactions are exothermic
Exothermic-
Endothermic-

19. State Functions
Usually we have no way of knowing the internal energy of a system; finding that value is simply too complex a problem.
However, we do know that the internal energy of a system is independent of the path by which the system achieved that state.
In the system below, the water could have reached room temperature from either direction.

20. Standard Enthalpy Changes
What is a state function?
Depends on the present state of the system (not the past or future of the system)
(state functions are dependent on the conditions, or present state, under which a change occurs so the conditions must be specified)
Standard enthalpy changes (ΔH°)
Pressure of 100kPa
Concentration of 1mol/dm3
All substances in their “standard states”

21. Enthalpy of Reaction
The change in enthalpy, H, is the enthalpy of the products minus the enthalpy of the reactants:
H = Hproducts − Hreactants

22. Enthalpy of Reaction
This quantity, H, is called the enthalpy of reaction, or the heat of reaction.

23. The Truth about Enthalpy
Enthalpy is an extensive property.
H for a reaction in the forward direction is __________ in size, but _____________ in sign, to H for the reverse reaction.
H for a reaction depends on the state of the ____________ and the state of the _____________.

24. Calorimetry
Since we cannot know the exact enthalpy of the reactants and products, we measure H through calorimetry, the measurement of heat flow.
The instrument used to measure heat flow is called a calorimeter.

25. Heat Capacity and Specific Heat
The amount of energy required to raise the temperature of a substance by 1 K (1 C) is its heat capacity, usually given for one mole of the substance.

26. What is specific heat?
Heat changes can be calculated from temperature changes
These changes generally depend on the mass of the object, the amount of heat added, and the nature of the substance
The specific heat capacity of a substance describes the amount of heat needed to increase the temperature of unit mass (g or kg) by one degree K or C

27. Heat Capacity and Specific Heat
We define specific heat capacity (or simply specific heat) as the amount of energy required to raise the temperature of 1 g of a substance by 1 K (or 1 C).

29. Heat Capacity and Specific Heat
Specific heat, then, is

30. Calorimetry!! Calorimetry is the science of measuring heat
A calorimeter is a device used to experimentally determine the heat associated with a reaction

31. Super fancy calorimeter (that we will use)

33. More on Enthalpy Changes
There is a natural direction in which change occurs
When you slip on a ladder, which way do you fall?
Changes happen in the direction of lower stored energy
Chemical changes take place to reduce enthalpy (products are more stable than reactants)

34. Great textbook example

35. Calorimetry problems
Type 1- Thermal Equilibrium -- you add a sample of something (usually solid) to something else (usually liquid) and the resulting mixture will have the same final temperature (ex. Adding a sample of a heated metal to water to find the specific heat of the metal)
The heat then flows from the metal into the water so the heat energy LOST by the metal is equal to the heat energy GAINED by the water
Thus, -qmetal = +qwater
–mcΔT(metal) = +mcΔT(water)

36. Type 1 - Example
A 28.4 g sample of aluminum is heated to 39.4 oC, then is placed in a calorimeter containing 50.0 g of water. Temperature of water increases from oC to oC. What is the specific heat of aluminum?
[the specific heat of water is 4.184J/g°C and the final temperature of water and aluminum are the same]

38. Calorimetry problems
Type 2- Combustion reaction problems (assume excess oxygen)
The heat from an exothermic reaction is absorbed by the water, and the temperature of the water rises
The temperature change depends on the amount of the reaction that took place, so the units for the answer will be kJ/mol
ΔHreaction = -ΔHwater/molreactant =
[-mwater*cwater*ΔTwater]/molreactant

39. Type 2 Practice Problem p. 219
Calculate the enthalpy of combustion of ethanol from the following data. Assume all the heat from the reaction is absorbed by the water. Compare your value with the IB data booklet value and suggest reasons for any differences.

41. Checking the answer IB data booklet value is -1367 kJ/mol
Possibilities for the difference:
Not all of the heat produced by the combustion reaction is transferred to the water. Some is needed to heat the copper calorimeter and some passes to the surroundings.
The combustion of the ethanol is unlikely to be complete owing to the limited oxygen available
The experiment was not performed under standard conditions

42. Calorimetry problems Type 3- Reactions/neutralization reactions
A reaction takes place in a calorimeter in aqueous solution (two solutions are combined, or a solid is added to an aqueous solution)
The enthalpy change can again be calculated by measuring the temperature change, but you must first find the limiting reactant

43. Type 3 - Practice Problem
1.00 dm3 of a 1.00 mol/dm3 solution of copper (II) sulfate is placed in a calorimeter with an initial temperature of 20°C moles of Zn(s) is added, and the temperature reaches a maximum of 70°C.

44. Assumptions No heat is lost from the system
All the heat goes from the reaction to the water
The density of the copper (II) sulfate solution is the same as water (4.184J/g°C) since the solutions are dilute
Water has a density of 1.00g/cm3