1. SOURCE

2. Energy and Chemical Change
Measuring and calculating the energy involved in chemical changes

3. Energy Energy is defined as the ability to do work or produce heat.
Kinetic Energy
Potential Energy
Energy is defined as the ability to do work or produce heat.
SI unit for energy is Joules (J).
There are 2 types of energy

4. Energy 1. Potential Energy (PE)
Energy that is based on an objects composition (chemical PE) or position;
associated with attractions and repulsions.
Chemical potential energy
the energy stored in a substance because of its composition (ex: plants, gasoline)
plays an important role in chemical reactions

5. Energy 2. Kinetic Energy (KE) - Energy of motion. Kinetic Energy
Can be measured as temperature or heat.
KE of a substance is directly related to the motion of its atoms and temperature.

6. Energy
* Every object has either potential or kinetic energy!!!

7. Energy
Law of Conservation of Energy: energy can be converted from one form to another, but it is neither created nor destroyed.

8. Heat
Form of energy associated with changing the temperature of an object.
An object’s temperature increases because energy is transferred into it
Temperature is a measure of the average kinetic energy of the particles in a sample of matter; it does not depend on the amount of matter in the sample.

9. Heat
The heat (energy) transfer occurs when 2 objects of different temperatures are brought into contact to reach equilibrium.
Depends on the mass of the sample.
Symbol = m Unit = grams
Reactions can be endothermic (heat/energy in) or exothermic (heat/energy out).

10. How Can I Measure Energy/Heat?
There are many units that can express an amount of energy.
calorie (cal) - amount of heat required to raise the temperature of 1.00 g of pure water by 1oC.

11. How Can I Measure Energy/Heat?
kilocalorie (kcal) calories or one Calorie – this is one food Calorie!
Joule - SI unit of heat and energy. (What we will mainly use)
Important conversions:
1 Calorie = 1 kcal
1 kcal = 1000 cal
1 cal = J

12. How Can I Measure Energy/Heat?
heat changes that occur during chemical and physical processes can be measured accurately and precisely using a calorimeter
a calorimeter is an insulated device used for measuring the amount of heat absorbed or released during a chemical or physical process

13. How Can I Measure Energy/Heat?
Inside most calorimeters you will need:
a known mass of water is placed in an insulated chamber in the calorimeter
it is designed to absorb the energy released from the reacting system (unknown/known substance) or to provide the energy absorbed by the system
the data collected is the change in temperature of the water inside the calorimeter

14. How Can I Measure Energy/Heat?
most calorimetric reactions occur at constant pressure
can also be used to determine the specific heat of an unknown metal
assuming that no heat is lost to the surroundings, the heat lost by the metal is equal to the heat gained by the water

15. What is Specific Heat? (Or, How do I Calculate Amount of Energy Transferred?)
Specific Heat is the amount of heat required to raise the temperature of one gram of any substance by 1oC.
Units = J/g x OC of kJ/g x oC
Each substance has its own specific heat.
Knowing the specific heat of a substance allows you to calculate the heat transferred by measuring the mass of the substance and the temperature change.

16. What is Specific Heat? (Or, How do I Calculate Amount of Energy Transferred?)
Specific heat of water = J/ g x oC
Formula: q = c x m x ∆T
c = specific heat
m = mass
∆T = Tfinal – Tinitial
Heat transferred = (specific heat)* (mass)* (change in T measured in OC )
q = (g)*(J/g x OC)*(∆T)

17. What is Specific Heat? (Or, How do I Calculate Amount of Energy Transferred?)
--This equation also indicates which direction the heat is traveling.
**If ∆T is (+) then heat is being transferred to or absorbed by the substance (endothermic)
**If ∆T (-) then heat is being released during the reaction (exothermic)

18. Keep in mind…
the specific heat of a substance is a measure of how efficiently that substance absorbs heat

19. PRACTICE!!!
Problem- How much heat is lost when solid Al with a mass of 4110 g cools from oC to 25 oC?
(Specific Heat of Al is J/g x oC)
First determine the ΔT = Tfinal -Tinitial
(25oC-660oC = -635oC)
Remember the equation is q= cxmxΔT
c = specific heat
m = mass

20. PRACTICE!!! Plug in the values q = (0.902 J/gxoC)*(4110g)*(-635oC)
Cancel like units
Multiply all values
J or divide by 1000 and get kJ
Significant figures gives us -2354kJ or
-2.35 x 106J
Type of reaction: EXOTHERMIC

21. PRACTICE!!!
Problem- How much heat is required to raise the temperature of 854 g of H2O from 23.5 oC to 85.0 oC?
ΔT = Tfinal – Tinitial
ΔT = 85.0oC – 23.5oC = 61.5oC
q = c x m x ΔT
q = (4.184 J/g x oC)*(854g)*(61.5oC)= J or 219.7kJ or 2.20 x 105J
Type of reaction: ENDOTHERMIC

22. So, What is Thermochemistry?
Thermochemistry is the study of heat changes that accompany chemical reactions and phase changes.
***Remember: Energy often takes the form of HEAT
As a chemical reaction progresses some of the energy in the bonds may be released as heat. These reactions are called EXOTHERMIC.

23. So, What is Thermochemistry?
In an exothermic reaction the products have less potential energy than the reactants.
Exothermic reactions release energy; more energy is released than is required to break bonds in the initial reaction.

24. So, What is Thermochemistry?
Examples: burning wood, any reaction in which the test tube or beaker gets warm, formation of water (explosion of hydrogen and oxygen), an athletic hot pack.
the contents of the athletic hot pack is the system which is the specific part of the universe that contains the reaction or process you wish to study
the universe is defined as the system plus the surroundings (which is everything in the universe other than the system)

25. So, What is Thermochemistry?
Some chemical reactions absorb heat as they progress. These reactions are called ENDOTHERMIC.
In an endothermic reaction the products have more heat energy (enthalpy) than the reactants.

26. So, What is Thermochemistry?
Reaction that needs energy to complete the reaction; a greater amount of energy is required to break the initial bonds than is released from the new bonds forming
Examples: cooking eggs, any reaction in which the beaker or test tube gets cooler, an athletic cold pack, decomposition of water (electrolysis).

27. Quick Quiz!!!
Identify each of the following as endothermic or exothermic:
a. burning natural gas (methane)
exothermic
b. baking a cake
endothermic
c. melting ice

28. What is Enthalpy?
Enthalpy is the chemistry word for heat content (remember that chemical compounds have POTENTIAL energy in their bonds – this is the heat content).
Enthalpy is represented by the symbol H
You cannot measure the actual enthalpy of a substance…instead you measure the change in enthalpy (the heat absorbed or released in a reaction)

29. What is Enthalpy?
The change in enthalpy is called the enthalpy heat of reaction and is represented by ∆H or ∆Hrxn.
ΔHrxn = Hproducts – Hreactants

30. Enthalpy of Exothermic Reactions
Enthalpy changes for exothermic reactions are always negative. ∆H is (-)
In an exothermic reaction the products have less potential energy than the reactants
Hprod. < Hreact.

31. Enthalpy of Exothermic Reactions
Energy is a product
Example:
CH4(g) + 2O2(g)  CO2(g) + 2H2O(g) kJ product
CH4(g) O2(g)  CO2(g) + 2H2O(g)
ΔH = kJ (because it is on the product side)

32. Enthalpy of Endothermic Reactions
Enthalpy changes for endothermic reactions are always positive. ∆H is (+)
In an endothermic reaction the products have more heat energy (enthalpy) than the reactants.
Hprod > Hreact

33. Enthalpy of Endothermic Reactions
Energy is a reactant
Example:
2H2O(l) kJ  2H2(g) O2(g)
2H2O(l)  2H2(g) O2(g)
ΔH = 572 kJ (because it is on the reactant side)

34. Thermochemical Equations
Thermochemical equations are balanced chemical equations that include the physical states of all reactants and products and the energy change.
2H2O(l)  2H2(g) + O2(g) ΔH = 572 kJ
CH4(g) + 2O2(g)  CO2(g) + 2H2O(g) ΔH = kJ

35. Enthalpy of Combustion
Enthalpy heat of combustion (ΔHcomb) is the enthalpy change for the complete burning of one mole of the substance.
-- carried out under standard conditions which are one atmospheric pressure (1 atm) and 298K (250C)
- C6H1206(s) + 6O2(g)6CO2(g)+ 6H2O(l)
ΔHcomb = -2808kJ

36. Changes of State
Molar enthalpy (heat) of vaporization (ΔHvap) is the heat required to vaporize one mole of liquid.
-- think of water vaporizing from your skin after you take a hot shower. Your skin provides the heat needed to vaporize the water and as the water absorbs the heat you feel cool (shiver)
ΔHvap = -ΔHcond (condensation)

37. Changes of State
Molar enthalpy (heat) of fusion (ΔHfus) is the heat required to melt one mole of a solid substance.
--think of ice in a drink. The drink cools as it provides the heat for the ice to melt
ΔHfus = - ΔHsolid (solidification—freezing)

38. Phase Changes
Occurs when energy is added or removed from a system and the substance can go from one physical phase to another

39. Endothermic Phase Changes
Melting
The energy absorbed to melt a solid is not used to raise the temperature of that solid
The energy instead disrupts the bonds holding the solid’s molecules together and cause the molecules to move into the liquid phase

40. Endothermic Phase Changes
The amount of energy required to melt one mole of a solid depends on the strength of the forces that hold the solid together
The melting point of a crystalline solid is the temperature at which the forces holding its crystal lattice together are broken and it becomes a liquid

41. Endothermic Phase Changes
Vaporization
Particle that escape from the liquid enter the gas phase and those liquids at room temperature the gas phase is called vapor
Vaporization is the process by which a liquid changes into a gas or vapor
Once the solid becomes a liquid then and only then does the temperature of the substance begin to increase

42. Endothermic Phase Changes
When vaporization takes place only at the surface of the liquid it is called evaporation
Evaporation is the method by which the human body maintains and controls its temperature

43. Endothermic Phase Changes
The pressure exerted by a vapor over a liquid is called the vapor pressure
The temperature at which the vapor pressure of a liquid equals the external or atmospheric pressure is called the boiling point

44. Endothermic Phase Changes
Sublimation
Is the process by which a solid changes directly to a gas without first becoming a liquid
Dry ice (CO2) and snow are the most common examples

45. Endothermic Phase Changes
If ice cubes are left in the freezer for extended periods of time, they will eventually sublime and become smaller
This process is also helpful in freeze drying foods for hikers and astronauts

46. Exothermic Phase Changes
Condensation
When a vapor molecule loses energy its velocity is reduced therefore colliding more with other molecules to form a liquid
Condensation is the process by which a gas or vapor becomes a liquid and it is the reverse action of vaporization

47. Exothermic Phase Changes
Deposition
Is the process by which a substance changes from a gas or vapor to a solid without first becoming a liquid
It is the reverse action of sublimation
The formation of snow crystals high up in the atmosphere is an example

48. Exothermic Phase Changes
Freezing Point
Is the temperature at which a liquid is converted into a crystalline solid
The same temperature as the melting point of a given substance

49. Phase Change Graph

50. Phase Change Graph
Graph shows the energy required to go from one phase to the other
Where the graph inclines, potential energy is at its greatest and temperature is increasing
Where the graph plateaus (flat region) kinetic energy is at its greatest but the temperature remains constant

51. Phase Diagrams
A phase diagram is a graph of pressure versus temperature that shows in which phase a substance exists under different conditions of temperature and pressure

52. Phase Diagrams
The triple point is the point on a phase diagram that represents the temperature and pressure at which three phases of a substance can coexist
The critical point is the point that indicates critical pressure and temperature above which water cannot exist as a liquid

53. Phase Diagrams
Different for each substance because of the different boiling/freezing points