1. Energy

2. Energy Energy is everywhere
It is in the form of light, heat, electricity and many other forms
Energy is the driving force for things to move and stop

3. Energy Molecules are held together by chemical bonds
In order to break bonds, we need to add energy to the bond to break it
Think about a rubber band. If you let the rubber band sit there, it will never break. But if you stretch it by putting energy in, it will break.
We add energy (heat, light, electricity…) to break chemical bonds
All the energy we deal with is potential energy

4. Breaking Bonds 2H2O + Energy → 2H2 + O2
We have to ADD energy to water to break its bonds to form H2 and O2.
This is why it is so hard to create enough fuel for hydrogen cars.
Because we added energy to the reactants, the products will have more energy than the reactants

5. Potential Energy Diagram

6. Forming Bonds
When two atoms or molecules join together, they must GIVE off energy.
If not, the molecules/atoms will have too much energy and will simply collide and fly apart.
Chemicals are joined/reacted together by collisions.
2H2 + O2 → 2H2O + Energy
We have energy as a product meaning it is produced and released.

7. Potential Energy Diagram

8. Enthalpy Enthalpy is the HEAT contained in a system.
This is the energy within the system.
We represent it with a number in kJ/mol
It is represented by the letter H
∆H is the change in enthalpy during a reaction as we go from the reactants to the products
∆H = Hproducts – Hreactants
Hproducts is the heat contained in the products
Hreactants is the heat contained in the reactants

9. Enthalpy
So the energy we have to form and break bonds will leave the system eventually
We will either feel something hot or cold
There are special names for these cases
Exothermic
Endothermic

10. Exothermic Exothermic reactions gives off heat to its surroundings.
To remember: EXothermic = Exit
Heat exits from the system into the surroundings
So things get hot around the system in the surroundings
CH4 + 2O2 → CO2 + 2H2O + Energy
CH4 + 2O2 → CO2 + 2H2O + Heat
CH4 + 2O2 → CO2 + 2H2O + 891kJ/mol
We call this notation thermochemical equation

11. Exothermic The -891kJ/mol is from our calculation
∆H = Hproducts – Hreactants
Note that it is negative as the product has less energy than the reactants
The heat term is on the product side
This means it is exothermic
Do not have to worry about how we get ∆H. We just need to know exothermic is on product side. Thus the products have LESS energy than the reactants.

14. Endothermic Endothermic reaction absorbs heat from its surroundings
To Remember: ENdothermic = Enter
Heat will enter the system from the surroundings
So things gets cold as heat from the surroundings (hand, bottle, beaker) is absorbed into the reaction
KClO3 + Energy → K+ + ClO3-
KClO3 + Heat → K+ + ClO3-
KClO kJ/mol → K+ + ClO3-

15. Note!
You may think that when energy is absorbed, the object should get hot/warm, as there is more energy. But remember, conservation of energy. The energy is going from the surroundings into the BONDS. So the surroundings get cooler as it has less energy
You are feeling the surroundings, not the system/reaction itself

16. Endothermic The 41.4kJ/mol is from our calculation
∆H = Hproducts – Hreactants
Note that is it positive as the products has more energy than the reactants
The heat term is on the reactant side for endothermic
The reactants has less energy than the products

19. The general rule
If MORE energy is used to BREAK the bonds in the reactants than it is given off when we PRODUCE the bonds, the reaction is ENDOTHERMIC
If LESS energy is used to BREAK the bonds in the reactants than it is given off when we PRODUCT the bonds, the reaction is EXOTHERMIC

20. Real Life Examples
Candle Flame – Exothermic because the heat is produced and released as the wax is being reacted and burned off.
Ice Cube – Endothermic because it is absorbing heat from your hands as you hold it to melt so it makes your hands cooler.
Explosions – Exothermic as the reaction is releasing heat
Ice Packs – Endothermic because as the reaction takes place, it is absorbing heat from the pack/bag.

21. Another way to represent enthalpy
Remember that the equation for enthalpy is ∆H = Hproducts – Hreactants
We put the ∆H outside of the reaction in parentheses.
Endothermic
A → B ; ∆H = +800kJ/mol
Remember that endothermic, the products have more energy than the reactants so positive H
Exothermic
A → B ; ∆H = -800kJ/mol
As the products have less energy than the reactants so we have negative H
We call this ∆H notation

22. Page 122 for summary box

23. Moles and Heat
If you have more particles reacting and releasing energy, you should assume more heat
Think about an explosion, more TNT, bigger the explosion
So how does moles affect heat?
It is directly proportional

24. Moles and Heat
If it takes 800kJ/mol to form a bond of O2, then it should take 1600kJ/mol to form 2 O2 bonds.
It should take 8000kJ/mol to form 10 O2 bonds.
The heat term written in equations is for the overall reaction

25. Example - 1
Convert the following ∆H notation into thermochemical equation using the smallest coefficient possible
(1/2)B + (3/2)H2O → (1/4)B2H6 + (3/4)O2 ∆H = 762kJ/mol
Convert the following thermochemical equation into ∆H notation using the smallest coefficient possible
F2 + (1/2)O2 → OF2 + 22kJ/mol

26. Moles and Heat
Remember, the heat term in equations is for the overall reaction
So if you have multiple molecules, the heat term is split up according to how many moles are present

27. Mole and Heat We can turn all of these into conversion factors!
CH4 + 2O2 → CO2 + 2H2O + 891kJ/mol
When 1 mole of CH4 is consumed, 891kJ/mol of heat is released
When 2 mole of O2 is consumed, 891kJ/mol of heat is released
When 1 mole of CO2 is produced, 891kJ/mol of heat is released
When 2 mole of H2O is produced, 891kJ/mol of heat is released
We can turn all of these into conversion factors!

28. Mole and Heat
So for CH4 , the conversion factor turns out to be 1molCH4/891kJ or 891kJ/1molCH4
O2 will be 2molO2/891kJ or 891kJ/2molO2
ETC ETC

29. Example - 2 CH4 + 2O2 → CO2 + 2H2O + 891kJ/mol
How much heat is released during the formation of 12.5 moles of CO2?
How much heat is released during the formation of mol of H2O?
If 26.79g of CO2 are produced, how much heat is released?

30. Practice - 1
Page #68-69
Page #70-75
Page #76-80