1. Thermochemistry
Heat and Energy

2. HEAT
The flow of energy from a material with a higher temperature to one with a lower temperature.
When heat is transferred (lost or gained), there is a change in the energy within the substance
Most common method of transfer is conduction

3. Conduction Conduction involves the transfer of energy by contact.
When particles with higher energy (higher T) contact particles with lower energy (lower T), energy is transferred.
This continues until all particles have approximately the same energy.

4. Why do objects feel “hot” or “cold”?
A “hot” object is simply one that transfers energy to your finger when you touch it.
A “cold” object is one that takes energy from your finger when you touch it.
“Hot” and “cold” are relative terms.

5. System vs. Surroundings
The system is what we are talking about:
A chemical reaction
A piece of metal
A beaker of water
The surroundings is all the substances around the system.

6. Direction of Heat Flow Surroundings
Sometimes, energy is transferred from the surroundings INTO the system.
Sometimes, energy is transferred OUT of the system to the surroundings.
System
System
This process is
ENDOTHERMIC
This process is
EXOTHERMIC
Heat (q)
– Heat is thermal energy that can be transferred from an object at one temperature to an object at another temperature
– Net transfer of thermal energy stops when the two objects reach the same temperature.
To study the flow of energy during a chemical reaction, one must distinguish between the system and the surroundings.
System — the small, well-defined part of the universe in which we are interested in studying (such as a chemical reaction)
Surroundings — the rest of the universe (including the container in which the reaction is carried out)
Thermochemical equations are chemical equations in which heat is shown as either a reactant or a product.
Exothermic reaction — process in which heat (q) is transferred from the system to the
surroundings: q < 0
Endothermic reaction — process in which heat is transferred to the system from the
surroundings: q > 0

7. Units for energy
calorie (cal) – defined as the amount of energy needed to raise the temperature of 1 g of water by 1°C.
How much energy is needed to raise the temp. of 5 g of water by 1°C?
How much energy is needed to raise the temp. of 5 g of water by 10°C?

8. Units for energy
Calorie (Cal) – this is a dietary unit. It is actually equal to 1 kilocalorie.
1 Cal = 1 kcal = 1000 cal
Joule (J) – this is the SI unit for energy. We will be using this unit primarily.
1 cal = J

9. Units for energy Heat energy is abbreviated with a lower case q.
In an exothermic process, q < 0.
In an endothermic process, q > 0.
The sign of q represents direction more than value.
-180 J is more energy than +120 J. The signs just indicate whether heat is going into the system or out of the system.

10. Enthalpy, H

11. Enthalpy is a measure of the stored energy in a system.
Enthalpy cannot be measured.
Instead, we measure the change in enthalpy (ΔH) as a process occurs.
We will learn ways to do this later in the unit. For now, you will be given Δ H, as you will see in the example problems.

12. Why are some reactions endothermic and others are exothermic?
To break a chemical bond (or to weaken attractive forces) requires energy. During a chemical reaction, the bonds of the reactants are broken.
When bonds are formed (or attractive forces take hold), energy is released. During a chemical reaction, energy is released as chemical bonds of the products are formed.

13. Why are some reactions endothermic and others are exothermic?
If more energy is used to break the bonds of the reactants than is released by the formation of bonds of the products, the reaction is endothermic.
The net result is energy being put into the system.
For an endothermic reaction, ΔH > 0.

14. ΔH>0

15. Why are some reactions endothermic and others are exothermic?
If less energy is used to break the bonds of the reactants than is released by the formation of bonds of the products, the reaction is exothermic.
The net result is energy being released from the system.
For an exothermic reaction, ΔH < 0.

16. ΔH<0

17. ΔH sounds a lot like q ?!
ΔH refers to a specific process. For that process, ΔH doesn’t change. ΔH is a ratio.
q is simply a measure of the total heat exchanged.
q = nΔH
As an analogy, ΔH is like molar mass and q is like mass.
You can have any mass of water, but the molar mass will always be 18.0 g/mol.

18. ΔH sounds a lot like q ?!
As an analogy, ΔH is like molar mass and q is like mass.
You can have any mass of water, but the molar mass will always be 18.0 g/mol.
You can release any amount of energy while burning methane (and therefore, get any value for q), but the ΔH for this process will always be kJ/mol.

19. Look at the following reaction
Fe2O3 + 3 CO  2 Fe + 3 CO2 ΔH = -23 kJ/molrxn
What kind of reaction is this?
How do you know?
Is energy absorbed or released by the system?
More energy was released in forming the bonds in the products than was used to break the bonds in the reactants.

20. Fe2O3 + 3 CO  2 Fe + 3 CO2 ΔH = -23 kJ/molrxn
For every 1 mole of Fe2O3 that reacts, 23 kJ of energy are released.
For every 3 moles of CO that react, 23 kJ of energy are released.
For every 2 moles of Fe that form, 23 kJ of energy are released.
For every 3 moles of CO2 that form, 23 kJ of energy are released.

21. Fe2O3 + 3 CO  2 Fe + 3 CO2 ΔH = -23 kJ/molrxn
What is the value of q when 3.50 moles of iron (III) oxide react?
According to the balanced equation, 23 kJ are released for every 1 mol of Fe2O3. ΔH = -23 kJ/mol Fe2O3

22. Fe2O3 + 3 CO  2 Fe + 3 CO2 ΔH = -23 kJ/molrxn
How much energy is released when 3.50 moles of iron (III) oxide react?
How is this question different?
Let’s use unit analysis this time.

23. Fe2O3 + 3 CO  2 Fe + 3 CO2 ΔH = -23 kJ/molrxn
1. How much energy is released when 14.5 grams of carbon dioxide are formed?
2. How many grams of Fe2O3 must react to release 140 kJ of heat?

24. CH4 + H2O  3 H2 + CO ΔH = +206 kJ/molrxn
Determine the amount of energy needed to react 3.25 moles of methane.
Determine the mass of H2O that will react when 95.2 kJ of energy are absorbed by the reaction.