1. Heat of Reaction & Enthalpy

2. General Information
All chemical reactions involve energy changes, whether energy is being absorbed or given off
The energy absorbed or given off comes from chemical bonds
When bonds are broken, energy is absorbed
Energy is needed in order to break the bonds
When bonds are formed, energy is released
Energy is given off when the compound becomes more stable after a bond forms

3. Temperature ≠ Heat Temperature Heat
The measurement of the average kinetic energy of the particles in a sample of matter
SI unit is Kelvin (K)
Temperature is a number - it is not energy
Heat
Energy that flows between samples of matter because of a difference in the temperature
Flows from hot to cold
SI unit is Joules (J)
Heat is energy
Example: Look at ice melting
Temperature stays at 0°C
But heat is being absorbed

4. Flow of Energy Exothermic Endothermic
The system releases heat to its surroundings
Feels warm or hot because the release of heat increases the kinetic energy of the surroundings
When energy is release, an energy term will appear on the product side of the equation
Endothermic
The system absorbs heat from its surroundings
Feels cold or cool because it is absorbing energy from the surroundings. This decreases the kinetic energy of the surroundings
Is indicated by writing the energy term on the reactant side of the equation

5. If a mixture of hydrogen and oxygen is ignited, water will form and energy will be released explosively
2H2(g) + O2(g)  2H2O(g)
Because energy is released/produced – the reaction is exothermic
This equation does not tell you that energy is evolved as heat during the reaction

6. 2H2(g) + O2(g)  2H2O(g) + 483.6 kJ But through experiments we know:
483.6 kJ of energy is released/produced
Adding the amount of energy makes it a thermochemical equation
2H2(g) + O2(g)  2H2O(g) kJ
States must be included because they influence the overall amount of energy exchanged

7. Enthalpy Enthalpy (H°)
Is the heat content in a system, OR, the total amount of potential and kinetic energy within a substance
Enthalpy cannot be measured
Note: The ° symbol indicates a temperature of 25°C and 101kPa, which is standard temperature and pressure in thermodynamics

8. Enthalpy Change
An enthalpy change (ΔH°) is the amount of energy absorbed or lost by a system as heat during a process at constant pressure
We can calculate the enthalpy changes during a reaction
We also call this change in enthalpy the heat of reaction

9. Thermochemical equations can also be written by designating the value of ΔH
For exothermic reaction, ΔH is always negative (-) because the system loses energy
Endothermic reactions, ΔH is always positive (+) because the system gains energy
Same as before – exothermic
2H2(g) + O2(g)  2H2O(g) ΔH° = kJ
Opposite process - endothermic
2H2O(g)  2H2(g) + O2(g) ΔH° = kJ

10. Potential Energy Graphs
Graphically, we can represent the potential energy for the reactants and products like:
Shows the potential energy of the reactants is more than the potential energy of the products
The drop in potential energy is the heat released or the heat of reaction (ΔH)
Heat is being released so the value is negative

11. Endothermic is the complete opposite
Shows the potential energy of the reactants is less than the potential energy of the products
The increase in potential energy is the heat gained or the heat of reaction (ΔH)
Heat is being absorbed so the value is positive

12. Questions? HW 1.8 (8-10) 1.9 (1-3) Tomorrow: Heat cont’d
Review – Let me know if you have specific topics you want to review

13. Heat of Reaction & Enthalpy
Day 2

14. Review An exothermic reaction releases energy
Energy can be seen as a product of the reaction
An endothermic reaction needs energy for the reaction to occur
Energy can be seen as a reactant
Enthalpy (H) is the energy (heat) content of a system at constant pressure
You cannot measure the actual energy or enthalpy of a substance
You can measure the change in enthalpy ΔH
An enthalpy change (ΔH°) is the amount of energy absorbed or lost by a system as heat during a process at constant pressure

15. Chemistry problems involving enthalpy changes are similar to stoichiometry problems
The amount energy that is absorbed or released in a reaction depends on the number of moles of reactants involved

16. 2H2O2(l)  2H2O(l) + O2(g) ΔH= -190kJ
Example
How much heat will be released if 1.0 g of hydrogen peroxide decomposes
2H2O2(l)  2H2O(l) + O2(g) + 190kJ OR
2H2O2(l)  2H2O(l) + O2(g) ΔH= -190kJ

17. 2 NaHSO4(s) ----> Na2SO4(s) + H2O(g) + SO3(g)
Another Example
Given the decomposition of sodium hydrogen sulfate reaction:
2 NaHSO4(s) ----> Na2SO4(s) + H2O(g) + SO3(g)
ΔH= kJ
If 3.60 g of NaHSO4(s) reacts, how much heat is released?

18. Given the following reaction:
C(s) + 2S(s) kJ  CS2(l)
How many atoms (particles) of Carbon can be burned if kJ of energy are available?

19. Time to work on HW 1.9 (4-10)
Review
Tomorrow: Move on to Unit 2 – Solutions
Monday: Unit 1 EXAM