1. Welcome to CH104 – General Chemistry
Prayer
Syllabus
Lab times and days
Review of Enthalpy

2. The Nature of Energy Systems and Surroundings
System: part of the universe we are interested in.
Surroundings: the rest of the universe.

3. Internal Energy Internal Energy: total energy of a system.
Cannot measure absolute internal energy.
Change in internal energy, DE = Efinal - Einitial

4. First Law of Thermodynamics
Energy cannot be created or destroyed.
Energy of (system + surroundings) is constant.
Any energy transferred from a system must be transferred to the surroundings (and vice versa).
From the first law of thermodynamics:
when a system undergoes a physical or chemical change, the change in internal energy is given by the heat added to or absorbed by the system plus the work done on or by the system:
DE = q + w

5. First Law of Thermodynamics
Relating DE to Heat and Work

6. First Law of Thermodynamics
Endothermic and Exothermic Processes
Endothermic: absorbs heat from the surroundings.
Exothermic: transfers heat to the surroundings.
An endothermic reaction feels cold.
An exothermic reaction feels hot.

7. First Law of Thermodynamics
State Functions
State function: depends only on the initial and final states of system, not on how the internal energy is used.

8. E = qengine + wengine
E = qstove
A block diagram of two different paths for the combustion of octane.
Occurring inside an engine (path1), the process transfers heat and work.
Occurring in a stove burner (path 2), the process does no work and
transfers only heat.

9. DH = Hfinal - Hinitial = qP
Enthalpy
Enthalpy, H: Heat transferred between the system and surroundings carried out under constant pressure.
Can only measure the change in enthalpy:
DH = Hfinal - Hinitial = qP

10. Enthalpies of Reaction
DHrxn = H(products) - H (reactants)
Enthalpy is an extensive property
(magnitude DH is directly proportional to amount):
CH4(g) + 2O2(g)  CO2(g) + 2H2O(g) DH = -802 kJ
2CH4(g) + 4O2(g)  2CO2(g) + 4H2O(g) DH = kJ
When we reverse a reaction, we change the sign of DH:
CO2(g) + 2H2O(g)  CH4(g) + 2O2(g) DH = +802 kJ
Change in enthalpy depends on state:
H2O(g)  H2O(l) DH = -88 kJ

11. Hess’s Law
In the above enthalpy diagram note that
H1 = H2 + H3

12. Enthalpies of Formation
If 1 mol of compound is formed from its constituent elements, then the enthalpy change for the reaction is called the enthalpy of formation, Hof .
Standard conditions (standard state): 1 atm and 25 oC
Standard enthalpy, Ho, is the enthalpy measured when everything is in its standard state.
Standard enthalpy of formation: 1 mol of compound is formed from substances in their standard states.

13. Enthalpies of Formation
Standard enthalpy of formation of the most stable form of an element is zero.

14. Enthalpies of Formation
Using Enthalpies of Formation to Calculate Enthalpies of Reaction
Hrxn =
H1 + H2 + H3

15. Enthalpies of Formation
Using Enthalpies of Formation to Calculate Enthalpies of Reaction
For a reaction:
DHOrxn = SnDHOf(products) -SnDHOf(reactants)