Chapter 11: Thermodynamics

Essential Definitions

State functions describes an aspect of a chemical system
- Pressure, volume, temperature, and moles are state functions in the ideal gas law
Thermodynamics state functions can be internal energy (E), enthalpy (H), entropy (S), and Gibbs free energy (G)
Changes in state function depend on the initial and final state of a system
Not state functions are heat (q) and work (w) because they rely on the steps used to transform matter from initial to final

Extensive properties of matter change as the amount of sample changes
- Such as ΔH, ΔS, ΔG, and ΔE
To make them intensive properties, the temperature, pressute, mass, and physical state must be defined
Standard state is 1 atm, 25 celsius, and 1 mol of compound present
- This gives them a superscript of zero to note that they are at standard state
- ΔH°, ΔS°, ΔG°, ΔE°

Exo-, or exothermic, means energy is being lost from the system to the surroundings
- Negative sign for quantities
- ΔH is negative if exothermic
Endo-, or endothermic, means energy is being gained from surroundings
- absorbs heat energy
- ΔH is positive

Energy can be heat, light, chemical, nuclear, electrical, or mechanical
Law of conservation of energy states energy cannot be created or destroyed. All forms of energy can be converted to heat energy
Can be called kinetic, KE, or potential, PE, energy
- Kinetic is the energy matter has because of its motion
- potential energy is stored and released under certain conditions
- gravitational energy is PE
- electrostatic attraction between oppositely charge ions is PE

Calorie is the amount of heat needed to raise the temperature of 1 gram of something by 1 degree celsius. 1 calorie = 4.184 joules
- The specific heat of water is the energy needed to raise 1 gram of water by 1 degree celsius = 4.184 J/g*C
Specific heat equation
- q = sp.ht x g x ΔT
- heat energy = (specific heat)(mass in grams)(temperature change)

First law of thermodynamics states energy is always conserved
ΔE = q + w
- change in energy = heat + work
- if q is positive, heat is added to the system
- if w is positive, work is done on the system

Work is the force applied to an object as it moves a certain distance
- work = force x distance moved
Can also be found over a given area
- work = pressure x area x volume change

First law of thermodynamics
- ΔE = qp - PΔV
- heat, q, have the subscipt, p, to indicate the pressure must be held constant
When heat energy is measure in a calorimeter that does not allow the volume to change, PΔV is zero, resulting in ΔE = qv, where the volume is held constant
- Calorimeters that do not let volume change are called bomb calorimeters
Enthalpy change is ΔH where H is the heat content of the substance and the Δ is the difference
- ΔH = Hproducts - Hreactants
ΔE = ΔH - PΔV

ΔH is an extensive property and to make it intensive, the amount of chemical reacting must be specified
Standard heat of reaction, ΔH°, is the heat produced when the number of moles in the balanced equation reacts. in kJ.
If a formation reaction occurs, ΔH°f, where only one mole of product is the equation, the units will be kJ/mol

When converting coefficients in a balanced equation by multiplying, dividing, or any mathematic conversion, that same conversion must be done to the heat of reaction
Summarized law
- If the coefficients are multiplied by a constant, that ΔH°reaction is multiplied by the same constant
- If two or more equations are added to create an overall reaction, the heats of the equations can be added to obtain the overall reaction heat
When trying to use Hess’s Law, follow the principles:
- Focus on complex molecules first
- focus on atoms and molecules that occur in just one reaction
- Focus on atom and molecules that are in the overall reaction
- Focus on finding atoms and molecules to cancel out unneeded ones from equations

Formation reactions are when reactants are in their standard state and there is only one mole of product. Examples:
- Fe (s) + ½O2 (g) → FeO (s)
- 2K (s) + ½H2 + 2O2 (g) + P (s) → Fe2O3 (s)
Heat of formation for elements is always zero

Entropy is the number of different ways a system can arrange the particles
- Increasing number of particles increases entropy
- Increasing volume increases entropy
- Changing state (liquid to gas) increasing entropy
Ice has the least entropy while gaseous water has the most
Entropy is assigned the symbol S with the units J/C
Standard entropy is S° with the units J/C mol
Entropy principles
- Gas formation increases entropy greatly. A higher Δng means a greater entropy
- If Δng = 0, solid to liquid phase is a high contributor to entropy increase
- solid formation decreases entropy
- increase in temperature increases entropy

Free energy change, ΔG°, is the max amount of energry available from any chemical reaction
- Driven by entropy and enthalpy values
ΔG° = ΔH° - TΔS°
Second law of thermodynamics is any physical or chemical change must result in the increase of entropy in the universe

ΔH°	ΔS°	ΔG° as T increases	Comment
Negative	Positive	Always negative	Always spontaneous
Positive	Negative	Always positive	Never spontaneous
Negative	Negative	Become positive	Becomes nonspontaneous as T increases
Positive	Positive	Becomes negative	Becomes spontaneous as T increases