 Chapter 16
Thermodynamics
 Part 1 Heat and Energy
 Heat is the quantity of thermal energy transferred from one substance to
another as a result of a difference in temperature.
Heat is measured in kJ and Joules
 Temperature is a measure of the average kinetic energy of the particles in
matter.
Temperature is measured in oC.
 Energy is the capacity to do work.
 Part 1 Heat and Energy
 As the temperature increases, the kinetic energy increases and the total
energy increases. Keep in mind that the total energy includes the number of
particles in the system.
example: a cup of water and a gallon of water
if at room temperature, which has more energy?
 Part 1 Heat and Energy
 The SI unit for energy is the joule (jewl) Most often we will use the kJ.
 Activity
 to a clean beaker add 100 ml of HCl and 100 ml of NaOH. Mix and record the
temperature.
 Predict the temperature change if you double the volume of both liquids.
 What did you discover?
 Specific Heat
 Specific heat is the measure of the amount of heat required to raise the
temperature of one gram of a substance one degree Celsius.
Unit are J/g x oC
Equation: q = m x C x Δ T
where q is heat, m is mass, C is specific heat and
ΔT is the change in temperature.
Practice Problems 1-3 page 536-537
 Specific heat lab( metals)
Practice Problems 4,5 page 539
Thermochemistry and Enthalpy
 Thermochemistry is the study of heat transfers that accompany chemical
reactions.
Easier to understand when you divide the universe into two parts. The system
and the surroundings.
Reactions that give off energy(heat) are said to be exothermic. Heat flows from
the system to the surroundings.
Reactions that require energy(heat) are said to be endothermic. Heat flows from
the surroundings in to the system.
Note: system + surroundings = universe
 Thermochemistry and Enthalpy
 Enthalpy(H) is a measure of the energy content of a system.
 C(s) + O2(g)  CO2(g) Energy released 393.5 kJ
 C(s) + O2(g)  CO2(g) ΔH = -393.5 kJ
Ba(OH)2 + 2NH4Cl BaCl2 +2NH4OH
ΔH = +130kJ
 Demo
 Enthalpy
 Standard Enthalpy of Formation is not an exact science. Rather it is like the oC
scale.
 “standard state” is at 25oC and one atm. Any change in either is a change in
enthalpy.
 So, what does this look like?
 ΔHfo where o is for standard stateand f is for formation of a mole
 See table 16-4 for standard enthalpies
 Enthalpy
 Determining Enthalpy Changes in Reactions
 Hess’s Law: When reactants are converted to products whether in one step
or more, the enthalpy change is the same.
 Like this: methane to water
 CH4 +2 O2  CO2 +2H2O
 C + 2H2 = CH4 -74.8
 C + O2  CO2 -393.5
 H2 + ½ O2  H2O -285.8
 Like this: methane to water
 CH4 +2 O2  CO2 +2H2O
 C + 2H2 = CH4 -74.8
 C + O2  CO2 -393.5
 H2 + ½ O2  2H2O -285.8
 Our equation is ΔH = Hproducts - Hreactants
 -393.5 + 2( -285.8) – (-74.8) = -890.3
 Better practice these.
 Pg 551 # 12 a, b, 17 page 553
 Part 3 Spontaneous Process
 Spontaneous reactions can occur, under specific conditions, without any
assistance.
 Burning Bunsen burner, melting of ice, food coloring in a glass all occur
spontaneously.
 Entropy and Spontaneity
 Entropy, S, is a measure of disorder. Nature likes disorder. Order requires
too much energy.
 Perfume in a room demonstrates entropy. It spread out, diffuses, on its
own(spontaneously.)
 Part 3 Spontaneous Process
 Factors that determine spontaneity
 ΔH (-) Exothermic tend to be spontaneous.
 ΔS (+) Increased disorder tend to be spontaneous
 ΔH (+) Endothermic tend to be nonspontaneous.
 ΔS (-) Increased order tend to be nonspontaneous.
 Confused???
 Spontaneous?
 ΔS is positive(increased disorder) when…
A solid changes into a liquid.
A liquid changes into a gas.
A solute dissolves in a solvent.
A system is heated.
A reaction increases the number of particles in the system. 2H2O  2H2O + O2
 Gibbs free energy
 Gibbs free energy, G, is a quantity that takes into account the contributions of
enthalpy and entropy when predicting if a reaction will take place.
 ΔG = Gfinal – Ginitial or ΔG = Gproducts– Greactants
Bringing it all together.
 ΔG = ΔH – TΔS
 a negative ΔG = spontaneous reaction
 Gibbs free energy
 To freeze one mole of water at -10oC you must remove 224J of energy or 224J. ΔG is negative as is ΔH so this occurs spontaneously.
 At temperatures greater than 0oC TΔS is greater than ΔH so ΔG is positive.
We know water won’t freeze spontaneously above 0oC so ΔG must be ___
 Gibbs free energy
 Exothermic reactions and reactions that increase entropy of a system
tend to be spontaneous. SO…, -ΔH, +ΔS and, -ΔG
 Table 16-6 is your friend. Practice #’s 20-22