Workshop Chem 201

Thermodynamics: general concepts of energy, entropy, enthalpy and spontaneity of reactions

Peer mentor: Gustavo Rodrigues de Moraes Sessions: Monday & Friday 2 – 4 PM [email protected] Discord: https://discord.com/invite/7C7Aew Zoom room: https://cccconfer.zoom.us/j/98202589993 STEM Center: tinyurl.com/Mesa-STEM-Center Heat – Chapter 6

• First Law of Thermodynamics: The energy of the universe is constant.

∆E = q + w where, ∆E = change in the system’s internal energy, Measuring heat! q = heat, w = work Enthalpy

• What is it? Enthalpy, the sum of the internal energy and the product of the pressure and volume of a thermodynamic system. Enthalpy is an energy-like property or state function.

Enthalpy is defined by the letter H We will basically study reactions under constant pressure and constant volume H = E + PV E = internal energy of the system P = pressure of the system V = volume of the system Enthalpy

• Thus, ∆H of the system is equal to the energy flow as heat…

∆H = Hproducts - Hreactants

• When Hproducts > Hreactants, ΔH is positive: heat is absorbed by the system, and the reaction is endothermic.

• When Hproducts < Hreactants, ΔH is negative: overall decrease in enthalpy is achieved by the generation of heat, and the reaction is exothermic ∆E = (−) Exothermic reaction ∆E = (+) Endothermic reaction Heat is increasing in the surroundings Heat is increasing in the system Increasing number of microstates Decreasing number of microstates Hess’s Law

• Hess’s Law stated that enthalpy is the same whether the reaction takes place in one step or in a series of steps.

∆H1 = ∆H2 + ∆H3

Practice time

1. Calculate the enthalpy of the following chemical reaction:

CS2(ℓ) + 3O2(g) ---> CO2(g) + 2SO2(g)

Given:

C(s) + O2(g) ---> CO2(g) ΔH = −393.5 kJ/mol

S(s) + O2(g) ---> SO2(g) ΔH = −296.8 kJ/mol

C(s) + 2S(s) ---> CS2(ℓ) ΔH = +87.9 kJ/mol

2. Calculate the enthalpy for this reaction given the following thermochemical equations:

2C(s) + H2(g) ---> C2H2(g) ΔH° = ??? kJ

5 C2H2(g) + ⁄2O2(g) ---> 2CO2(g) + H2O(ℓ) ΔH° = −1299.5 kJ

C(s) + O2(g) ---> CO2(g) ΔH° = −393.5 kJ 1 H2(g) + ⁄2O2(g) ---> H2O(ℓ) ΔH° = −285.8 kJ

3. If C2H5OH + 3 O2(g) à 2 CO2(g) + 3 H2O(l) ∆H = -1418 kJ/mol

How much heat is released when 10.00g of ethanol is burned? Entropy (∆S)

• What is it? Entropy is the measure of a system’s thermal energy per unit temperature that is unavailable for doing useful work (Britannica) . Entropy can be viewed as a measure of molecular randomness or disorder.

• The Second Law of Thermodynamics: In any spontaneous process there is always an increase in the entropy of the universe. Entropy

∆Suniv = ∆Ssys + ∆Ssurr

ΔSuniv is positive: • Entropy of the universe increases • Process is spontaneous in the direction written

ΔSuniv is negative: • Process is spontaneous in the opposite direction

ΔSuniv is zero: • Process has no tendency to occur • System is at equilibrium Spontaneity

• Processes are spontaneous when they result in an increase in disorder. • Entropy changes in the surroundings are determined primarily by the flow of energy into or out of the system as heat.

• ∆Ssurr = (+) energy is being added up to the surroundings. Reaction is exothermic/ exergonic. Surroundings: greater disorder.

• ∆Ssurr = (−) energy is being absorbed from the surroundings. Reaction is endothermic/ endergonic. Surroundings: more organized.

H2O1(l) à H2O(g) 18 mL of liquid water becoming to 31 L of gaseous water • A system tends to undergo changes that lower its energy because when a system at constant temperature moves to a lower energy, the energy it gives up is transferred to the surroundings, leading to an increase in entropy there. • Temperature has a huge impact on spontaneity: • Low temperatures: energy transfer plays an important role • High temperature: energy transfer is not so important

∆� ∆� = − ���� �

4. Arrange the compounds in order of increasing entropy (∆S) and justify your order. a) CH3OH(s), b) CH3OH(l), c) CH3OH(g), d) CH3CH2OH(g)

5. Predict the sign of ∆S∘ for each of the following reactions. a. The thermal decomposition of solid calcium carbonate:

CaCO3(s) à CaO(s) + CO2(g) b. The oxidation of SO2 in air:

2SO2(g) + O2(g) à 2SO3(g) Free Energy (∆G)

∆G = ∆H − (T × ∆S) ∆G = variation in free energy ∆H = variation in enthalpy T = temperature in Kelvins ∆S = variation in entropy ∆ ∆S =− univ

Reactions are spontaneous when ∆G is negative. Since energy is being given off. Gibbs equation

∆ ( ) • ��� = �

Keq > 1 Keq ≈1 Keq < 1 Product favored Equilibrium Reactant favored ∆G∘ = (−) ∆G∘ = 0 ∆G∘ = (+) Free energy is released Reversible Free energy is required Exergonic Endergonic Ex: batteries, combustion Ex: extremely weak acids, extremely insoluble salts

• ∆� = � × � × ��� Thermo scenarios Practice time

o 6. Using the thermodynamic tables, calculate ΔH rxn for each chemical reaction. a. 2Mg(s) + O2(g) → 2MgO(s) b. CaCO3(s) → CaO(s) + CO2(g) c. AgNO3(s) + NaCl(s) → AgCl(s) + NaNO3(s) 7. At what temperature would the following reaction occur spontaneously?

2Cr2O3→ 4Cr + 3O2 (∆H°= 1256.4 kJ; ∆S°= 587 J/K)

8. Given the dissolving of CaCl2(s)

2+ − CaCl2(s)→Ca (aq)+2Cl (aq) ΔG° = CaCl2(s) = -748.1 kJ/mol, ΔH = CaCl2= -795.8 kJ ∆� − ∆�° o � = Δs = � • Ca2+= -53.1 J/Kmol, • CaCl2= 104.6 J/Kmol, • Cl- = 56.5 J/Kmol

Calculate the temperature of this reaction.

9. Calculate ΔGo using only the following ΔG values for the reaction:

CS2(l)+2O2(g)→CO2(g)+2SO2(g)

o C(s)+O2(g)→CO2(g) with ∆G =−394.39kJ/mol o S(s)+O2(g)→SO2(g) with ∆G = −300.13kJ/mol o C(s)+2S(s)→CS2(l) with ∆G = 67.1kJ/mol o o 10. Assess the feasibility of the reaction below by evaluating (∆S rxn, ∆H rxn o and ∆G rxn) for this reaction at 25°C. If the reaction is spontaneous, is it enthalpically favored, entropically favored or both?

(NH)4Cl(s)→NH3(g)+HCl(g)

11. Determine the Kp at 298.15 K for the reaction:

2CO(g) + O2(g) ⇌ 2CO2(g) 12. ∆G° for the reaction below is 31.4 kJ/mol at 298.15 K. Use the thermodynamics quantity to decide in which direction the reaction is spontaneous when the concentrations of − + H2SO4(aq), HSO 4(aq), and H3O (aq) are 0.2 M, 0.1 M, and 0.1 M respectively.

+ − H2SO4(aq) + H2O(l) ⇌ H3O (aq) + HSO 4(aq)

Useful resources: • Thermo videos: https://ch301.cm.utexas.edu/thermo/index.php

Exercises: https://www2.southeastern.edu/Academics/Faculty/wparkinson/help/ enthalpy/test.html Copyright

• https://www.britannica.com/science/enthalpy • https://www.britannica.com/science/entropy-physics • https://www.chemteam.info/Thermochem/HessLawIntro1a.html • https://chem.libretexts.org/Bookshelves/General_Chemistry/Map%3A_Ch emistry_- _The_Central_Science_(Brown_et_al.)/05._Thermochemistry/5.E%3A_The rmochemistry_(Exercises) • https://chem.libretexts.org/Courses/Mount_Royal_University/Chem_1202 /Unit_7%3A_Principles_of_Thermodynamics/7.E%3A_Exercises_on_Entrop y_and_Gibbs_Energy • Zumdahl, Steven S.. Chemistry (p. 225). Cengage Learning. Kindle Edition.