Spontaneous Processes, Entropy and Thermodynamics

Spontaneous Process

• It's a physical or chemical change that occurs by itself
• This process happens without needing energy from an external source
• Example 1: Methane combustion (CH4 + 2O2 → 6CO2 + 2H2O)
• Example 2: Ammonium Nitrate dissolution in water (NH4NO3 (s) → NH4+ (aq) +NO3- (aq))

Entropy

• Entropy measures how spread out or dispersed the energy of a system is
• It indicates randomness and disorder
• There is increasing randomness of the universe
• The SI unit for entropy is Joules per Kelvin (J/K)

Entropy in Phase Changes and Dissolving

• Entropy increases when a substance changes from solid to liquid during melting
• Entropy increases when a substance changes from liquid to vapor during vaporization
• Entropy increases when a solute dissolves into a solution

Second Law of Thermodynamics

• This law states the universe's entropy increases in a spontaneous process
• Entropy remains unchanged during equilibrium
• Mathematically, it means ∆Suniv = ∆Ssys + ∆Ssur > 0

Entropy Types

• Positive Entropy occurs during melting, evaporation, and sublimation
• Entropy change is affected by the production of more gas molecules
• Change in entropy may be positive or negative

Negative Entropy

• Negative entropy occurs during deposition, solidification, and condensation
• It also occurs when the number of gas molecules decreases

Entropy Examples

• Melting a solid has positive entropy
• Freezing a liquid has negative entropy
• Converting a vapor into solid has negative entropy
• (H_2(g) + Br_2(g) \rightarrow 2HBr(g)) has positive entropy.
• (H_2(g) + Br_2(g) \rightarrow 2HBr(g)) has negative entropy.

Calculating Entropy Changes

• Use the standard entropy of reaction equation: ∆S° = Σ nS° (products) − Σ nS° (reactants)
• Variables m and n are the stoichiometric coefficients

Example 1 of Entropy Calculation

• Reaction example: H2 (g) + I2 (s) → 2HI (g)
• S° values: H2(g) = 130.6 J/K⋅mol, I2(s) = 116.7 J/K⋅mol, HI(g) = 206.3 J/K⋅mol
• Calculation: ∆S° = +165.3 J/K

Example 2 of Entropy Calculation

• Reaction example: N2 (g) + 3 H2 (G) → 2 NH3 (g)
• Calculate using the equation: ∆S° = -199 J/K

Josiah William Gibb

• American scientist famous for development of CHEMICAL THERMODYNAMICS
• He combined concepts of entropy and enthalpy
• He coined the term “Free energy,” defined as energy available to do WORK

Gibb's Free Energy

• It is used to predict spontaneity
• G = H – TS, where G is Gibb's Free Energy, H is enthalpy, T is temperature, and S is entropy
• For constant temperature, the change in free energy is: ΔG = ΔH – TΔS
• ΔH represents change of enthalpy, ΔS is change of entropy, T is standard temperature (25°C)

Spontaneity Conditions

• ΔG < 0 indicates the reaction is spontaneous in the forward direction
• ΔG > 0 indicates the reaction is non-spontaneous
• ΔG = 0 indicates the system is at equilibrium, with no net change

Calculating ∆H and ∆S

• Calculations at 25°C help determine reaction characteristics

Example 1 of Enthalpy Calculation

• For the reaction: CaCO3(s) → CaO (s) + CO2 (g)
• Given H° values: CaCO3(s) = -1206.9 kJ/mol, CaO (s) = -635.6 kJ/mol, CO2 (g) = -393.5 kJ/mol
• ΔH°rxn = 177.8 kJ/mol

Example 2 of Entropy Calculation

• For the reaction: CaCO3(s) → CaO (s) + CO2 (g)
• Given S° values: CaCO3(s) = 92.9 kJ/mol, CaO (s) = 39.8 kJ/mol, CO2 (g) = 213.6 kJ/mol
• ΔS°rxn = 160.5 kJ/mol

Reaction Characteristics Overview

Chemical Equilibrium

• Chemical equilibrium occurs when the rate of the forward and reverse reactions are equal
• Shown by: aA + bB ⇌ cC + dD
• The amounts of reactants and products remain constant

Chemical Equilibrium Dynamics

• Chemical reactions continue even at equilibrium
• Indicators of chemical reactions still taking place:
  • Color change
  • Gas formation
  • Formation of a Precipitate
  • Temperature variations

Reversible Reactions

• Reactants do not fully convert to products
• Reverse reactions may occur
• Claude Louis Berthollet first discovered this

Le Chatelier's Principle

• When stress is applied on system at equilibrium
• System adjusts to reduce stress effect
• The system will act to reestablish equilibrium

Concentration Changes

• Adding reactant shifts equilibrium to the right to produce more products
• Reducing product concentration also shifts equilibrium to the right
• Adding a product shifts equilibrium to the left, favoring reactant formation
• 2H2(g) + O2(g) ⇌ 2H2O(g) - BACKWARD

Pressure Changes

• Increasing pressure shifts the reaction towards the side with fewer gas molecule
• Consider: H2(g) + I2(g) ⇌ 2HI(g) - has no effect
• CaCO3(s) + I2(g) ⇌ CaO(s) + CO2(g) = increasing pressure causes a reverse reaction

Temperature Changes

• In Endothermic reactions: Heat is absorbed (ΔH is positive) , consider heat as a reactant
  • Increase of temperature will shift toward the production equilibrium to right, to counteract
  • N2O4 + heat ⇌ 2 NO2
  • Decrease of temperature will shift toward the production equilibrium to the left
  • N2O4 + heat ⇌ 2 NO2
• In Exothermic reactions: Heat is released (ΔH is negative), consider heat as a product
• Increase of temperature will shift toward the consumption equilibrium to the left, to counteract - N2O4 + heat ⇌ 2 NO2 - Decrease of temperature will shift toward the consumption equilibrium to the right - N2O4 + heat ⇌ 2 NO2

Acids and Bases

• Johannes Bronsted (1879-1947) & Thomas Lowry (1874-1936) created the Bronsted-Lowry Theory
• Acid strength relies on donating protons
• Base strength relies on accepting protons
• Bronsted-Lowry Acid-Base Reaction: a proton is transferred from an acid to a base
• Bronsted-Lowry Theory describes acid-base interactions in terms of proton transfer

Acid-Base Properties of Water

• Water can be both a Bronsted-Lowry acid and base and is amphoteric
• Water undergoes autoionization
• ionization constant of water is 1x10-14 at 25 degrees
• In ionization, [H3O+] and [OH-] concentrations are always equal
• Kw = [H3O+] [OH-] = 1x10-7

Bronsted-Lowry Acid

• A Bronsted-Lowry acid donates a proton (H+)
• H2O + NH3 → NH4+ + OH- (aq)

Bronsted-Lowry Base

• Bronsted-Lowry base accept a proton
• H2O(l) + NCl(aq) → H30+(aq) + CL- (aq)

Conjugate Base

• Conjugate base is a Bronsted-Lowry acid that forms after an acid donates a proton

Conjugate Acid

• Conjugate acid is a Bronsted-Lowry base that forms when a base accepts a proton
• A conjugate acid-base pair has identical molecular formula, except the acid gains an extra H+

The pH Scale

• The pH scale measures acidity or basicity in solutions
• Molar concentrations of H3O+ and OH⁻ in aqueous solutions are usually small
• pH is a more practical term to use
• pH is defined as the negative logarithm of [H3O+]
• pH less than 7 indicates acidity
• pH equal to 7 indicates neutrality
• pH greater than 7 indicates basicity

pH Calculations

• The pH can be caculated with the negative log of H3O+
• If the pOH of the solution is 4.5, the pH of the solution is 9.5
• If the POH is 3.8, then [OH-] = 1.58 x 10-4 M

Redox Reactions: Oxidation and Reduction

• These are reduction-oxidation reactions involving electron transfer
• Electron transfer relies on the reduction or oxidation potential of species
• The oxidation number changes by gaining or losing electrons

Oxidation Reactions

• Oxidation reactions are where electrons are lost by a molecule, atom, or ion (REDUCING AGENT)
• These reactions are when the oxidation state increases, they donate electrons (LEionOra)

Reduction Reactions

• Reduction reactions are where electrons are gained by a molecule, atom, or ion (OXIDIZING AGENT)
• These reactions are when the oxidation state decreases, they accept electrons (GEdonRoa)

Rules for Assigning Oxidation States

• Oxidation states of elements in their free, uncombined state is zero
• The oxidation state of hydrogen (H) in compounds is +1 (except in hydrides where it is -1)
• Oxidation state of oxygen (O) in compounds is -2 (except in peroxides where it is -1; and with fluorine where it is +2)
• Oxidation state of a representative element is equal to its valence
• The sum of oxidation states in a radical equals its charge
• The sum of oxidation states in a compound is zero

Balancing Redox Reactions: Example 1

• Assign oxidation numbers to all atoms
• Identify oxidized and reduced atoms, Mn reduced, Fe oxidized
• Use a bracketing line to connect atoms that undergo oxidation and reduction
• Write the oxidation-number change at the midpoint
• Multiply by factors to equate lost electrons with gained electrons
• Balance for atoms and moles (check the equation is balanced

Balancing Redox Reactions: Example 2

• Assign oxidation numbers to all atoms
• Identify which atoms are oxidized and which are reduced (Mn and Fe)
• Use a bracketing line to connect atoms that undergo oxidation and reduction
• Multiply by factors to equate lost electrons with gained electrons
• Balance for atoms and moles (check the equation is balanced)