Lectures 24-33 Chem II

Spontaneous process

a process that occurs without intervention

Spontaneity is…

independent of the speed or rate of reaction; a spontaneous process proceeds slowly

Factors of spontaneity

• Energy change

• Temperature ΔH (measures the total heat content of a system)

• Entropy change

Gibbs Free energy equation with entropy

ΔG= ΔH-TΔS

negative ΔG means…

spontaneous

positive ΔG means…

non-spontaneous

Entropy

measures the disorder of a system

• Greater disorder = greater entropy

ΔS ΔH and ΔS -ΔH

spontaneous at high temp. and spontaneous at all temp.

-ΔS ΔH and -ΔS -ΔH

non-spontaneous at any temp. and spontaneous at low temp.

First Law of Thermodynamics

energy of the universe is constant, but the various forms of energy can be interchanged in physical and chemical processes

Second Law of Thermodynamics

in any spontaneous process, there is always an increase in entropy

Third Law of Thermodynamics

a perfect crystal at absolute zero (0 Kelvin) has an entropy value of zero

S (entropy) equation

S = k*ln(W), where k is the Boltzmann constant (1.38E-23 J/K) and W is the number of microstates (energetically equivalent ways to arrange components in a system)

Entropy of gas, liquid, solid (allotropes- different forms of the same element)

S_gas > S_liquid > S_solid

Entropy of pure substance vs mixture

S_mixture > S_{pure substance}

Entropy with temperature

S_{high temperature} > S_{low temperature}

Entropy with mass

S_{increasing mass} > S_{decreasing mass}

More complex structures have

have greater entropy

Equation of entropy with universe

S_universe = S_sys + S_surroundings

▪ System: the reaction

▪ Surroundings: everything else

S(system) S(surroundings) S(universe)

Spontaneous

-S(system) -S(surroundings) -S(universe)

Non-spontaneous (reaction will proceed in reverse)

S(system) -S(surroundings) S(universe)?

Spontaneous if system is larger in magnitude than surroundings

-S(system) S(surroundings) S(universe)?

Spontaneous if surroundings are larger in magnitude than system

Tabulated standard entropy ∆S° values

Standard state = at 25 degrees Celsius (298K), 1M, 1 atm, etc.

Standard entropy equation and example

∆S°(rxn) = Σn∆S°(products) - Σm∆S°(reactants)

Change in entropy of the surroundings is and equation…

directly proportional to the enthalpy of the system

• ∆Ssurr = -∆Hsys/T

Exothermic process corresponds to

positive entropy change in surroundings

What is a major difference between chemical kinetics and thermodynamics?

Chemical kinetics focuses on the pathway between reactants and products (path-dependent); thermodynamics only considers initial and final states (state functions).

What do you notice about entropy values for elements and compounds?

Entropy values for elements are much smaller than entropy values for compounds, since compounds typically have more microstates.

What is the difference between ∆S(universe) in spontaneous and equilibrium processes?

∆S(universe) > 0 for spontaneous processes, but ∆S = 0 in equilibrium processes.

Define Gibbs Free Energy.

Gibbs Free Energy: the maximum amount of non-expansion work that can be extracted from a closed system

∆G°(rxn) equation

∆G°(rxn)= Σn∆G°f(products) - Σm∆G°f(reactants) at 25 degrees Celsius

∆G°(rxn)

is a measure of the driving force

The free energy of formation of pure elements in their standard states…

is zero

Negative ∆G(rxn) indicates…

forward reaction has excess energy and will occur spontaneously

Positive ∆G(rxn) indicates…

there is not enough energy in the forward direction, so the backward reaction will occur

∆G(rxn) equation

= ∆G°(rxn) + RTln(K)

∆G(rxn) at equilibrium

= 0 at equilibrium; ∆G°(rxn) = -RTln(K) at equilibrium

Hess’s Law for ∆G.

• If a reaction is reversed, the sign of its ∆G value reverses

• If the amount of materials is multiplied by a factor, the value of the ∆G is multiplied by the same factor

∆G under nonstandard conditions equation

∆G = ∆G° + RTln(Q) (yes, equilibrium never leaves)

▪ Remember Q and K are calculated the same way – [products]/[reactants]

The larger the value of K (equilibrium constant)

the more negative the value of ∆G°

If ∆G° = -RTln(K), then you should be able to solve for K using algebra

▪ K = e^(-∆G°/RT), where K can be any equilibrium constant (Keq, Ksp, etc.)

Reactant-favored vs product-favored

+∆G° vs -∆G°

Favorable entropy is

is positive; unfavorable entropy is negative

Favorable enthalpy is

is negative; unfavorable enthalpy is positive

Gibbs free energy with entropy and enthalpy equation

∆G = ∆H-T∆S

Determine the minimal temperature for spontaneity, which means the “threshold” temperature where a reaction becomes spontaneous

T = ∆H/∆S you can derive this; prove this to yourself (hint: what does ∆G have to be for a reaction to be spontaneous?)

Why are most real reactions irreversible?

In a real reaction, some (if not most) of the free energy is “lost” as heat.

Electrochemistry:

the study of the interchange between chemical change and electrical work

Electrochemical cells

systems utilizing redox reactions to produce or use electrical energy

Oxidation

loss of electrons; Reducing Agent: the species that gets oxidized

▪ Electron(s) is a product

Reduction

gain of electrons; Oxidizing Agent: the species that gets reduced

▪ Electron(s) is a reactant

Electrons are negatively charged, so if you gain electrons (reduction)…

The oxidation number becomes more negative; if you lose electrons (oxidation), the oxidation number becomes more positive

Cathode

Where oxidation occurs

Where electrons are produced

Anions migrate towards

Has a negative sign

Anode

Where reduction occurs

Where electrons are consumed

Cations migrate towards

Has a positive sign

How to write half-reactions

• Electrons remain constant

How to write net reaction

• No electrons here

Physical meaning of an oxidation number

• Indicates a degree of oxidation and reduction

• Tracks electron loss (oxidation) or gain (reduction) during chemical reactions, particularly in covalent compounds where actual charge transfer does not occur

Oxidation number rules (that are a hierarchy!).

• Pure element has an oxidation number of zero

• Monatomic ion has an oxidation number equal to the charge on the ion

• Polyatomic ions must have the sum of the oxidation numbers equal to the charge on the whole ion

• Hydrogen = +1 (unless a hydride- NaH, LiH, CH4, LiAlH4)

• Oxygen = -2 (unless peroxide- H2O2, Na2O2, BaO2)

• Fluorine = -1

Active Metals (Group 1 and Group 2)

Lose electrons easily

Easily oxidized

Strong reducing agents

Active nonmetals (Halogens in Group 17)

Gain electrons easily

Easily reduced

Strong oxidizing agents