Week 8.2 thermodynamics and equilibria

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Prepared by: Mrs Faraziehan Senusi PA-A11-7C Physical Transformation of Pure Substances Chemical Equilibrium Chapter 4 Thermodynamic and Equilibria First Law of Thermodynamics Reference: Chemistry: the Molecular Nature of Matter and Change, 6 th ed, 2011, Martin S. Silberberg, McGraw- Second Law of Thermodynamics Simple Mixtures

Transcript of Week 8.2 thermodynamics and equilibria

Page 1: Week 8.2   thermodynamics and equilibria

Prepared by:Mrs Faraziehan Senusi

PA-A11-7CPhysical Transformation of Pure Substances

Chemical Equilibrium

Chapter 4Thermodynamic and Equilibria

First Law of Thermodynamics

Reference: Chemistry: the Molecular Nature of Matter and Change, 6th ed, 2011, Martin S. Silberberg, McGraw-Hill

Second Law of Thermodynamics

Simple Mixtures

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THE SECOND LAW OF THERMODYNAMICS

• The second law of thermodynamics is a general principle which places constraints upon the direction of heat transfer and the attainable efficiencies of heat engines.

• The first law reflects the observation that energy is conserved, but it imposes no restriction on the process direction.

• The inadequacy of the first law to identify whether a process can take place is remedied by introducing the second law of thermodynamics.

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A process must satisfy both the first and second laws of thermodynamics to proceed.

A cup of hot coffee left on a table eventually cools off.First law: amount of energy lost by the coffee is equal to the amount gained by the surrounding air.BUT a cup of cool coffee in the same room never gets hot by itself. This process never takes place. Doing so would not violate the first law as long as the amount of energy lost by the air is equal to the amount gained by the coffee.

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Spontaneous process

• Predict whether or not a reaction will occur when reactants are brought together under a specific set of conditions (for example, at a certain temperature,pressure, and concentration).

• Reaction that does occur under the given set of conditions is called a spontaneous process

• Reaction that does NOT occur under specified conditions is nonspontaneous

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Example:• A waterfall runs downhill, but never up, spontaneously.• A lump of sugar spontaneously dissolves in a cup of coffee, but dissolved sugar

does not spontaneously reappear in its original form.• Water freezes spontaneously below 0°C, and ice melts spontaneously above 0°C

(at 1 atm).• Heat flows from a hotter object to a colder one, but the reverse never happens

spontaneously.

• Iron exposed to water and oxygen forms rust, but rust does not spontaneously

change back to iron.

These examples show that processes that occur spontaneously in one direction cannot, under the same conditions, also take place

spontaneously in the opposite direction.

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MAJOR USES OF THE SECOND LAW1. The second law may be used to identify the direction of

processes. • heat does not flow spontaneously from a cold to a hot body

2. The second law of thermodynamics is also used in determining the theoretical limits for the performance of commonly used engineering systems, such as heat engines and refrigerators, (Clausius Statement & Kelvin–Planck Statement)

• heat cannot be transformed completely into mechanical work• It is impossible to construct an operational perpetual motion

machine

3. In order to predict the spontaneity of a process, we need to introduce a new thermodynamic quantity called entropy.

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ENTROPY

The first law of thermodynamics deals with the property energy and the conservation of it.

The second law leads to the definition of a new property called entropy. Entropy is a measure of molecular disorder or molecular randomness. As a system becomes more disorder, the positions of the molecules

becomes less predictable and the entropy increases.

For any substance, the particles are more highly ordered in the solid state than in the liquid state. Liquid are more highly ordered than in the gaseous state. Thus, the entropy of any substance increases as the substance goes from solid to liquid to gas.

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Melting: Sliquid > Ssolid

Vaporization: Svapor > Sliquid

Dissolving: Ssolution > Ssolute

Heating: ST2 > ST1

In a solid, the atoms or molecules are confined to fixed positions and the number of microstates is small. Upon melting, these atoms or molecules can occupy many more positions as they move. Consequently, the number of microstates increases because there are now many more ways to arrange the particles. Therefore, we predict this “order disorder” phase transition to result in an increase in entropy because the number of microstates has increased.The vaporization process will also lead to an increase in the entropy of the system. The increase will be considerably greater than that for melting, however, because molecules in the gas phase occupy much more space, and therefore there are far more microstates than in the liquid phase.

The solution process usually leads to an increase in entropy. When a sugar crystal dissolves in water, the highly ordered structure of the solid and part of the ordered structure of water break down. Thus, the solution has a greater number of microstates than the pure solute and pure solvent combined.

Heating also increases the entropy of a system. As the temperature is increased, the energies associated with all types of molecular motion increase. This increase in energy is distributed or dispersed among the quantized energy levels. Consequently, more microstates become available at a higher temperature; therefore, the entropy of a system always increases with increasing temperature.

Microstates : each arrangement of the energy of each molecule in the whole system

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• Spontaneity of process predicted by change in entropy (S)

• Greater disorder therefore greater entropy

• Entropy state function

∆S = Sf – Si

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• The entropy of the universe increases in a spontaneous process and remains unchanged in an equilibrium process

Universe = system + surroundings

∆Suniverse = ∆Ssystem + ∆Ssurrounding

• For a spontaneous process:

∆Suniverse = ∆Ssystem + ∆Ssurrounding > 0

• For an equilibrium process:

∆Suniverse = ∆Ssystem + ∆Ssurrounding = 0

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• For a spontaneous process, the second law says that ΔSuniv must be greater than zero, but it does not place a restriction on either ΔSsys or ΔSsurr . Thus, it is possible for either ΔSsys or ΔSsurr to be negative, as long as the sum of these two quantities is greater than zero.

• For an equilibrium process, ΔSuniv is zero. In this case, ΔSsys and ΔSsurr must be equal in magnitude, but opposite in sign.

• What if for some hypothetical process we find that ΔSuniv is negative? What this means is that the process is not spontaneous in the direction described. Rather, it is spontaneous in the opposite direction.

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Entropy changes in the system

aA + bB cC + dD

The standard entropy of reaction ΔS°rxn is given by the difference in standard entropies between products and reactants:

∆Sorxn = [cSo

(C) + dSo (D)] – [aSo

(A) + bSo (B)]

∆Sosystem = ∆So

rxn = ∑n∆So (Product) – ∑m∆So (Reactant)

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Continued…

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Use the values of standard molar entropies in Appendix to calculate the entropy change at 25°C and one atmosphere pressure for the reaction of hydrazine with hydrogen peroxide.

The balanced equation for the reaction is

Example 7

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Entropy changes in the surrounding

• Exothermic process:

heat transferred to surrounding enhance motion of molecules in surrounding, increase disorder, increase entropy of surrounding

• Endothermic process:

heat absorbed from surrounding molecular motion decreases, decreases entropy surrounding

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• The minus sign is used because if the process is exothermic, ΔHsys is negative and ΔSsurr is a positive quantity, indicating an increase in entropy.

• On the other hand, for an endothermic process, ΔHsys is positive and the negative sign ensures that the entropy of the surroundings decreases.

The change in entropy of the surroundings, ΔSsurr is proportional to ΔHsys

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Example 8

Determine whether the synthesis of ammonia is spontaneous at 25oC.

N2 (g) + 3H2(g) 2NH3(g) ∆Ho= -92.6kJ/mol

Because ΔSuniv is positive, we predict that the reaction is spontaneous at 25°C.

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Third law of thermodynamics

The entropy of a pure, perfect crystalline substance (perfectly ordered) is zero at absolute zero (0 K).

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Gibbs Free Energy

• The second law of thermodynamics tells us that a spontaneous reaction increases the entropy of the universe; that is, ΔSuniv > 0.

• In order to determine the sign of ΔSuniv for a reaction, however, we would need to calculate both ΔSsys and ΔSsurr .

• In general, we are usually concerned only with what happens in a particular system.

• Therefore, we need another thermodynamic function to help us determine whether a reaction will occur spontaneously if we consider only the system itself.

• In order to express the spontaneity of a reaction more directly, we introduce another thermodynamic function called Gibbs free energy (G), or free energy.

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• Gibbs free energy, GG = H – TS

• Gibbs free energy change, ΔG,

• Free energy is the energy available to do work. • G is a state function.• The relationship between ∆G and spontaneity may be summarized

as follows:

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Standard free energy changes

aA + bB cC + dD

∆Gorxn = [c∆Gf

o (C) + d∆Gf

o (D)] – [a∆Gf

o (A) + b∆Gf

o (B)]

∆Gorxn =∑n∆Gf

o (Product) –∑m∆Gf

o(Reactant)

• ∆Gfo = standard free energy of formation

• ∆Gorxn is the free-energy change for a reaction when it occurs

under standard-state conditions, when reactants in their standard states are converted to products in their standard states.

• Free energy change that occurs when 1 mole of the compound is synthesized from its elements at standard states.

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Example 9

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Example 10 9

9

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• In order to predict the sign of ΔG, according to

ΔG = ΔH - TΔS

we need to know both ΔH and ΔS. • Temperature may also influence the direction of a

spontaneous reaction. If both ΔH and ΔS are positive, then ΔG will be negative only when the

TΔS term is greater in magnitude than ΔH. This condition is met when T is large.

If ΔH is positive and ΔS is negative, ΔG will always be positive, regardless of temperature.

If ΔH is negative and ΔS is positive, then ΔG will always be negative regardless of temperature.

If ΔH is negative and ΔS is negative, then ΔG will be negative only when TΔS is smaller in magnitude than ΔH. This condition is met when T is small.

The Temperature Dependence of Spontaneity

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Free energy and chemical equilibrium

• As mentioned earlier, during the course of a chemical reaction not all the reactants and products will be at their standard states.

• Start reaction at standard state and when reaction progress, then no longer at standard state

• Therefore use ∆G instead of ∆Go to predict reaction. • Under this condition, the relationship between ΔG and

ΔG°, which can be derived from thermodynamics, is

where R is the gas constant (8.314 J/K.mol), T is the absolute temperature of the reaction, and Q is the reaction quotient

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• At equilibrium, by definition, ΔG = 0 and Q = K, where K is the equilibrium constant. Thus,

KP is used for gases Kc for reactions in solution

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Example 11• Calculate the equilibrium constant for the following reaction

at 25oC

∆Gfo(H2O)= -237.2kJ/mol

This extremely small equilibrium constant is consistent with the fact that water does not spontaneously decompose into hydrogen and oxygen gases at 25°C. Thus, a large positive ΔG° favors reactants over products at equilibrium.