chemistry 123: physical and organic chemistry · 3/22/09 3 chemistry 123: physical and organic...

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Chemistry 123: Physical and Organic Chemistry Topic 4: Gaseous Equilibrium http://people.ok.ubc.ca/orcac/chem123out.html

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Topic 4: Gaseous Equilibrium Text: Chapter 6 & 15 4.0 Brief review of Kinetic theory of gasses (Chapter 6) 4.1 Concept of dynamic equilibrium 4.2 General form & properties of equilbrium const. and reaction quotient 4.3 Extent of reaction at equilibrium; degree of dissociation 4.4 Variation of equilibrium constant with temperature 4.5 Le Chatelier's Principal and it's applications

Topic 4: Introduction,


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Topic 4.0: Kinetic Theory Kinetic Molecular Theory

Particles are point masses in constant, random, straight line motion.

Particles are separated by great

distances.

Collisions are rapid and elastic.

No force between particles.

Total energy remains constant.

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Topic 4.0: Kinetic Theory Pressure – Assessing Collision Forces

Translational kinetic energy,

Frequency of collisions,

Impulse or momentum transfer,

Pressure proportional to impulse times frequency


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Topic 4.0: Kinetic Theory Pressure and Molecular Speed

Three dimensional systems lead to:

um is the modal speed uav is the simple average urms

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Topic 4.0: Kinetic Theory Pressure

Assume one mole:

PV=RT so:

NAm = M:

Rearrange:


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Topic 4.0: Kinetic Theory Distribution of Molecular Speeds

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Topic 4.0: Kinetic Theory Temperature

Modify:

PV=RT so:

Solve for ek:

Average kinetic energy is directly proportional to temperature!


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Topic 4.0: Equilibrium Dynamic Equilibrium

♦  Equilibrium – two opposing processes taking place at equal rates.

H2O(l) H2O(g)

I2(H2O) I2(CCl4)

NaCl(s) NaCl(aq) H2O

CO(g) + 2 H2(g) CH3OH(g)

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Topic 4.0: Equilibrium The Equilibrium Constant Expression ♦ Methanol synthesis is a reversible reaction.

CO(g) + 2 H2(g) CH3OH(g) k1

k-1 CH3OH(g) CO(g) + 2 H2(g)


k-1


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Topic 4.0: Equilibrium Three Approaches to the Equilibrium

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Topic 4.0: Equilibrium


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k-1

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Forward: CO(g) + 2 H2(g) → CH3OH(g)

Reverse: CH3OH(g) → CO(g) + 2 H2(g)

At Equilibrium:

Rfwrd = k1[CO][H2]2

Rrvrs = k-1[CH3OH]

Rfwrd = Rrvrs k1[CO][H2]2 = k-1[CH3OH]

[CH3OH] [CO][H2]2

=k1 k-1

= Kc

k1

k-1


k-1


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Topic 4.0: Equilibrium Activity

•  Thermodynamic concept introduced by Lewis. •  Dimensionless ratio referred to a chosen reference

state.

aB = [B] cB0

cB0 is a standard reference state = 1 mol L-1 (ideal conditions)

= γB[B]

• Accounts for non-ideal behaviour in solutions and gases. • An effective concentration.

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Topic 4.0: Kinetic Theory

a A + b B …. → g G + h H ….

Equilibrium constant = Kc= [A]m[B]n …. [G]g[H]h ….

Thermodynamic Equilibrium constant = Keq=

(aG)g(aH)h ….

(aA)a(aB)b ….

[A]m[B]n …. [G]g[H]h …. ( γG)g(γH)h ….

(γA)a(γB)b …. × =

≈ 1 under ideal conditions


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Topic 4.0: Kinetic Theory Relationships Involving the Equilibrium Constant •  Reversing an equation causes inversion of K. •  Multiplying by coefficients by a common factor raises the equilibrium constant to the corresponding power. •  Dividing the coefficients by a common factor causes the equilibrium constant to be taken to that root.

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Topic 4.0: Kinetic Theory Gases: The Equilibrium Constant, KP

Mixtures of gases are solutions just as liquids are. Use KP, based upon activities of gases.


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Topic 4.0: Kinetic Theory Gases: The Equilibrium Constant, KC

In concentration we can do another substitution

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Topic 4.0: Kinetic Theory


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Topic 4.0: Equilibrium CaCO3(s) CaO(s) + CO2(g)

Kc = [CO2] KP = PCO2(RT)

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Topic 4.0: Equilibrium The Reaction Quotient, Q:


k-1 Equilibrium can be approached various ways. Qualitative determination of change of initial conditions as

equilibrium is approached is needed.

At equilibrium Qc = Kc Qc = [A]tm[B]tn [G]tg[H]th

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Topic 4.0: Equilibrium Altering Equilibrium Conditions: Le Châtelier’s Principle

When an equilibrium system is subjected to a change in temperature, pressure, or concentration of a reacting species, the system responds by attaining a new equilibrium that partially offsets the impact of the change.

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Le Châtelier’s Principle

Q = = Kc [SO2]2[O2] [SO3]2

Q > Kc

Kc = 2.8×102 at 1000K 2 SO2(g) + O2(g) 2 SO3(g) k1

k-1


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• Adding a gaseous reactant or product changes Pgas. • Adding an inert gas changes the total pressure.

• Relative partial pressures are unchanged. • Changing the volume of the system causes a change in the equilibrium position.

•  When the volume of an equilibrium mixture of gases is reduced, a net change occurs in the direction that produces fewer moles of gas.

•  When the volume is increased, a net change occurs in the direction that produces more moles of gas.

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Topic 4.0: Equilibrium Effect of Temperature on Equilibrium

• Raising the temperature of an equilibrium mixture shifts the equilibrium in the direction of the endothermic reaction. • Lowering the temperature causes a shift in the direction of the exothermic reaction.


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Topic 4.0: Equilibrium Effect of a Catalyst on Equilibrium

• A catalyst changes the mechanism of a reaction to one with a lower activation energy. • A catalyst has no effect on the condition of equilibrium. • But does affect the rate at which equilibrium is attained.

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Topic 4.0: Equilibrium Nitrogen Cycle and the Synthesis of Nitrogen Compounds.

N2(g) + O2(g) NO(g) k1

k-1 KP = 4.7 10-31 at 298K and

1.3 x 10-4 at 1800K

chemistry 123: physical and organic chemistry · 3/22/09 3 chemistry 123: physical and organic...

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