Chemical Equilibrium - libvolume4.xyzlibvolume4.xyz/.../chemicalequilibriumpresentation2.pdf · Equilibrium Equilibrium is a state in which there are no observable changes as time

Chemical Equilibrium

Chapter 14

14.1-14.5

Equilibrium

Equilibrium is a state in which there are no observable changes as time goes by.

Chemical equilibrium is achieved when:

1.) the rates of the forward and reverse reactions are equal and

2.) the concentrations of the reactants and products remain constant

Equilibrium

• There are two types of equilibrium:

Physical and Chemical.

– Physical Equilibrium

• H20 (l) ↔ H20 (g)

– Chemical Equilibrium

• N2O4 (g) ↔ 2NO2

Physical Equilibrium


Chemical EquilibriumN2O4 (g) ↔ 2NO2 (g)

Law of Mass Action

• Law of Mass Action- For a reversible reaction

at equilibrium and constant temperature, a

certain ratio of reactant and product

concentrations has a constant value (K).

• The Equilibrium Constant (K)- A number equal

to the ratio of the equilibrium concentrations of

products to the equilibrium concentrations of

reactants each raised to the power of its

stoichiometric coefficient.

Law of Mass Action

• For the general reaction:

K = [C]c[D]d

[A]a[B]b

aA (g) + bB (g) cC (g) + dD (g)

Equilibrium Constant

N2O4 (g) ↔ 2NO2 (g)


• Chemical equilibrium is defined by K.

• The magnitude of K will tell us if the

equilibrium reaction favors the reactants or

the products.

• If K » 1::..favors products

• If K « 1::..favors reactants

Equilibrium Constant Expressions

• Equilibrium constants can be expressed

using Kc or Kp.

• Kc uses the concentration of reactants and

products to calculate the eq. constant.

• Kp uses the pressure of the gaseous

reactants and products to calculate the eq.

constant.


• Equilibrium Constant Equations

Kc = [NO2]

2

[N2O4]Kp =

NO2P2

N2O4P

aA (g) + bB (g) cC (g) + dD (g)

Homogeneous Equilibrium

• Homogeneous Equilibrium- applies to

reactions in which all reacting species are

in the same phase.

N2O4 (g) ↔ 2NO2 (g)

Kp = NO2P2

N2O4P

In most cases

Kc ≠ Kp

Kc = [NO2]

2

[N2O4]


• Relationship between Kc and Kp

Kp = Kc(RT)∆n

∆n = moles of gaseous products – moles of gaseous reactants

= (c + d) – (a + b)

Equilibrium Constant Calculations

The equilibrium concentrations for the reaction between carbon monoxide and molecular chlorine to form COCl2 (g) at 740C are [CO] = 0.012 M, [Cl2] =

0.054 M, and [COCl2] = 0.14 M. Calculate the equilibrium constants Kc and Kp.

CO (g) + Cl2 (g) COCl2 (g)

Kc = [COCl2]

[CO][Cl2]=

0.14

0.012 x 0.054= 220

Kp = Kc(RT)∆n

∆n = 1 – 2 = -1 R = 0.0821 T = 273 + 74 = 347 K

Kp = 220 x (0.0821 x 347)-1 = 7.7


• The equilibrium constant Kp for the reaction is 158 at

1000K. What is the equilibrium pressure of O2 if the PNO

= 0.400 atm and PNO = 0.270 atm?

Kp = 2

PNO PO2

PNO2

2

PO2 = KpPNO

2

2

PNO2

PO2= 158 x (0.400)2/(0.270)2 = 347 atm

Heterogeneous Equilibrium

• Heterogeneous Equilibrium- results from a reversible reaction involving reactants and products that are in different phases.

• Can include liquids, gases and solids as either reactants or products.

• Equilibrium expression is the same as that for a homogeneous equilibrium.

• Omit pure liquids and solids from the equilibrium constant expressions.


Constant

CaCO3 (s) CaO (s) + CO2 (g)

[CaCO3] = constant

[CaO] = constantKp = PCO2

The concentration of solids and pure liquids are not

included in the expression for the equilibrium constant.


Constant


Consider the following equilibrium at 295 K:

The partial pressure of each gas is 0.265 atm. Calculate Kpand Kc for the reaction.

NH4HS (s) NH3 (g) + H2S (g)

Kp = PNH3 H2SP = 0.265 x 0.265 = 0.0702

Kp = Kc(RT)∆n

Kc = Kp(RT)-∆n ∆n = 2 – 0 = 2 T = 295 K

Kc = 0.0702 x (0.0821 x 295)-2 = 1.20 x 10-4

Multiple Equilibria

• Multiple Equilibria- Product molecules of one equilibrium

constant are involved in a second equilibrium process.

A + B C + D

C + D E + F

A + B E + F

Kc‘

Kc‘‘Kc

Kc = Kc‘‘Kc‘ x

Kc =‘[C][D]

[A][B]

Kc =‘‘[E][F]

[C][D]

[E][F]

[A][B]Kc =

Writing Equilibrium Constant Expressions

• The concentrations of the reacting species in the condensed phase are expressed in M. In the gaseous phase, the concentrations can be expressed in

M or in atm.

• The concentrations of pure solids, pure liquids and solvents do not appear in

the equilibrium constant expressions.

• The equilibrium constant is a dimensionless quantity.

• In quoting a value for the equilibrium constant, you must specify the balanced equation and the temperature.

• If a reaction can be expressed as a sum of two or more reactions, the equilibrium constant for the overall reaction is given by the product of the

equilibrium constants of the individual reactions.

14.2

What does the Equilibrium

Constant tell us?

• We can:

– Predict the direction in which a reaction

mixture will proceed to reach equilibrium

– Calculate the concentration of reactants and

products once equilibrium has been reached

Predicting the Direction of a

Reaction

• The Kc for hydrogen iodide in the following equation is

53.4 at 430ºC. Suppose we add 0.243 mol H2, 0.146 mol

I2 and 1.98 mol HI to a 1.00L container at 430ºC. Will

there be a net reaction to form more H2 and I2 or HI?

H2 (g) + I2 (g) → 2HI (g)

[HI]02

[H2]0 [I2]0Kc =

[1.98]2

[0.243] [0.146]Kc = Kc = 111

Reaction Quotient

The reaction quotient (Qc) is calculated by substituting the

initial concentrations of the reactants and products into

the equilibrium constant (Kc) expression.

IF

• Qc > Kc system proceeds from right to left to

reach equilibrium

• Qc = Kc the system is at equilibrium

• Qc < Kc system proceeds from left to right to

reach equilibrium

Reaction Quotient

Calculating Equilibrium

Concentrations

• If we know the equilibrium constant for a

reaction and the initial concentrations, we can

calculate the reactant concentrations at

equilibrium.

• ICE method

Reactants Products

Initial (M):

Change (M):

Equilibrium (M):


Concentrations

• At 1280ºC the equilibrium constant (Kc) for the reaction

is 1.1 x 10-3. If the initial concentrations are [Br2] = 0.063

M and [Br] = 0.012 M, calculate the concentrations of

these species at equilibrium.

Br2 (g) 2Br (g)

Let x be the change in concentration of Br2

Br2 (g) 2Br (g)

Initial (M)

Change (M)

Equilibrium (M)

0.063

-x

0.063 - x

0.012

+2x

0.012 + 2x


Concentrations[Br]2

[Br2]Kc =

Kc = (0.012 + 2x)2

0.063 - x= 1.1 x 10-3

4x2 + 0.048x + 0.000144 = 0.0000693 – 0.0011x

4x2 + 0.0491x + 0.0000747 = 0

ax2 + bx + c = 0 -b ± b2 – 4ac√2a

x = x = -0.00178

x = -0.0105


Concentrations

Br2 (g) 2Br (g)

Initial (M)

Change (M)

Equilibrium (M)

0.063 0.012

-x +2x

0.063 - x 0.012 + 2x

At equilibrium, [Br] = 0.012 + 2x = -0.009 M

At equilibrium, [Br2] = 0.063 – x = 0.0648 M

or 0.00844 M


Concentrations

• Express the equilibrium concentrations of all species in terms of the initial concentrations and a single unknown x, which represents the change in concentration.

• Write the equilibrium constant expression in terms of the equilibrium concentrations. Knowing the value of the equilibrium constant, solve for x.

• Having solved for x, calculate the equilibrium concentrations of all species.

Factors that Affect Chemical

Equilibrium

• Chemical Equilibrium represents a balance between forward and reverse reactions.

• Changes in the following will alter the direction of a reaction:

– Concentration

– Pressure

– Volume

– Temperature

Le Châtlier’s Principle

• Le Châtlier’s Principle- if an external

stress is applied to a system at

equilibrium, the system adjusts in such a

way that the stress is partially offset as the

system reaches a new equilibrium

position.

• Stress???

Changes in Concentration

• Increase in concentration of reactants

causes the equilibrium to shift to the

________.

• Increase in concentration of products

causes the equilibrium to shift to the

________.


Change Shift in Equilibrium

Increase in [Products] left

Decrease in [Products] right

Increase in [Reactants] right

Decrease in [Reactants] left


FeSCN2+(aq) ↔ Fe3+(aq) +

SCN-(aq)

a.) Solution at equilibrium

b.) Increase in SCN-(aq)

c.) Increase in Fe3+(aq)

d.) Increase in FeSCN2+(aq)

Changes in Volume and Pressure

• Changes in pressure primarily only

concern gases.

• Concentration of gases are greatly

affected by pressure changes and volume

changes according to the ideal gas law.

PV = nRT

P = (n/V)RT

Changes in Pressure and Volume

Change Shift in Equilibrium

Increase in Pressure Side with fewest moles

Decrease in Pressure Side with most moles

Increase in Volume Side with most moles

Decrease in Volume Side with fewest moles

Changes in Pressure and Volume

Changes in Temperature

• Equilibrium position vs. Equilibrium constant

• A temperature increase favors an endothermic reaction and a temperature decrease favors and exothermic reaction.

Change Endo. Rx Exo. Rx

Increase T K decreases K increases

Decrease T K increases K decreases

Changes in Temperature

Consider: N2O4(g) ↔ 2NO2(g)

The forward reaction absorbs heat; endothermic

heat + N2O4(g) ↔ 2NO2(g)

So the reverse reaction releases heat; exothermic

2NO2(g) ↔ N2O4(g) + heat

Changes in temperature??

Effect of a catalyst

• How would the presence of a catalyst

affect the equilibrium position of a

reaction?

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Chemical Equilibrium - libvolume4.xyzlibvolume4.xyz/.../chemicalequilibriumpresentation2.pdf · Equilibrium Equilibrium is a state in which there are no observable changes as time