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Chapter 18 Chemical Equilibrium. Collision Theory of Reactions. Collision Theory of Gas-phase Reactions A chemical reaction can occur only when two molecules collide with a kinetic energy at least equal to certain energy E a , called activation energy of the reaction.

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collision theory of reactions
Collision Theory of Reactions

Collision Theory of Gas-phase Reactions

A chemical reaction can occur only when two molecules collide with a kinetic energy at least equal to certain energy Ea , called activation energy of the reaction.

The success of a collision also depends the relative orientation of molecules. This direction-dependence is called the steric requirement of the reaction.

collision theory of reactions1
Collision Theory of Reactions
  • Sufficient energy

proper orientation

  • Proper orientation

not sufficient energy

  • Sufficient energy

poor orientation

Chemical reaction is the overall effect of collisions

between reacting molecules

collision theory of reactions2
Collision Theory of Reactions

A conversion of kinetic energy to potential energy

occurs during formation of an intermediate complex that can either go on to form products or fall apart into the unchanged reactants.

This can be shown by a graph that traces the energy of the system before, during, and after the collision.

energy changes during a reaction1
Energy Changes During a Reaction

Transition State Complex

In the transition state complex, the original bonds have weakened, whereas the new bonds are only partially formed .

Activation Energy Ea:

The difference between the

energy of the transition state

complex and the reactant

energy.

rate of a chemical reaction
Rate of a Chemical Reaction

Three important factors influence

the speed of chemical reactions:

Temperature

The higher the temperature, the faster the rate of reaction.

Catalysis

A catalyst increases the rate of reaction.

Concentration of Reactants

The greater the concentration, the greater the rate of reaction.

effect of temperature on reaction rates
Effect of Temperature on Reaction Rates

Kinetic energy distribution curves at two temperatures explain

the effect of temperature on reaction rates.

Ea, the activation energy,

is the same at both temperatures.

Only the fraction of the particles

in the sample represented by the

area beneath the curve to the

right of Ea is able to react.

The fraction of molecules that is

able to react increases rapidly as the temperature is raised.

effect of catalyst on the rate of a chemical reaction change in activation energy
Effect of Catalyst on the Rate of a Chemical Reaction: Change in Activation Energy

A catalyst speeds up a reaction by providing a new pathway that has a lower activation energy.

effect of change in activation energy on the rate of a chemical reaction1
Effect of Change in Activation Energy on the Rate of a Chemical Reaction

The fraction of molecules that collide with kinetic energy that is at least equal to the activation energy, Ea , is bigger in a catalyzed reaction because the activation energy barrier is lowered.

Since Ea‘< Ea

The catalyzed

reaction rate

is faster.

effect of concentration on reaction rate
Effect of Concentration on Reaction Rate

Reaction rate depends on the frequency of effective collisions:

The more particles there are in a given volume,

the more frequently collisions will occur and

the more rapidly the reaction will take place.

development of equilibrium
Development of Equilibrium

For a reversible reaction in a closed system,

the equilibrium is established when

the forward reaction rate is

equal to the reverse reaction rate.

development of equilibrium1
Development of Equilibrium

If the system is not in equilibrium, the concentration of the species in the faster reaction will decrease, and thus the reaction will become slower; the concentration of the species in the slower reaction will increase, and thus the reaction will become faster.

Opposite rates

will eventually

become equal,

and an equilibrium

will be established.

the equilibrium constant
The Equilibrium Constant

Consider the reaction H2(g) + I2(g) 2 HI(g)

At equilibrium the following ratio is a constant

K is called equilibrium constant

the equilibrium constant1
The Equilibrium Constant

For the general equilibrium

a A + b B c C + d D

When writing an equilibrium constant expression,

use only the concentrations of gases, (g),

or dissolved substances, (aq).

Do not include solids, (s), or liquids, (l).

the equilibrium constant2
The Equilibrium Constant

Equilibrium Constant, K

For any equilibrium at a given temperature, the ratio of the product of the concentrations of the species on the right side of the equilibrium equation, each raised to a power equal to its coefficient in the equation, to the corresponding product of the concentrations on the left side of the equation, each raised to a power equal to its coefficient in the equation, is a constant.

The equilibrium constant is both

equation-dependent and temperature-dependent.

significance of the value of k
Significance of the Value of K

Example:

Is the forward reaction favored, the reverse reaction favored, or are appreciable quantities of all species present at equilibrium in the following reaction?

HC2H3O2(aq) H+(aq) + C2H3O2–(aq)

K = 1.8 × 10–5.

Solution:

Since K is very small, the reverse reaction is favored.

significance of the value of k1
Significance of the Value of K

Consider the general reaction: Reactants Products

If the equilibrium constant is very large (K > 100),

[Products] > [Reactants], so the forward reaction is favored.

If the equilibrium constant is very small (K < 0.01),

[Products] < [Reactants], so the reverse reaction is favored.

If the equilibrium constant is neither larger nor small,

[Products] ≈ [Reactants], so appreciable quantities of all species are present at equilibrium.

le chatelier s principle
Le Chatelier’s Principle

Le Chatelier’s Principle

If a system is in equilibrium, any change imposed on the system tends to shift the equilibrium in a direction that tends to counteract the initial change.

Le Chatelier’s principle only suggests an outcome; it does not provide an explanation.

le chatelier s principle the pressure volume effect
Le Chatelier’s PrincipleThe Pressure (Volume) Effect

A gas-phase equilibrium responds to compression-a reduction in volume of the reaction vessel.

If a gaseous equilibrium is compressed, the equilibrium will be shifted in the direction of formation of fewer molecules, thus minimizes the increase in pressure.

If the system is expanded, the shift will be in the direction of formation of more molecules.

3 H2 (g) + N2 (g) ↔ 2 NH3 (g)

More molecules less molecules

le chatelier s principle the pressure volume effect1
Le Chatelier’s PrincipleThe Pressure (Volume) Effect

3 H2 (g) + N2 (g) ↔ 2 NH3 (g)

If the pressure increases, the equilibrium will be shifted to the right (less molecules).

If the pressure decreases, the shift will be to the left.

To increase the yield of ammonia, industrial process uses pressures of 250 atm or higher.

le chatelier s principle the temperature effect
Le Chatelier’s PrincipleThe Temperature Effect

If a reaction is exothermic, the reverse reaction is endothermic.

If the temperature increases, the equilibrium will be shifted to the direction of consuming heat (endothermic, to the left for NH3 reaction lelow).

If the temperature decreases, the shift will be in direction of producing heat (exothermic, to the right).

3 H2 (g) + N2 (g) ↔ 2 NH3 (g) + 92kJ

le chatelier s principle the temperature effect1
Le Chatelier’s PrincipleThe Temperature Effect

The left tube at 25 0C contains very little brown gas compared to the tube on the right at 80 0C

N2O4 (g)+ heat ↔ 2 NO2 (g)

colorless brown

le chatelier s principle the concentration effect
Le Chatelier’s PrincipleThe Concentration Effect

Let us consider the reaction equilibrium:

3H2 (g) + N2 (g) ↔ 2 NH3 (g)

If H2 is added to the reaction chamber, the shift will be in the forward direction to counteract the increase in the number of hydrogen molecules thus producing more NH3.

If H2 is removed, the equilibrium will shift to the reverse direction to increase the H2 concentration.

le chatelier s principle the concentration effect1
Le Chatelier’s PrincipleThe Concentration Effect

Let us consider the reaction equilibrium:

3H2 (g) + N2 (g) ↔ 2 NH3 (g)

If H2 is added (increase of [H2]), the shift will be in the forward direction to increase in the concentration [NH3] and decrease the concentrations [H2] and [N2].

Adding an inert gas has no effect on the equilibrium, although the total pressure increases.

solubility equilibria
Solubility Equilibria

The equation for dissolving AgCl , a low-solubility compound, is

AgCl (s) ↔ Ag+ ( aq) + Cl- (aq)

This equilibrium is characterized by the solubility product constant Ksp

Ksp = [Ag+ ] [Cl-]

solubility equilibria1
Solubility Equilibria

Calculation of solubility product from the solubility: The chloride ion concentration of a saturated solution of silver chloride is 1.3 x 10-5 M . Calculate the solubility product for silver chloride .

Ksp = [Ag+ ] [Cl-]

In saturated solution of pure silver chloride the concentration of [Ag+ ] and [Cl-] are equal. Therefore

Ksp = [Ag+ ] [Cl-] =(1.3 x 10-5) x(1.3 x 10-5)

= 1.7 x 10-10

solubility equilibria2
Solubility Equilibria

Solubility and Solubility Product.

For compounds of similar structure, the smaller the solubility product, the smaller the solubility. For example the solubility of silver bromide ( Ksp = 5.2 x 10-13 ) is lower than the solubility of silver chloride.

Common Ion Effect

  Suppose that a soluble chloride, such as NaCl were to be added to the saturated solution of silver chloride. According to the Le Chatelier’s principle the equilibrium would shift the equilibrium in the reverse direction, reducing the solubility of silver chloride.

.

ionization equilibria weak acid
Ionization Equilibria: Weak acid

Weak acids ionize only slightly when dissolved in water.

For a general weak acid HA,

HA(aq) H+(aq) + A–(aq)

Major species: HA(aq)

Minor Species: H+(aq) + A–(aq)

ionization equilibria weak acid1
Ionization Equilibria: Weak acid

The ionization of a weak acid is usually so small that it is negligible compared with the initial concentration of the acid.

We assume that all ionization concentrations are negligible

when subtracted from the initial concentration.

In other words, the initial concentration of the weak acid is also the final concentration after the acid ionizes.

ionization equilibria weak acid2
Ionization Equilibria: Weak acid

Example:

Find the pH of 0.1 M nitrous acid. Ka = 4.5 × 10–4.

Solution:

HNO2(aq) H+(aq) + NO2–(aq)

Let x = [H+] = [NO2–]; [HNO2] = 0.1 M

ionization equilibria weak acid3
Ionization Equilibria: Weak acid

Find the pH of 0.1 M nitrous acid. Ka = 4.5 × 10–4.

Solution:

x2 = (0.1) (4.5 × 10–4)

x = [H+] = 7 × 10–3

pH = – log [H+] = – log (7 × 10–3) = 2.2

ionization equilibria buffer solution
Ionization Equilibria: Buffer Solution

Buffer Solution

A solution that resists changes in pH because

it contains relatively high concentrations of both

a weak acid and a weak base.

The acid reacts with any added OH–;

The base reacts with any added H+.

ionization equilibria buffer solution1
Ionization Equilibria: Buffer Solution

Determine the pH of a solution that is 0.25 M in HAc and 0.35 M in NaAc. Ka = 1.8 × 10–5.

NaAc(aq)  Na+(aq) + Ac–(aq)

HAc(aq) H+(aq) + Ac–(aq)

pH = – log [H+] = – log (1.3 × 10–5) = 4.89

ionization equilibria buffer solution2
Ionization Equilibria: Buffer Solution

Example:

Determine the acid-to-base concentration ratio that will yield a buffer solution with a pH of 4.50 if the acid has

Ka= 1.0 × 10–5.

Solution:

HA(aq) H+(aq) + A–(aq)

[H+] = antilog (–pH) = antilog (–4.50) = 3.2 × 10–5 M

homework
Homework
  • Homework: 25, 27, 29, 31, 33, 35, 37, 41, 43, 47, 61, 63, 71, 75