Equilibrium Notes

The Concept of Equilibrium

• Ex) elevator, football game, moving walkway

1. Reversible Reactions: Chemical reaction in which the products can regenerate the original reactants . A double arrow is used to express a reversible reaction.

Ex)

2NO 2(g) ↔ N 2 O 4(g)

reactants

yield

products ↔ or or forward reaction reverse reaction Forward Rxn:

2NO 2 → N 2 O 4

Reverse Rxn:

N 2 O 4 → 2NO 2

2. Some reactions can reverse on their own , some under certain conditions , ( temp, pressure or a catalyst ) and some will not reverse. Ex) Single Replacement Rxns:

3CuCl 2(aq) + 2Al (s) → 3Cu (s) + 2AlCl 3(aq)

3. Chemical Equilibrium:

State of a product’s and reactant’s concentrations remaining constant , not equal . This is because the rate of the forward and reverse reactions are equal . The symbol “[ ]” denotes concentration .

A. Reaction rates are affected by concentration . If the concentration of the reactant increases, the rate of the forward reaction will increase .

B. As the reaction proceeds, the concentration of the reactants decrease and the products increases . This will cause the forward reaction to decrease .

4. Reaching equilibrium does not mean the reaction has stopped . It only means the rate of the forward and reverse reactions are equal . This is not a static equilibrium, but a dynamic equilibrium that stays constant over time.

A ↔ B

Concentration becomes Constant

NOT

Equal

Rates

become Equal

Reaction Rates & Equilibrium:

A. Collision Theory: (Theoretical Model) 1. In order for reactions to occur between substances, the particles must

collide

. 2. According to the

Collision Theory

, a successful collision occurs when A. the collision is

energetic

enough, and B. the particles collide with the correct

orientation

.

3.

Effective

collisions

lead

to the formation of products;

ineffective

collisions do

not

lead to the formation of products.

Ineffective Collision – Insufficient Energy - No Products

Effective and Ineffective Collisions

Effective Collision – Sufficient Energy – Forms Products

Effective collisions have enough energy, and the correct orientation to form products. Ineffective collisions revert to the original reactants.

4. A successful collision results in bond

breaking

(endothermic) and bond

forming

(exothermic).

5. The minimum energy needed to produce an

effective energy

collision is called the

activation

for the reaction. It’s abbreviation is

E a

.

Activation Energy

6. A transitional structure results from a successful collision. The structure is present while old bonds are

breaking

and new bonds are

forming

. It is called an

activated complex

and is unstable and short-lived. It is neither

reactant

nor

product

. 7. A substance that increases the rate of a chemical reaction by providing a mechanism with a lower energy of activation is called a

catalyst

.

8. A system that has just one phase is called a

homogeneous

system, (g→g→g) while a system that has more than one phase is called a

heterogeneous

(s→l→g) system.

Endothermic

9. Energy Diagram: shows changes in

energy

during a reaction.

Exothermic

Energy Diagram for a Chemical Rxn.

Reaction Pathway Activated Complex E a forward reaction

(80 kJ)

Energy of Reactants Added Catalyst ∆H = -60 kJ (for.) +60 kJ (rev.) E a reverse reaction

(140 kJ)

Forward Reaction (Exo) Energy of Products Reverse Reaction (Endo)

Factors Affecting Reaction Rates

1. Chemical Kinetics is concerned with the

rate

which a reaction occurs.

at 2.

Reaction rate

is a

measure

of the rate or speed of a chemical reaction. Reaction rate is determined by measuring the change in

concentration

of reactants and products over a certain amount of

time

. These reaction rates are determined experimentally.

3. Rate-influencing factors are those factors that affect

rate

of reactions by altering the

frequency

,

orientation

, or

energy

at which particles collide.

4. According to the Collision Theory, the following 5 factors affect the rate of a reaction: 1.

Nature of Reactants:

A. Structure:

complexity

of bonds broken and formed & the

orientation

.

B. State: Homogeneous systems have reactants and products in the

same

state. Heterogeneous systems have reactants and products in

different

states. Homogeneous systems usually react

faster

and heterogeneous systems react

slower

.

2 liquids would usually react quickly Iron rusts in air very slowly

2.

Temperature

:

(Average kinetic energy!)

of thumb is that every

10°C increase

in temperature

doubles the reaction rate

.

Rule

3.

Concentration

:

increased

concentration

,

increases

# collisions.

Aqueous solutions can

change [conc] or

molarity

,

gases can change

pressure

. Solids and pure liquids like

water

cannot change concentration.

4.

Surface Area

:

Increased

surface area

,

increases

frequency of collisions. This is especially true for

heterogeneous

systems

(s→l→g)

.

5.

Catalyst:

Increases the rate of reaction without being

consumed

in the reaction. Catalysts speed up the reaction rate by

lowering

the

activation energy

needed for the reaction to occur.

Uncatalyzed reaction (slow) Catalyzed reaction (fast)

5.

Inhibitors:

Decrease

the rate of the reaction by taking the place of a

reactant

and

stopping

the reaction (opposite of a catalyst.)

Energy diagram involving a catalyst:

11. What is in the reaction vessel at time = 0?

H 2 + N 2 12. Write the forward reaction: 3H 2 + N 2 → 2NH 3 13. What kind of reaction is the forward reaction?

synthesis it slows down 14. Write the reverse reaction: 2NH 3 → 3H 2 + N 2 15. What kind of reaction is the reverse reaction?

decomposition it speeds up 16. Over time the concentration of which substance(s)

decreases

?

H 2 + N 2 17. Over time the concentration of which substance(s)

increases

?

NH 3 18. Mark on the graph with a dashed line when equilibrium is reached.

19. At equilibrium which substance(s) is(are) present in the greater concentration?

H 2 + N 2 20. Is the

forward

or

reverse

reaction favored?

reverse

(min)

Equilibrium Systems and Stress – Le Chatlier’s Principle

1. Le Chatlier’s Principle states that if a stress is imposed on a system at equilibrium, the equilibrium position will

shift

toward the direction that tends to minimize the stress. (Pure

liquids

and

solids

are not affected by changes in equilibrium.) 2. In chemistry, Le Chatlier’s Principle is used to manipulate the outcome of reversible reactions to

maximize

the amount of

product

produced by altering the temperature, pressure, or removing a reactant or product from a reaction.

Crash Course in Chemistry – Equilibrium Video Khan Academy – LeChatlier’s Principle

The Haber Process

3. An application of Le Chatlier’s Principle is

The Haber Process

. This was used by

Germany

in WWI to produce

ammonia

(NH 3 ) for nitrogen containing explosives .

Today this process is used to create ammonia (

NH 3

) for household cleaners and especially

fertilizers

allowing for a four-fold increase in food production from 1900 to 2000.

4. There are

4

factors which determine whether a reaction favors making reactants

“lies to the left”

or favors making products

“lies to the right”

.

A. Changes in Concentration -Adding/Removing Reactants -Adding/Removing Products B. Changes in Pressure C. Changes in Temperature

Changes in Concentration - (Reactants and Products)

1. A dding

a substance to a system at equilibrium drives the system

AWAY

from that substance and makes more of the substance(s) on the opposite side.

2. R emoving

a substance from a system at equilibrium drives the system toward

REPLACING

that substance.

Ex)

N 2 O 4(g) ↔ 2NO 2(g)

↑ N 2 O 4(g) ↑ NO 2(g) ↓ N 2 O 4(g) ↓ NO 2(g) shift

right

shift shift shift

left left right

making more

NO 2

making more

N 2 O 4

making more

N 2 O 4

making more

NO 2

B. Changes in Pressure

1. Avogadro’s Law

– Equal volumes of gases at the same temperature and pressure have the

same

number of molecules. (

1 mole of any gas at STP =

22.4L

= 6.02 x10 23 particles

) 2. According to

Boyle’s Law

an increase in pressure means a

decrease

in volume. So if the pressure is

increased

on a system at equilibrium, the side which occupies the

lower

volume will be favored. (inverse relationship) 3. If pressure is increased, the reaction will shift in the direction that produces

fewer

moles.

Ex)

2NO 2(g) ↔ N 2 O 4(g)

↑ pressure shift to

right

making more

N 2 O 4 2H 2 O 2(g) ↔ 2H 2 O (g) + O 2(g)

↑ pressure shift to

left

making more

H 2 O 2 H 2(g) + Cl 2(g) ↔ 2HCl (g)

↑ pressure shift to

neither!

making more

same # of moles each side

4.Pressure changes will only affect

gases!

.

C. Changes in Temperature

1. Exothermic Reactions release

heat. If temperature is increased on this system then the reaction which absorbs or uses heat will

increase

. If temperature is decreased on an exothermic reaction, then the reaction which releases heat will

decrease

.

Ex)

H 2(g) + I 2(g) ↔ 2HI (g) + Heat

Increasing the temperature will produce a higher yield of

H 2 + I 2

.

Lowering the temperature will produce a higher yield of

HI

.

C. Changes in Temperature

2. Endothermic Reactions absorb

heat. If temperature is increased on this system then the reaction which absorbs or uses heat will

increase

. If temperature is decreased on an endothermic reaction, then the reaction which releases heat will

decrease

.

Ex)

HEAT + NH 4 Cl (s) ↔ NH 3(g) + HCl (g)

Increasing the temperature will produce a higher yield of

NH 3 + HCl

.

Lowering the temperature will produce a higher yield of

NH 4 Cl

.

Conclusion: These examples have been

illustrations of

Le Chatlier’s Principle

which states that a system at

equilibrium

, when subjected to a

stress

, will temporarily adjust itself to relieve the stress. This means that the

shift

to the right or left, or the

increased

forward or reverse reaction, will be temporary and a new equilibrium will be reestablished.

Practice: 1. 2SO 2(g) + O 2(g)



2SO 3(g) + heat

What conditions of temperature and pressure favor high equilibrium concentrations of SO 3 ?

(

high / low

) pressure; (

high / low

) temperature

1.

3H 2(g) + N 2(g)



2NH 3(g) + hea

t The commercial production of ammonia uses the Haber Process which is expressed by the above equation.

What condition of temperature and pressure will provide a maximum yield of NH 3 ?

(

high / low

) pressure: (

high / low

) temperature

1.

4HCl (g) + O 2(g)



2H 2 O (g) + 2Cl 2(g) + heat

Increasing the temperature

of the reaction will (

increase / decrease

) the forward reaction.

Decreasing the pressure

on the system will (

increase / decrease

) the forward reaction.