Skip to document
Welcome to StudocuSign in to access the best study resources
Guest user
0followers
HomeMy LibraryAsk AI
  • You don't have any recent items yet.
My Library
  • You don't have any modules yet.
  • You don't have any books yet.
  • You don't have any Studylists yet.

Chemical equilibrium and kinetics

notes on kinetics and chemical equilibrium
Module

General and Organic Chemistry (AOC106DI)

26 Documents
Students shared 26 documents in this course
Academic year: 2020/2021
Uploaded by:
Anonymous Student
This document has been uploaded by a student, just like you, who decided to remain anonymous.
Coventry University

Comments

Please sign in or register to post comments.

Students also viewed

Related documents

Related Studylists

as chemistry

Preview text

Catalysis:

A catalyst is a chemical that speeds up a reaction without being consumed in the process. The activation energy of most catalysts is reduced. With the activation energy reduced, more collisions will have enough relative kinetic energy to cause a reaction. As a result, more reactions will occur, and the overall reaction rate will increase. Because a catalyst boosts both forward and reverse processes, it does not change the equilibrium constant.

Heterogeneous catalysts are catalysts that are in a different phase than the reactants or products. The reactants and products are usually gases or liquids, while the catalysts are usually solids. In the same phase as the reactants and products is a homogeneous catalyst. These catalysts are frequently liquid or gaseous in nature. Homogeneous catalysts are usually aqueous acids and bases.

Equilibrium:

Many chemical reactions are reversible, and forward and reverse reactions can happen simultaneously. Chemical equilibrium occurs when the rate of the forward reaction is equal to the rate of the backward reaction.

A constant K characterises the state of equilibrium:

aA + bB cC + dD ←→

where a, b, c, and d are the corresponding stoichiometric coefficients.

K = [C]c [D]d / [A]a [B]b

The reaction quotient:

For reactions that are not in equilibrium, the following equation is used:

Q stands for reaction quotient, which is used to anticipate which way a reaction would go. Q will constantly fluctuate toward K because reactions always progress toward equilibrium. For a reaction, Q and K can be compared to determine which direction the reaction will go.

  1. The reaction is at equilibrium if Q equals K.

  2. When Q exceeds K, the ratio of products to reactants is larger than when the system is at equilibrium. This indicates that the reverse reaction rate will be higher than the forward rate.

  3. If Q is smaller than K, the product-to-reactant ratio is lower than it is at equilibrium. This demonstrates that the forward response rate will outnumber the reverse reaction rate.

The principle of Le Chatelier can be applied to systems in equilibrium. When a system in equilibrium is strained, the principle states that the system will change in a way that reduces stress.

Le Chatelier's approach applies to the following sorts of stresses:

  • Heating or cooling the system

  • Adding or removing a product or reactant

  • Changing the system's pressure

The Haber process is the reaction that follows. This method is employed in the production of ammonia from nitrogen and hydrogen. With temperatures between 400 and 450 degrees Celsius and a pressure of 200 atm, an iron catalyst is utilised.

N 2 (g) + 3H 2 (g)  2NH 3 (g) + Heat

What can happen to the Haber process as a result of various stresses:

The reaction is exothermic, which means it generates heat (heat on right hand side of reaction). The reaction will shift to the left (reverse reaction) and become endothermic as the temperature rises. The additional heat on the right side of the reaction will be reduced as a result of this. There will be a drop in NH3 and an increase in both N2 and H2 as a result of this shift.

N2, H2, and NH3 gas are held in equilibrium in a container. If N2 gas is supplied, the system will use the forward reaction to reduce the partial pressure of N2 to compensate for the higher concentration of N2. The partial pressures of N2 and H2 will both fall as a result of this forward reaction. As a result, NH3 and heat levels will rise (forward reaction).

Total pressure rises when the temperature of a container remains constant while the container's size is reduced. On the left side of the process, there are four gas molecules and two on the right. To bring the pressure closer to equilibrium, the equilibrium will move to the side with the fewest gas molecules. As the equilibrium shifts to the right side, there will be an increase in NH3 and heat will be generated.

Was this document helpful?

Chemical equilibrium and kinetics

Module: General and Organic Chemistry (AOC106DI)

26 Documents
Students shared 26 documents in this course

University: Coventry University

Was this document helpful?
Catalysis:
A catalyst is a chemical that speeds up a reaction without being consumed in the process.
The activation energy of most catalysts is reduced. With the activation energy reduced,
more collisions will have enough relative kinetic energy to cause a reaction. As a result,
more reactions will occur, and the overall reaction rate will increase. Because a catalyst
boosts both forward and reverse processes, it does not change the equilibrium constant.
Heterogeneous catalysts are catalysts that are in a different phase than the reactants or
products. The reactants and products are usually gases or liquids, while the catalysts are
usually solids. In the same phase as the reactants and products is a homogeneous catalyst.
These catalysts are frequently liquid or gaseous in nature. Homogeneous catalysts are
usually aqueous acids and bases.
Equilibrium:
Many chemical reactions are reversible, and forward and reverse reactions can happen
simultaneously. Chemical equilibrium occurs when the rate of the forward reaction is
equal to the rate of the backward reaction.
A constant K characterises the state of equilibrium:
aA + bB cC + dD ←→
where a, b, c, and d are the corresponding stoichiometric coefficients.
K = [C]c [D]d / [A]a [B]b
The reaction quotient:
For reactions that are not in equilibrium, the following equation is used:
Q stands for reaction quotient, which is used to anticipate which way a reaction would go.
Q will constantly fluctuate toward K because reactions always progress toward
equilibrium. For a reaction, Q and K can be compared to determine which direction the
reaction will go.
1. The reaction is at equilibrium if Q equals K.
2. When Q exceeds K, the ratio of products to reactants is larger than when the system is at
equilibrium. This indicates that the reverse reaction rate will be higher than the forward
rate.

Students also viewed

Related documents