Chemistry 3202 – Test 2 Review:

Topics Covered:

Collision Theory: For a chemical reaction to occur, reacting particles must collide. Not all collisions are successful however, they must meet two main conditions:

i. The collision must have sufficient force. As discussed in the previous unit, it takes energy to break a chemical bond. For new chemicals to form, old bonds must be broken. When reactants collide, the kinetic energy the particles possess is turned into potential energy (straining the chemical bonds). If the strain is sufficient, the reacting molecule will break its original bonds and may possibly form new ones.

ii. The collision must be of the correct orientation. When particles collide, they must collide with the correct positioning so that as old bonds are broken, the reactants are positioned in such a way that the new bonds of the products can form. Without correct positioning, a reaction cannot happen even if the collision has sufficient force.

Kinetic Molecular Theory: Gas particles are in constant motion in random directions. They move in straight lines and rarely collide. Gas particles spread out as far apart as possible, this is known as diffusion. Evidence for KMT includes:

i. Diffusion. Gas particles spread out as far apart as possible, which is why we can smell perfume across a room.

ii. Pressure. The force that is known as pressure is caused by the continual collision of gas particles against an object. i.e. a balloon stays inflated because of the gas particles constantly hitting up against the walls of the balloon.

Measuring Rates of Reaction:

It is important to be able to measure some observable property when determining reaction rates. The method used depends on the nature of the reaction. Common indicators of changes of rate include, but are not limited to: pH, Temperature, Electrical Conductivity, Colour, volume, mass, pressure. Generally, the rate of a reaction is determined as such:

Rate = Change in property / Change in time

Factors Affecting Reaction Rate:

There are 5 main factors that affect a reaction rate:

i. Temperature. Temperature affects a reaction rate because temperature is the ‘average kinetic energy’ of a sample of matter. As the overall average kinetic energy increases, the proportion of molecules that have enough energy to successfully collide increases as well. The amount of energy needed to sufficiently collide is known as the threshold energy.

ii. Concentration of reactants. Concentration affects reaction rate because with increased concentration the number of collisions between reactants increases as well. As the number of collisions increases, the likelihood of having a successful collision increases too. ( In the gaseous phase, the pressure of a reactant acts similarly.) Example: Under certain conditions, 1% of collisions in a reaction are successful. With an increase in concentration, the number of collisions that are successful is still 1% however the overall number of collisions increases. A dilute solution may have 100 collisions per minute 1 successful collision , while a concentrated solution may have 500 collisions per minute 5 successful collisions.

iii. Surface area of a solid reactant. Increasing the surface area of a solid reactant allows for more of the reactant to be available to be collided with. Therefore, a solid that has a higher surface area (eg. A powder) will react faster than a solid with a lower surface area.

iv. Addition of a catalyst. A catalyst alters the mechanism of a reaction, so that the amount of energy needed for a successful collision is decreased. Catalysts cause a reaction that normally needs one step to take several steps, with each elementary process requiring less energy overall that the one step process. Examples: Enzymatic processes in the human body are controlled by catalysts known as enzymes.

v. Nature of the reactants. There are a number of ways in which a reaction can be affected by the types of reactants that are used. a. State of Matter: Gases react faster than liquids, which react faster than solids.

Aqueous ionic compounds will react almost instantlyb. Bond type: Ionic bonds are weaker than molecular ones, so ionic compounds

generally react faster than molecular compoundsc. Bond strength: In molecular compounds, a molecule with a strong bond (i.e. a triple

bond) will react slower than a weaker bond (a single bond)d. Acid Strength: Strong acids react quicker than weak acids, because they are

completely ionized. This means they produce more H+ ions than weak acids, which are what react during a reaction with an acid. There are 6 strong acids in total (HCl, HBr, HI, H2SO4, HNO3, HClO4)

Potential Energy Diagrams:

A potential energy diagram shows the amount of potential energy during the course of a reaction. A potential energy diagram has similarities to an enthalpy diagram in that the reactants and products are shown at the start and end of the reaction. The difference along the Y axis shows the energy change, and whether the reaction was exothermic or endothermic. Between the reactants and products is a region where there is an increase in potential energy. This occurs when the reactants collide. During the collision, the kinetic energy of the reactants becomes converted to potential energy. If the collision has sufficient force, the reaction will reach the transition state (aka activated complex). This is a high energy state whereby the reactants start becoming products (hence the name transition state). At this point, 1 of 2 things will occur

1. The reactants will reform, and the collision will be considered unsuccessful.2. The products will form, and the collision will be successful.

The amount of energy required for the forward reaction is considered the Forward Activation Energy (Eafwd ) The amount of energy required for the reverse reaction is the Reverse Activation Energy (Earev) Both the forward and reverse activation energies are always positive.

To calculate the forward, reverse or overall energy:

Eafwd - Earev = ∆E

Some reactions will have multiple steps, each having their own activation energy. Notably, catalyzed reactions will have multiple steps, and the activation energy for a catalyzed reaction will be less than that of an uncatalyzed one.

Reaction Mechanisms:

Many reactions do not occur in a single step; rather they occur in multiple steps. Each step is known as an elementary process. Each elementary process has initial reactants and includes a single successful collision. The elementary process may produce true products, which do not react further, or alternatively a reaction intermediate may be formed. A reaction intermediate will then be used up in a subsequent elementary process. As well, when a catalyst is used in a reaction, it will be used in one elementary process, but will be produced once again in a later elementary process. In both cases the reaction intermediate and the catalyst are not included in the overall reaction, as their net production/consumption is zero.

One important feature of a reaction mechanism is that in many cases the rates of each elementary process are stated. The slowest elementary process is known as the rate determining step. The rate determining step can also be identified on a potential energy diagram because it has the greatest activation energy.

Example:

Step 1 A + B C fast

Step 2 C + B D + E slow

Step 3 D + B E + A very fast

In the above reaction mechanism, A is a catalyst, D and C are reaction intermediates. They are not produced nor consumed in the reaction. A is first consumed, then produced later on, while C & D are produced first then consumed.

The overall reaction is :

3B 2E

Drawing a Diagram: In this reaction, it is not stated whether the process is exothermic or endothermic, so the overall energy change does not necessarily have to be drawn. However, in certain circumstances values for a multi-step process will be given (Ea for forward and reverse processes, ∆E) and the diagram will have to reflect those values.

Equilibrium: A Dynamic Process

Equilibrium is a phenomenon whereby a reaction can be reversed, so that the forward and reverse of a reaction happen at the same time at the same rate. Two conditions must be met before equilibrium can be established.

1. The process must occur in a closed system. If either the products or reactants are allowed to be removed in a process, equilibrium cannot be established. Either the forward or reverse reaction will always exceed the other, with one of the processes going to completion.

2. The process is so highly exothermic that the activation energy of the reverse reaction cannot be reached.

It is important to note that equilibrium is a dynamic process; while no net change is occurring, both processing are continuing to happen.

Establishing Equilibrium:

Equilibrium must be established over time, if a reaction starts with only the reactants present. Also, when equilibrium is established there may not be equal amounts of reactant and product present, only the rates of the forward and reverse reaction are equal.

Consider the following example:

A + B C + D

1. (Starting with only A and B present)

The forward reaction is at its fastest, since this is when the concentrations of A and B are at their greatest. The reverse reaction does not occur, since the concentrations of C and D are zero.

2A + B C + 3D ‘Fast’

C + 3D 2A + B ‘Zero’

2. After time passes

The forward reaction slows down continuously, since the concentrations of A and B are decreasing. The rate of the reverse reaction is increasing, as C and D are being produced. **It is important to note however, that the forward reaction is still faster than the reverse. The reverse reaction will never become faster than the forward reaction while equilibrium is being established initially.

2A + B C + 3D ‘Slowing Down’

C + 3D 2A + B ‘Speeding Up’

3. Equilibrium is established

The forward reaction has slowed down, while the reverse reaction has speeded up to the point where the reactants are being produced at the same rate as the products. No net change in the concentration of any of the 4 compounds will occur, unless the equilibrium is somehow disturbed.

2A + B C + 3D ‘Equal rates’

C + 3D 2A + B ‘Equal rates’

Disturbing Equilibrium:

Once an equilibrium is established, it will remain at equilibrium until it is disturbed in some manner. Equilibrium will not be affected by changes to surface area nor the presence of a catalyst, since both of these will influence the rate of a reaction equally. Equilibrium will be disturbed if there is a change to the concentration of one or more species present, the pressure (if one or more reactants is a gas) or the temperature.

1. Changing the concentration

When the concentration of a species present is increased, it results in speeding up the reaction of which it is a reactant. This will cause the equilibrium to be disturbed, since one reaction will be faster than the other. This will cause a net increase in products, with a net decrease in reactants (from the point at which equilibrium was disturbed. The reactant that was increased will have an overall increase in concentration in comparison to when equilibrium was first established)

Example:

2A + B C + 3D

Increasing the Concentration of B will increase the forward rate of reaction. A net increase of C and D will result. The concentration of A will decrease. The concentration of B will be between its concentration before equilibrium was disturbed and its concentration immediately after equilibrium was disturbed.

Thought question: How would the graph look if the concentration of a product was decreased?

Effect of a pressure change (gaseous equilibrium):

In a gaseous equilibrium a pressure increase will cause the reaction to increase the concentration of the side where there is the least gas produced. Hence, the reaction (forward or reverse) that will be favored in high pressure conditions produces the least number of moles of gas, while in low pressure conditions the side producing the greatest number of moles of gas will be favored.

Example:

4 moles of gas 8 moles of gas

In high pressure conditions, equilibrium will favor the reactants, while in low pressure conditions equilibrium will favor the products.

Effect of Temperature Change:

An increase in temperature change will result in the forward and reverse reaction speeding up. However, the endothermic process will speed up to a greater extent. Therefore, an increase in temperature will cause a net increase in the endothermic products and a net decrease in the exothermic products. It may be useful to think of the equilibrium in terms of a thermochemical equation, and increasing the temperature is analogous to increasing the ‘concentration’ of heat.

Increasing the temperature:

Increasing the temperature speeds up the endothermic process more than it speeds up the exothermic process. The result is that there is a net gain of the products of the endothermic process, and a decrease in the products of the exothermic process.

In this example, CO and O2 are produced in greater quantities, since they are the product of the endothermic process, which has speeded up to a greater extent than the exothermic process.

Decreasing the Temperature:

Decreasing the temperature slows down the exothermic and endothermic reaction, however once again temperature change affects the endothermic process more. As a result, the exothermic process is faster than the endothermic process, and there is a net gain of the products of the exothermic reaction.

In this example, a decrease in temperature will cause more of the exothermic products to be produced (CO2) and a net decrease in the amount of the endothermic products (CO and O2)

Le Chatalier’s Principle:

“A Dynamic Equilibrium tends to respond so as to relieve the effect of an change in the conditions that affect the equilibrium.”

The idea behind Le Chatalier’s Principle is that when a change is made to an equilibrium, the system will respond to counter that change. For example, if the concentration of a compound is increased, the system will respond to decrease the concentration of that compound. If the concentration is decreased, the system will respond to increase that concentration. Temperature and pressure work the same way.

Temperature increases favor endothermic processes (which ‘consume’ heat) and temperature increases favor exothermic processes (which release heat)

Pressure increases favor the reaction side with the fewest moles of gas (to reduce pressure) while pressure increases favor the side of the reaction with the most moles of gas (to increase pressure).

IMPORTANT NOTE: Adding a catalyst and increasing surface area have no impact on a system at equilibrium. While they will allow a system to reach equilibrium faster, the end equilibrium that is reached will be the same either with or without the catalyst / altering surface area. This is because both the forward and reverse reaction are impacted equally.

Practice Problems: Le Chatalier’s Principle

P 529 # 33-35, P. 533 #1-5

The following are found online at : http://www.saskschools.ca/curr_content/chem30_05/3_equilibrium/assignments/assignment3.htm

Le Châtelier's Principle1.

State Le Châtelier's Principle.

2.

What are three stresses that can affect the position of an equilibrium? Identify the one stress that will cause the value of Keq to change.

3.

State the effect of a catalyst on equilibrium.

4.

Methanol (methyl alcohol; CH3OH) can be manufactured using the following equilibrium reaction:

CO(g) + 2 H2 (g) ↔ CH3OH(g) + energy

Predict the effect of the following changes on the equilibrium concentration of CH3OH(g). Will it’s concentration increase, decrease, or remain the same?

  a. The temperature of the system is decreased.

  b. The pressure of the system is increased.

  c. More H2 (g) is added.

  d. A catalyst is added to the system.

5.

Use Le Châtelier's Principle to predict how the changes listed will affect the following equilibrium reaction:

2 HI(g) + 9.4 kJ ↔ H2 (g) + I2 (g)

  a. Will the concentration of HI increase, decrease, or remain the same if more H2 is added?

  b. What is the effect on the concentration of HI if the pressure of the system is increased?

  c. What is the effect on the concentration of HI if the temperature of the system is increased?

  d. What is the effect on the concentration of HI if a catalyst is added to the system?

6.

Suggest four ways to increase the concentration of SO3 in the following equilibrium reaction:

2 SO2 (g) + O2 (g) ↔ 2 SO3 (g) + 192.3 kJ

7.

In the equilibrium reaction:

4 HCl(g) + O2 (g) ↔ 2 H2O(g) + 2 Cl2 (g) + 114.4 kJ

Predict the direction of equilibrium shift (forward, reverse, no change) if the following changes occur:

  a. The pressure is increased.

  b. Heat is added.

  c. Oxygen is added.

  d. HCl is removed.

  e. A catalyst is added.

8.

Nitric oxide, NO, releases 57.3 kJ/mol when it reacts with oxygen to give nitrogen dioxide.

  a. Write a balanced equation for this reaction.

  b. Omit….

c. Predict the effect that increasing the temperature will have on:

the equilibrium concentration of all reaction participants (NO, O2, and NO2)

Practice on Reaction Rates:

Chemistry 30

Unit 2: Chemical KineticsAssignment 3: 4-1 to 4-3

1. Explain why the rate of a chemical reaction increases as the concentration of the reactants increases.

2. Consider two gases A and B in a container at room temperature. What effect will the following changes have on the reaction rate between these gases (increase, decrease, no effect)?

a) The pressure is increased.

b) The number of molecules of gas A is doubled

c) The temperature is decreased

3. Which of the following reactions will have the fastest rate? The slowest? Explain

a) C12H26 (s) + O2 (g) ® 12 CO2 (g) + 13 H2O (g)

b) S2O82-(aq) + 2 I-(aq) ® 2 SO4

2-(aq) + I2(s)

c) Ba2+(aq) + SO42-(aq) ® BaSO4(s)

Explanation:

4. The series of steps by which an overall chemical reaction takes place is called the

5. The slowest step in the series of steps in a chemical reaction is called the

6. Define activation energy.

7. Consider the following potential energy diagrams which represent two chemical reactions. On the basis of these diagrams, which reaction would you expect to occur at a faster rate? Why?

A B

8. Which will react faster, zinc with 3 M hydrochloric acid or zinc with 1 M hydrochloric acid? Why?

9. White phosphorus reacts immediately and rapidly with oxygen when exposed to air. What can you say about the amount of activation energy required for this reaction?

10. Hydrogen peroxide reacts with hydrogen ions and iodide ions according to the following reaction mechanism:

step 1. H+ + H2O2 ® H3O2+ fast

step 2. H3O2+ + I- ® H2O + HOI slow

a) Write the overall reaction described by this mechanism

step 1. H+ + H2O2 ® H3O2+

step 2. H3O2+ + I- ® H2O + HOI

overall:

b) If you wanted to increase the rate of the overall reaction, would it be better to increase the concentration of H+ or I-. Why?

12. Consider the following potential energy diagram:

a) Is the forward reaction endothermic or exothermic: _____________________

b) Determine H for the forward reaction: _____________________

c) Determine H for the reverse reaction: _____________________

d) Determine Ea for the forward reaction: _____________________

e) Determine Ea for the reverse reaction: _____________________

f) Label the location of the activation complex in the diagram.

g) Add to the diagram a possible pathway for a catalyzed reaction.

13. Sketch a potential energy curve for an reaction based on the information provided below.

Label the parts representing the activated complex, activation energy, and change in enthalpy, H.

Hforward = -30 kJ Ea reverse = +50 kJ