Unit 9: Kinetics, Thermodynamics, & Equilibrium-Lecture ... U… · Unit 9: Kinetics,...

Unit 9: Kinetics, Thermodynamics, & Equilibrium-Lecture Regents Chemistry ’14-‘15 Mr. Murdoch

Unit 9a (Kinetics & Energy Changes) of 43

Unit 9a:

Kinetics and Energy Changes

1.

Student Name: _______________________________________

Class Period: ________



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Unit 9 Vocabulary:

1. Activated Complex: The species that are formed and decomposed

during the mechanism; also called the intermediate.

2. Activation Energy: The energy that must be added to allow the

reactants to start the reaction and form the activated complex.

3. Catalyst: A chemical that is added to a reaction to eliminate steps in

the mechanism, increase the reaction rate, and decrease the activation

energy without itself being consumed by the reaction.

4. Effective Collision: A collision between reactant particles that results

in a chemical reaction taking place.

5. Enthalpy: The total amount of potential energy stored in a substance.

6. Endothermic: A reaction that absorbs and stores energy from the

surrounding environment.

7. Entropy: A system’s state of disorder. Entropy increases as

temperature increases. Entropy increases as a substance goes from

solid to liquid to gas.

8. Equilibrium: A system where the rate of forward change is equal to

the rate of reverse change. At equilibrium there is no net change.

9. Exothermic: A reaction that releases stored energy into the

surrounding environment.

10. Favored: A change in a thermodynamic property that contributes

towards the reaction being spontaneous.

11. Gibbs Free Energy: The total amount of energy available in a

system to do work. Free Energy is a combination of both enthalpy

and entropy.

12. Heat of Reaction: The net gain or loss of potential energy during a

chemical reaction.

13. Inhibitor: A chemical that is added to a reaction to add steps to the

mechanism, decrease the reaction rate, and increase the activation

energy without itself being consumed by the reaction.

14. Kinetics: The study of reaction mechanisms and reaction rates.

15. Nonspontaneous: A reaction that requires a constant input of energy

to occur, or the reaction will reverse or stop.



16. Reaction Rate: The amount of reactant consumed in a given unit of

time.

17. Spontaneous: A reaction that continues independently once started.

18. Thermodynamics: The study of heat flow during physical and

chemical changes.

19. Unfavored: A change in a thermodynamic property that contributes

towards the reaction being nonspontaneous.



Notes page:



Unit 9a Homework Assignments:

Assignment: Date: Due:



Chemical Kinetics:

Chemical Kinetics are the study of:

i. Reaction mechanisms (how reactions occur) and;

ii. Reaction rates (how long the reaction takes to complete).

Mechanism:

A mechanism is the pathway the reaction takes.

i. Each mechanism is a series of steps that leads from reactants to

products;

ii. Particles of reactant material must collide to react;

iii. Collisions must occur with enough activation energy to react;

iv. Molecules must be orientated (positioned) properly to react.

Collisions that satisfy these requirements and lead to the initiation of

the reaction are labeled effective collisions.

The effective mechanism can occur in a series of steps, each

involving electron shifts as old bonds are broken and new bonds are

formed.

How to speed up chemical reactions (and get a date) - Aaron Sams - YouTube - 4:56

Topic: Kinetics

Objective: How do reactions occur, and how fast do they occur?

https://www.youtube.com/watch?v=OttRV5ykP7A



Reaction Example: CO(g) + NO2(g) CO2(g) + NO(g)

This reaction does NOT go directly as stated above. It first starts with

the decomposition of nitrogen dioxide.

Step #1: 2 NO2(g) NO3(g) + NO(g)

Step #2: CO(g) + NO3(g) CO2(g) + NO2(g)

Final step: CO(g) + NO2(g) CO2(g) + NO(g)

Reaction Example (Reference Table I): H2 + I2 2 HI

For the reaction above, a possible reaction might be:

1. Step #1: H2(g) 2 H(g) (the covalent bond in the diatomic

hydrogen molecule absorbs energy to break apart into hydrogen

atoms;

2. Step #2: I2(g) 2 I(g) (the covalent bonds in the diatomic iodine

molecule absorb energy to break apart into iodine atoms;

3. Step #3: 2 I(g) + 2 H(g) 2 HI(g) (H would rather bond to I than H -

electronegativity). New H-I bonds form, releasing energy;

4. Net Reaction: H2(g) + I2(g) 2 HI(g) (2 H and 2 I are intermediates

(activated complex) formed in steps #1 & #2, which are then used

in step #3.

Topic: Kinetic Mechanisms

Objective: What types of steps may make a reaction mechanism?



Reaction mechanisms:

A reaction mechanism is similar to different roads that lead to the

same destination. Depending on which road taken, it requires a

different amount of time and/or energy to reach your destination.

The only way to learn exactly which mechanism is truly responsible

for the reaction occurring is to complete an experiment.

The overall rate of any mechanism is determined by the slowest step

in the mechanism.

o The slowest step is the rate-limiting step or rate-determining step.

For any reaction, whichever step is the slowest is the part of the

mechanism that controls the speed of the overall reaction.

Watch Bozeman Chemistry The Rate-Limiting Step - YouTube - 6:46

Topic: Kinetic Mechanisms

Objective: What types of steps may make a reaction mechanism?

https://www.youtube.com/watch?v=MEEg7aHqk6A



Catalysts:

A catalyst is something that speeds up a reaction by modifying a

step in the reaction mechanism with the result being a shortened

mechanism.

i. A catalyst generally lowers the activation energy required to

initiate a mechanism, or a step in a mechanism, allowing the

overall reaction to proceed at a faster rate.

ii. Catalysts are not consumed by the reaction. A catalyst becomes a

temporary part of the mechanism, then is released, and is able to

work again within the mechanism.

iii. When you dissolve an ionic solid (crystal) in water to form an

aqueous ionic solution before undergoing a double replacement

reaction, the water acts as a catalyst as the moving water

molecules increase the rate that the dissolved ions will collide.

Faster collisions = faster reactions!

Topic: Catalysts

Objective: How might we increase the rate of a reaction?



Inhibitor:

An inhibitor is something that slows down a reaction by adding

additional steps to the mechanism with the result generally

lengthening the amount of time needed for the reaction to complete.

i. Inhibitors generally increase the activation energy; if you require

more energy to initiate a reaction, it will take longer to add that

extra energy.

ii. Inhibitors are not consumed by the reaction.

iii. Certain non-ferrous metals form a patina (a combination of oxides,

carbonates, or sulfides) that will inhibit additional oxygen from

gaining access to the metal and causing further oxidation. The

Statue of Liberty is made of copper sheets; the green color we see

is the result of oxygen reacting with the copper and forming the

greenish patina.

Topic: Inhibitors

Objective: How might we decrease the rate of a reaction?

http://thebrightnessofbeing.com/art/copper-and-patina-new-versus-old/attachment/copper-patina-color-weathering/



Reaction Rate:

The speed at which a reactant is chemically changed into a product is

known as the Reaction Rate.

i. The rate of reaction is generally measured in terms of amount of

reactants consumed over a period of time.

ii. Reaction rate is measured by the number of effective molecular

collisions that occur per unit of time. The greater the number of

effective collisions in a period, the faster the reaction rate.

iii. Any action that speeds up the rate of effective collisions will

increase the rate of the reaction.

Topic: Reaction Rates

Objective: How do we measure the amount of a reaction over time?



Ionic Reactions:

1. Ionic reactions occur quickly;

2. Ionic reactions occur in liquid water, and the nature of liquid water

allows a large number of collisions in aqueous solution;

3. Aqueous solutions of ions have the ionic bonds already broken, and

they can readily make new ionic bonds;

4. Precipitate reactions occur almost immediately upon mixing

solutions.

Covalent reactions:

1. Covalent reactions occur slowly;

2. Covalently bonded molecules have many different types of bonding,

and that can make some covalent reactions very slow;

3. Covalently bonded molecules may have single, double, or even triple

bonds that require large amounts of energy to break;

4. Covalent molecules used in cellular processes use biological

catalysts, known as enzymes, to speed up the process of breaking

down these complex bonds more efficiently and more quickly.

Topic: Rate and Reactant Nature

Objective: How may the type of reactants change reaction rate?



Determining Molecular Polarity:

Temperature:

You should recall that temperature is the measure of the average

kinetic energy of particles in a system. Kinetic energy is moving

energy; if you increase the temperature, you increase the kinetic

(moving) energy of the system.

i. If you increase the kinetic (moving) motion, you will increase the

number of possible effective collisions.

ii. We add heat energy (Bunsen burner; hotplate) to increase the

temperature (kinetic energy) so we can complete a reaction in a

reasonable amount of time.

Lower Temperature Higher Temperature

Topic: Rate and Temperature

Objective: How may the temperature change reaction rate?



Concentration:

Concentration of reactants affects the orientation of colliding

particles. The more particles there are in a given volume, the greater

the chance they will collide.

i. As particles will only react if they collide in an orientation that

allows them to bond, having more particles increases the random

odds of the orientation being correct.

ii. If the greater concentration allows for more random orientations, a

greater number of possible effective collisions may result.

iii. Think of cars in a parking lot; if more cars are moving around, the

greater the chance for a collision than if the parking lot is empty.

Lower Concentration Higher Concentration

Topic: Rate and Concentration

Objective: How may concentration affect the rate of reaction?



Surface Area:

Surface area of the reactants affects the rate of reaction much like

concentration. If you increase the surface area, you increase the

number of properly orientated collision sites.

i. This property only works with solids, as liquids and gases have

maximum surface area already.

ii. Crushing a solid will increase the surface area while the volume

stays the same, effectively increasing the collision locations.

iii. This property you already learned in Earth Science in regards to

weathering; larger rocks weather slower than smaller rocks, as the

surface area to volume ratio for larger rocks is lower.

Larger Particles (less area) Smaller particles (more area)

Topic: Rate and Surface Area

Objective: How may surface area affect the rate of reaction?

Larger particles have a lower surface area : volume ratio

Smaller particles have a higher surface area : volume ratio



Notes page:



Kinetics Practice Regents Questions: (ungraded)

1. In most aqueous reactions as temperature of the system increases, the

effectiveness of particle collisions

a) Increases b) Decreases c) Remains the same

2. Given the reaction of Mg(s) + 2 H2O(l) Mg(OH)2(aq) + H2(g), at which

temperature will the reaction occur at the greatest rate?

a) 25°C b) 50°C c) 75°C d) 100°C

3. Given the reaction of A(s) + B(aq) C(aq) + D(s), which change below would

increase the rate of this reaction?

a) A decrease in pressure

b) An increase in pressure

c) A decrease in temperature

d) An increase in temperature

4. As the concentration of reacting particles increases, the rate of reaction

generally

a) Decreases b) Increases c) Remains the same

5. Given the reaction of A2(g) + B2(g) 2 AB(g) + heat, an increase in the

concentration of A2(g) will

a) Increase the production of B2(g)

b) Decrease the production of AB(g)

c) Decrease the frequency of collisions between A2(g) and B2(g)

d) Increase the frequency of collisions between A2(g) and B2(g)

6. At STP, which 4.0 gram sample of zinc metal will react the fastest with dilute

hydrochloric acid?

a) Bar

b) Flat sheet

c) Powdered

d) Irregular lump



Student name: _________________________ Class Period: _______

Please carefully remove this page from your packet to hand in.

Kinetics homework

1. Which statement below explains why the speed of a chemical reaction is

increased when the surface area of a reactant is increased?

a) This change increases the concentration of the reactant.

b) This change increases the density of the reactant particles.

c) This change exposes more reactant particles to a possible collision.

d) This change alters the electrical conductivity of the reactant particles.

2. A catalyst works by

a) Increasing the energy released during a reaction.

b) Increasing the potential energy of the reactants.

c) Decreasing the potential energy of the reactants.

d) Decreasing the activation energy of the reaction.

3. An increase in the surface area of reactants in a reaction will result in

a) A decrease in the heat of reaction.

b) An increase in the heat of reaction.

c) A decrease in the rate of the reaction.

d) An increase in the rate of the reaction.

4. For a given reaction mechanism (A to E) and the table of observed durations,

which step is the rate-determining step?

Step Duration Rate Determining step?

A 0.0002 secs

B 0.0343 secs

C 1.7771 secs

D 0.0089 secs

E 0.9880 secs



5. Explain how the following would affect the reaction rate for the below

examples.

(Increase, Decrease, Remains the same)

a) Adding additional N2(g) to the reaction N2(g) + 3 H2(g) 2 NH3(g): ___

b) Removing H2(g) from the reaction N2(g) + 3 H2(g) 2 NH3(g): ___

c) Increasing the pressure on the reaction N2(g) + 3 H2(g) 2 NH3(g): ___

d) Use powdered NaCl(s) and not large crystals of NaCl(s) in the reaction: ___

e) Adding water as a catalyst to a double replacement reaction: ___

f) Adding CuSO4(aq) inhibitor to a reaction of acid and metal: ___

g) Increasing pressure on Na(s) + ZnSO4(aq) Na2SO4(aq) + Zn(s): ___

6. A student places three separate samples of sugar in three different insulated

cups each containing 50.0 mL of distilled water at 50.0°C. Sample A is a single

cube of sugar, Sample B is granulated sugar, and Sample C is powered sugar.

Which sample will dissolve the slowest? Explain your answer in terms of

effective collisions.

7. A student places three identical 5.0 g gram samples of sugar into three different

insulated cups each containing 50.0 mL of distilled water. Cup A contains

water at 10°C, Cup B contains water at 20°C, and Cup C contains water at

30°C. In which cup will the sugar dissolve the fastest? Explain your answer

in terms of effective collisions.



Energy Changes in Reactions:

In chemical reactions, reactants form products with either an

associated release or an associated absorption of energy.

Enthalpy is the total amount of Potential Energy (PE) stored in a

substance.

i. By changing the amount of PE involved in a chemical reaction,

we can change the enthalpy of the system.

a) Exothermic reactions release energy, so the enthalpy of the

system lowers.

b) Endothermic reactions store energy, so the enthalpy of the

system rises.

A + B C + D A + B C + D

Reactants Products

Topic: Energy Changes & Enthalpy

Objective: How does energy change during chemical reactions?

Energy Energy

Enthalpy

increases

Enthalpy

decreases



Endothermic Reactions:

During an endothermic reaction, energy is absorbed as a reactant.

A + B + Energy C + D

i. Energy can be considered as an additional reactant;

ii. Heat energy is absorbed by the reactants;

iii. This heat energy is absorbed from the environment around the

reaction, and the heat in the surroundings decreases;

iv. The products of an endothermic reaction have more energy than

did the starting reactants;

v. Kinetic energy is stored in the created bonds of the products;

vi. The excess stored energy makes the products of endothermic

reactions chemically unstable and highly reactive;

vii. Examples of endothermic reactions include nitroglycerine,

trinitrotoluene (TNT), and other explosives that have high amounts

of energy stored within their chemical bonds and are easily reactive.

Topic: Endothermic Reactions

Objective: How is energy absorbed during a reaction?

http://www.google.com/url?sa=i&source=images&cd=&cad=rja&uact=8&ved=0CAgQjRw&url=http://pixshark.com/exothermic-reaction-examples.htm&ei=AugAVY7lKorLsAStnoDwCw&psig=AFQjCNHt_tY8PjmysgM9GJygXiUH-NRmUQ&ust=1426209154788404



Endothermic Reaction example:

For the General Reaction of A + B C:

i. If HA (Heat Energy of Reactant ‘A’) is equal to 40 kJ, and HB = 20 kJ,

then the reactants together have a total heat energy of 60 kJ. If the

given HC = 110 kJ, then (110 kJ – 60 kJ) = 50 kJ of heat energy was

absorbed by the reactants from the surroundings to form the product.

ii. This reaction can therefore be rewritten to show the energy change as:

A(40kJ) + B(20 kJ) + Absorbed Energy(50 kJ) C(110 kJ)

(40 kJ + 20 kJ + 50 kJ = 110 kJ) 110 kJ

iii. Note that there is a total of 110 kJ of heat energy on both sides of the

equation. Equal energy amounts on both sides of the equation

satisfies the Law of Conservation of Energy.

110 kJ 110 kJ

http://www.google.com/url?sa=i&rct=j&q=&esrc=s&source=images&cd=&cad=rja&uact=8&ved=0CAcQjRw&url=http://chem103csu.wikispaces.com/First%2BLaw%2Bof%2BThermodynamics&ei=OPUAVfTUJovfUb2GhNgH&bvm=bv.87920726,d.cWc&psig=AFQjCNG_u49xl-MUAjGLhjNxbsrfkcsVpw&ust=1426212449394221



Endothermic Reaction example:

For a reaction from the NYS Chemistry Reference Table I:

N2(g) + O2(g) 2 NO(g) ∆H = + 182.6 kJ

i. The plus (+) sign in front of the 182.6 kJ indicates that this reaction is

endothermic, and that energy was absorbed. You can add the energy

for the reaction listed in the Reference Table I to the reactants side, as

it becomes a part of the products.

N2(g) + O2(g) + 182.6 kJ 2 NO(g)

ii. This new reaction equation shows that 1 mole of N2(g) and 1 mole of

O2(g) absorb 182.6 kJ of energy during the formation of 2 moles of

NO(g).

iii. 2 moles of NO(g) therefore have 182.6 kJ more energy stored in bonds

than 1 mole of N2(g) and 1 mole of O2(g) have in their combined bonds.

iv. NO(g) is more unstable than N2(g) and O2(g).



Exothermic Reactions:

During an endothermic reaction, energy is released as a product.

A + B C + D + Energy

i. Energy can be considered as an additional product;

ii. Heat energy is released from the reactants;

iii. This heat energy is released into the environment around the

reaction, and the heat in the surroundings increases;

iv. The products of an exothermic reaction have less energy than did

the starting reactants;

v. Potential energy in the bonds of the reactants is lost from the

products;

vi. The lost energy makes the products of exothermic reactions

chemically more stable and less reactive than their reactants;

vii. An example of an exothermic reaction is the burning of paper or

wood. After they have burned, the ash that remains has less stored

energy and is more stable (inflammable).

Topic: Exothermic Reactions

Objective: How is energy released during a reaction?

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Exothermic Reaction example:

For the General Reaction of A + B C:

i. If HA = 60 kJ and HB = 40 kJ, then the reactants together have a total

heat energy of 100 kJ. If HC = 30 kJ, then (100 kJ - 30 kJ) = 70 kJ of

heat energy was released from the reactants to the surroundings while

forming the products.

ii. This reaction can therefore be rewritten to show the energy change as:

A(60kJ) + B(40 kJ) C(30 kJ) + Released Energy(70 kJ)

(60 kJ + 40 kJ = 100 kJ) (30 kJ - 70 kJ = 100 kJ)

iii. Note that there is a total of 100 kJ of heat energy on both sides of the

equation. Equal energy amounts on both sides of the equation

satisfies the Law of Conservation of Energy.

100 kJ 100 kJ

http://www.google.com/url?sa=i&rct=j&q=&esrc=s&source=images&cd=&cad=rja&uact=8&ved=0CAcQjRw&url=http://chem103csu.wikispaces.com/First%2BLaw%2Bof%2BThermodynamics&ei=OPUAVfTUJovfUb2GhNgH&bvm=bv.87920726,d.cWc&psig=AFQjCNG_u49xl-MUAjGLhjNxbsrfkcsVpw&ust=1426212449394221



Exothermic Reaction example:

For a reaction from the NYS Chemistry Reference Table I:

C(s) + O2(g) CO2(g) ∆H = -393.5 kJ

i. The minus (-) sign in front of the 393.5 kJ indicates that this reaction

is exothermic, and that energy was released. You can add the energy

for the reaction listed in the Reference Table I to the products side, as

it was lost from the reactants.

C(s) + O2(g) CO2(g) + 393.5 kJ

ii. This new reaction equation shows that C(s) + O2(g) together release

393.5 kJ of energy during the mechanism that forms CO2(g).

iii. 1 mole of CO2(g) therefore has 393.5 kJ less energy stored in its bonds

than 1 mole of C(s) and 1 mole of O2(g) have in their combined bonds.

iv. CO2(g) is more stable than N2(g) and O2(g) are alone.

Watch Energy & Chemistry: Crash Course Chemistry #17 - YouTube - 9:25

https://www.youtube.com/watch?v=GqtUWyDR1fg



The reaction on the previous page shows the energy released during

the SYNTHESIS of one mole of CO2(g). The mechanism for this

synthesis reaction releases 393.5kJ of energy into the environment.

We can use this quantitative amount of energy for the synthesis of

one mole of CO2(g) to calculate the energy released by any molar

ratio.

Example:

i. How much energy is released during the mechanism to synthesize

2.3 moles of CO2(g)?

2.3 moles of CO2(g) x = - 905.05 kJ - 910 kJ

ii. The above equation states that 910 kJ of energy are released when

2.3 moles of CO2(g) are synthesized.

iii. What of the decomposition of CO2(g)? Decomposition is the

reverse process of synthesis, so the same given energy would be

used, but the sign would be the OPPOSITE.

Example:

2.3 moles of CO2(g) x = + 905.05 kJ + 910 kJ

i. The above equation states that 910 kJ of energy are absorbed when

2.3 moles of CO2(g) are decomposed.

∆H of Decomposition = (opposite sign) ∆H of Synthesis

Topic: Molar Heat Energy

Objective: How could you calculate molar heat energy?

-393.5 kJ

mole of CO2(g)

+ 393.5 kJ

mole of CO2(g)



Kinetics Practice Regents Questions: (ungraded)

1. Given the reaction of A + B C + D + heat, which statement best describes

this reaction?

a) The forward reaction is exothermic, and the reverse reaction is always

exothermic.

b) The forward reaction is exothermic, and the reverse reaction is always

endothermic.

c) The forward reaction is exothermic, and the reverse reaction can be either

exothermic or endothermic.

d) The forward reaction is endothermic, and the reverse reaction can be either

endothermic or exothermic.

Salt A and Salt B were dissolved in separate 100-mL samples of water. The

water temperatures were measured and recorded as shown in the table below.

Salt A Salt B

Initial water temperature: 25.1˚C 25.1˚C

Final water temperature: 30.2˚C 20.0˚C

2. Which statement is a correct interpretation of these results?

a) The dissolving of only salt B was exothermic.

b) The dissolving of only salt A was endothermic.

c) The dissolving of both salt A and salt B was endothermic.

d) The dissolving of salt A was exothermic and the dissolving of salt B was

endothermic.

3. Given the reaction of Fe(s) + S(s) FeS(s) + energy, which statement about this

reaction is true?

a) It is exothermic.

b) It is endothermic.

c) The potential energy of the reactants is the same as the potential energy of

the product.

d) The potential energy of the reactants is lower than the potential energy of the

product.



Potential Energy Diagrams:

A Potential Energy (PE) Diagram is a graphical depiction of the flow

of potential energy as a reaction goes from start to finish.

A basic Potential Energy flow process consists of these steps:

i. All reactions start with reactants, therefore the PE begins at the Heat

of Reactants (HR) level;

ii. Activation Energy (EA) must be added to the reactants to initiate the

reaction mechanism:

a) EA is the Energy needed to have effective collisions between

reactant particles;

b) EA raises the level of PE in the reaction to the Heat of Activated

Complex (HAP - intermediate step) level;

iii. The Activated Complex is VERY short-lived and breaks down

quickly to form the products of the reaction. This final step lowers

the level of PE to the Heat of Products (HP) level.

Watch Bozeman Chemistry Activation Energy - YouTube - 4:51

Topic: Potential Energy Diagrams

Objective: How do we depict the flow of energy within a reaction?

https://www.youtube.com/watch?v=YacsIU97OFc



Creating a Potential Energy Diagram:

Seven items must be included and labeled in a PE Diagram:

A. The axes:

1. y-axis is PE in kJ;

2. x-axis is the Reaction Coordinate (time from start to finish; this

time is unmeasured, therefore it has NO UNITS!

B. The Potential Energy Levels:

3. Heat of Reactants;

4. Heat of Products;

5. Heat of Activated Complex

C. The Energy Changes during the reaction:

6. Activation Energy (EA) as an arrow from HR to HAC;

7. ∆H as an arrow from HR to HP (∆H shows NET energy change)

Topic: Creating PE Diagrams

Objective: How do we make a Potential Energy Diagram?



Interpreting Endothermic Potential Energy Diagrams:

Endothermic Reactions: Energy is absorbed by the reactants as they

form products, so the NET amount of PE will increase.

For the PE Diagram for reaction A + B + 50 kJ C as shown below:

Experiments have determined that the combined heats of reactants

‘A’ and ‘B’ is 60 kJ. This is the combined Heat of Reactants (HR).

Experiments have also determined that the heat of product ‘C’ is 110

kJ. This is the Heat of Products (HP). Additional experiments show

that ‘A’ and ‘B’ require 70 kJ of energy to be input before ‘A’ and

‘B’ will react. This 70 kJ is the Activation Energy (EA). For this

Topic: Interpreting PE Diagrams

Objective: How do we read a Potential Energy Diagram?



reaction, adding 70 kJ to reactants ‘A’ and ‘B’ yields the Activated

Complex, which has its own Heat of Activated Complex (HAC),

which is ALWAYS the highest energy level in ANY reaction. This

reaction started with reactants totaling 60 kJ and a product having 110

kJ, which means we have a NET increase of 50 kJ from the reactants.

The + 50 kJ is called the Heat of Reaction, and shown as ∆H. This is

taking energy from the surroundings and storing WITHIN the bonds

of product ‘C’, meaning it is an endothermic reaction.

i. The bonds in ‘C’ have more energy and are less stable than bonds

in ‘A’ and ‘B’;

ii. The temperature of the surroundings has DECREASED as

surrounding kinetic energy was stored within the bonds of ‘C’.

If we added a catalyst to this reaction, the activation energy (EA)

would decrease by removing step(s) from the mechanism. As a

result, the HAC would also decrease. The opposite would happen

should an inhibitor be added to the reaction.



Interpreting Exothermic Potential Energy Diagrams:

Endothermic Reactions: Energy is released from the reactants as they

form products, so the NET amount of PE will decrease.

For the PE Diagram for reaction X + Y Z + 40 kJ as shown below:

Experiments have determined that the combined heats of reactants

(HR) for ‘X’ and ‘Y’ is 70 kJ. Experiments have also determined that

the heat of product (HP) of ‘Z’ is 30 kJ. Additional experiments show

that ‘X’ and ‘Y’ require 80 kJ of energy to be input before ‘X’ and

‘Y’ will react. This 80 kJ is the Activation Energy (EA). For this

reaction, adding 80 kJ to reactants ‘X’ and ‘Y’ yields the Activated

Topic: Interpreting PE Diagrams

Objective: How do we read a Potential Energy Diagram?



Complex, which has its own Heat of Activated Complex (HAC),

which is ALWAYS the highest energy level in ANY reaction. This

reaction started with reactants totaling 70 kJ and a product having 30

kJ, which means we have a NET decrease of 40 kJ from the reactants.

The - 40 kJ is called the Heat of Reaction, and shown as ∆H. This is

taking energy from the bonds in ‘X’ and ‘Y’ and releasing that energy

into the surroundings as kinetic energy, meaning it is an exothermic

reaction.

i. The bonds in ‘Z’ have less energy and are more stable than bonds

in ‘X’ and ‘Y’;

ii. The temperature of the surroundings has INCREASED as kinetic

energy was released from the bonds of ‘X’ and ‘Y’.

If we added a catalyst to this reaction, the activation energy (EA)

would decrease by removing step(s) from the mechanism. As a

result, the HAC would also decrease. The opposite would happen

should an inhibitor be added to the reaction.



Catalysts and Inhibitors:

For the PE diagram to the LEFT, Curve 1 would be the control (normal) PE curve of Activation Energy. Curve 2 would be the catalyzed (lower) PE curve of Activation Energy.

For the PE Diagram to the RIGHT, Curve 2 would be the control (normal) PE curve of Activation Energy. Curve 1 would be the inhibited (higher) PE curve of Activation Energy.



Reversing Chemical Reactions:

Chemical reactions are almost all reversible.

Synthesis reactions (putting together) may be reversed as

decomposition (breaking down) reactions.

With enough kinetic energy, almost any chemical reaction may be

reversed.

When a reaction is reversed, the products become the reactants. The

reaction coordinate goes from right to left instead of left to right.

Note that the forward (exothermic) reaction has a very small EA, but

the reverse (endothermic) reaction has a much greater EA. On the

diagram below the forward (exothermic) reaction should have a

downward ∆H arrow, as the forward reaction goes from higher PE

(Point H) to lower PE (Point E). On this diagram, the ∆H arrow for

the reverse (endothermic) reaction goes upwards instead.

Topic: Reversing Reactions

Objective: What can be done to reverse a chemical reaction?

Forward

(exothermic)

Reverse

(endothermic)

H ^

E ^



Endothermic Reaction:

Exothermic Reaction:



Potential Energy Practice Regents Problems: (ungraded)

1. Which information about a chemical reaction is provided by a potential energy

diagram?

a) The change in solubility of the reacting substances.

b) The oxidation states of the reactants and the products.

c) The average kinetic energy of the reactants and the products.

d) The energy released or absorbed during a chemical reaction.

The potential energy diagram below represents a chemical reaction.

2. Which arrow represents the activation energy of the forward reaction above?

a) A b) B c) C d) D

3. Given the reaction of S(s) + O2(g) SO2(g) + energy, which diagram shown below

best represents the potential energy changes for this reaction?

4. The activation energy required for a chemical reaction may be decreased by

a) Adding more reactant to the mechanism

b) Increasing the surface area of the reactant

c) Increasing the temperature of the reactant

d) Adding a catalyst to the reaction mechanism Cont’d next page



In the diagram below, which letter represents the activation energy for the reverse

reaction?

a) A b) B c) C d) D

5. In a potential energy diagram, the difference between the potential energy of the

products and the potential energy of the reactants is equal to the

a) Heat of reaction

b) Entropy of reaction

c) Activation energy of the reverse reaction

d) Activation energy of the forward reaction

The potential energy diagram below shows the reaction for X + Y Z:

6. When a catalyst is added to the reaction, it will change the value of

a) 1 and 2

b) 2 and 3

c) 1 and 3

d) 3 and 4



Student name: _________________________ Class Period: _______

Please carefully remove this page from your packet to hand in.

Potential Energy Diagrams Homework:

Below are three partially completed reactions. Find the appropriate Heat of Reaction in

Reference Table I and write the Heat of Reaction in the proper space, leaving the other space

BLANK. State if the reaction is EXOTHERMIC or ENDOTHERMIC based on Table I.

Identify the products as being STABLE or UNSTABLE when compared to the reactants for that

reaction. 9 pts.

Given Reaction:

EXOthermic or

ENDOthermic?

Products STAble or

UNSTAble?

1. N2(g) + 3 H2(g) + _______ 2 NH3(g) + _______

2. N2(g) + 2 O2(g) + _______ 2 NO2(g) + _______

3. H2(g) + I2(g) + _______ 2 HI(g) _______

Answer questions #4 through #13 using Reference Table I for the formation of NO(g) from its

elements. 1 pt. ea. except where noted.

4. How many moles of NO(g) are formed in the reaction on Table I? _______moles

5. What is the Heat of Reaction for the formation of NO(g) per mole? _______ kJ/mole

6. Show your calculations for the above answer:

7. How much energy is absorbed as 4.5 moles of NO(g) are formed? _______ kJ


9. How much energy is released as 2.7 moles of NO(g) are decomposed? _______ kJ


Samples of N2(g) and O2(g) are reacted to form 0.010 moles of NO(g).

11. How much energy will be absorbed by the reaction given above? _______ kJ


Cont’d next page



13. If the reaction in question #11 was carried out in a calorimeter in 50.0 g of H2O at an initial

temperature of 20.0˚C, what would the final water temperature be? Show ALL work. 3 pts.

Using the Potential Energy Diagram below, answer the following questions.

14. What is the heat of the reactants?

a) 30 kJ b) 110 kJ c) 140 kJ d) 160 kJ

15. What is the heat of the products?

a) 30 kJ b) 110 kJ c) 140 kJ d) 160 kJ

16. What is the heat of the activated complex for the uncatalyzed reaction?

a) 30 kJ b) 110 kJ c) 140 kJ d) 160 kJ

17. What is the activation energy for the catalyzed reaction?

a) 30 kJ b) 110 kJ c) 140 kJ d) 160 kJ

18. What is the ∆H of this reaction?

a) +80 kJ b) -80 kJ c) +130 kJ d) -130 kJ

19. What type of reaction is shown on this diagram?

a) Exothermic

b) Endothermic

c) Both types

d) Neither type

20. Using the dashed catalyzed reaction line as a guide, draw a dotted (● ● ●) curved line on the

above diagram to indicate what the PE curve would look like if a catalyst were added.



Notes page:

Unit 9: Kinetics, Thermodynamics, & Equilibrium-Lecture ... U… · Unit 9: Kinetics,...

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