Reaction Energy and Reaction Kinetics Chapter 17 Notes.

Reaction Energy and Reaction Kinetics

Chapter 17 Notes

Thermochemistry• Thermochemistry is the

study of the transfers of energy as heat that accompany chemical reactions and physical changes.– Endothermic and exothermic

reactions– Phase changes and specific

heat

Calorimeter

• The energy absorbed or released as heat in a chemical or physical change is measured in a calorimeter.

Heat and Temperature

• Heat (q): energy transferred between samples of matter because of a difference in their temperatures.

• Temperature: a measure of the average kinetic energy of the particles in a sample of matter.– Higher temperature = faster particles– Lower temperature = slower particles

Heat• As an object absorbs heat

– the particles increase in kinetic energy– the temperature of the object increases

• Heat always flows from a __________object to a __________ object

hottercolde

r

Thermochemical Systems• System: the collection of matter that

you are studying– The liquid in which a reaction is taking

place– The interior of a calorimeter

• Surroundings: the area directly around the system that is capable of absorbing energy from or releasing energy to your system– The beaker holding the liquid and the air

around the beaker– The walls of the calorimeter and the air

around them

Law of Conservation of Energy

• Energy cannot be created or destroyed.

• It can be moved from place to place

Endothermic Process

• An endothermic process absorbs heat from its surroundings– Instant cold packs– Vinegar and baking

soda– Boiling water

Exothermic Process

• A system that releases energy to its surroundings– Diluting sulfuric

acid– Water freezing into

ice– The heating unit

included in MREs

Heat of Reaction

• The energy transferred as heat during a chemical reaction is called the heat of reaction.

• This heat is shown as either a reactant or a product in a thermochemical equation.– CH4 + 2 O2 → CO2 + 2 H2O + 1565 J– CH3COOH + NaHCO3 + 1245 J → NaC2H3O2 + CO2 + H2O

Exothermic

Endothermic

Enthalpy

• Enthalpy change: the amount of energy transferred as heat during a process occurring at constant pressure

• ΔH = Hproducts – Hreactants

– Positive for endothermic reactions– Negative for exothermic reactions

Enthalpy and Thermochemical Equations

• Exothermic reaction

• CH4 + 2 O2 → CO2 + 2 H2O + 1565 J

ΔH =-1565 J

• Endothermic reaction• CH3COOH + NaHCO3 + 1245 J → NaC2H3O2 + CO2 + H2O

ΔH =1245 J

Entropy• Entropy: the amount of disorder

within a system– High entropy = chaotic

• The plasma inside the sun• Typical kid’s bedroom

– Low entropy = very orderly• Ice• My movie collection

• Processes occur spontaneously (without an external driving force) if they increase the entropy of the system.

Units for Heat• calorie (cal) = the amount of heat

needed to raise the temperature of 1 gram of water by 1oC. The “c” is ALWAYS lowercase for this unit.

• Calorie (Cal) = 1 kilocalorie (kcal). This is what we mean when we talk about food – how much energy is in the food. The “C” is ALWAYS capitalized!

• Joule = SI unit for energy, including heat

Heat Capacity

• The amount of heat needed to change an object’s temperature.

• This depends on– Size of the object– Material of the object– How big is the temperature change

• To allow more direct comparisons we use specific heat capacity instead.

Specific Heat Capacity

• The amount of heat need to change the temperature of 1 gram of a material by 1oC– For water, c = 4.18 J/g∙oC– This means it takes 4.18 joules of energy to

raise the temperature of 1 gram of water from 20oC to 21oC

• You have been given a table of specific heats. This needs to go in the references section of your binder.

Calculating Specific Heat

ΔTm

qc

Heat (J or cal)

Change in temperature =

Tfinal – Tinitial (oC or K)

Mass (g)

Specific heat (J/g∙oC or cal/g∙oC)

Example 1

• The temperature of a piece of copper with a mass of 95.4 grams increases from 25.0oC to 48.0oC when the metal absorbs 849 J of heat. What is the specific heat of copper?

ΔTm

qc

C0.25C0.4895.4g

849Jc

oo

Cg

J387.0c

o

Example 2

• How much heat is required to raise the temperature of 250.0 g of mercury 52.0oC?

Tcmq C0.52g0.250

Cg

J140.0q o

o

J1820q

Example 3

• What is the temperature change when 2543 J of heat are added to 76.4 g of alcohol?

mc

qT

g4.76CgJ

44.2

J2543T

o

C6.13T o

Example 4

• The temperature of a 283 gram sample increases 11.2oC by absorbing 1204 J of heat. What material is the sample?

ΔTm

qc

C2.11283g

1204Jc

o Cg

J380.0c

o

Reaction Kinetics

• Chemical kinetics: the study of reaction rates and mechanisms– How quickly reactions take place– The actual process of the reaction

Reaction Mechanisms

• Mechanism: the step-by-step sequence of reactions by which the overall chemical change occur.

• Rate determining step: the slowest part of the reaction mechanism– It takes the longest, thus determining the

speed of the overall reaction– When baking slice and bake cookies, the

actual baking is the longest step. It determines how long it takes to bake those cookies.

Reaction Mechanisms: An Example

The reaction 2 A2B + CD4 → CB2 + 4 AD doesn’t happen all at once. Instead it follows the following reaction mechanism.

1st: 2 A2B → 4 A- + 2 B2+ (very fast)

2nd: 4 A-+CD4 → C4+ + 4 AD (very slow)

3rd: C4+ + 2 B2+ → CB2 (very fast)

What is Needed for a Reaction?

• The reactants in question must collide with:– Enough energy– Favorable orientation

• Anything that increases the likelihood of those two conditions being met will increase the reaction rate

Kinetics and Conditions

1. Nature of reactants – some elements react more quickly than other elements

2. Surface area – more surface area means more opportunities for favorable collisions

3. Temperature – hotter reactants will collide with more energy

4. Concentration – more concentrated solutions allow for more frequent collisions

Chemical Reaction Rates

• Reaction rate: how quickly the concentration of a substance in a chemical reaction is changing. It’s measured in M/s.

• Instantaneous reaction rate: the rate at any given instant of time

• Initial reaction rate: the rate when a reaction is first started

Reaction Rates

• For reactants, the reaction rate is how quickly the concentration is decreasing

• For products, the reaction rate is how quickly the concentration is increasing

Rate Expressions

• Rate law/rate expression: an equation that relates reaction rate and concentrations of reactants

**Important note**You cannot predict the values of n, m, or

p from the reaction equation. The rate law must be determined experimentally!

pmn CBAkrate

Rate Expressions• Rate constant, k:

– A proportionality constant that allows the direct calculation of the rate from the concentrations of the reactants and catalyst

– The units of k vary according to the rate law: you use the units needed to make the units for the rate what you need. They will usually be one of the following:

s

1

sM

1

sM

12

The Order of Reaction

• Total reaction order: the sum of all of the exponents in the rate law

• Common orders – – 0th: rate = k– 1st: rate = k[A]– 2nd: rate = k[A][B] or rate = k[A]2

– 3rd: rate = k[A][B][C], rate = k[A]2[B] or rate = k[A]3

NOTE: higher reaction orders have more possibilities for the rate law

Determination of the Rate Expression

• We need to determine how each change in concentration affected the reaction rate.

Run[A] (M)

[B] (M)

[C] (M)

Rate (M/s)

1 0.10 0.10 0.10 0.00080

2 0.20 0.10 0.10 0.0016

3 0.10 0.20 0.10 0.00080

4 0.10 0.10 0.20 0.0032


• First, what changes between run 1 and run 2?– The concentration of A is doubled

Run[A] (M)

[B] (M)

[C] (M)

Rate (M/s)

1 0.10 0.10 0.10 0.00080

2 0.20 0.10 0.10 0.0016

3 0.10 0.20 0.10 0.00080

4 0.10 0.10 0.20 0.0032


• What change does this bring in the rate?– The rate is also doubled– So a change in [A] causes the same change in

rate– Rate [A]

Run [A] (M) [B] (M) [C] (M) Rate (M/s)

1 0.10 0.10 0.10 0.00080

2 0.20 0.10 0.10 0.0016

3 0.10 0.20 0.10 0.00080

4 0.10 0.10 0.20 0.0032


• Next, what changes between run 1 and run 3?– The concentration of B is doubled

Run[A] (M)

[B] (M)

[C] (M)

Rate (M/s)

1 0.10 0.10 0.10 0.00080

2 0.20 0.10 0.10 0.0016

3 0.10 0.20 0.10 0.00080

4 0.10 0.10 0.20 0.0032


• What change does this bring in the rate?– The rate is not changed– So a change in [B] does not change the rate


1 0.10 0.10 0.10 0.00080

2 0.20 0.10 0.10 0.0016

3 0.10 0.20 0.10 0.00080

4 0.10 0.10 0.20 0.0032


• Next, what changes between run 1 and run 3?– The concentration of C is doubled

Run[A] (M)

[B] (M)

[C] (M)

Rate (M/s)

1 0.10 0.10 0.10 0.00080

2 0.20 0.10 0.10 0.0016

3 0.10 0.20 0.10 0.00080

4 0.10 0.10 0.20 0.0032


• What change does this bring in the rate?– The rate is quadrupled– So a change in [C] is squared in the rate– Rate [C]2


1 0.10 0.10 0.10 0.00080

2 0.20 0.10 0.10 0.0016

3 0.10 0.20 0.10 0.00080

4 0.10 0.10 0.20 0.0032


• When we put it all together we get – Rate = k[A][C]2

Run[A] (M)

[B] (M)

[C] (M)

Rate (M/s)

1 0.10 0.10 0.10 0.00080

2 0.20 0.10 0.10 0.0016

3 0.10 0.20 0.10 0.00080

4 0.10 0.10 0.20 0.0032

Reaction Energy Diagrams

• Shows the changes in energy as a reaction takes place

• Can be used to analyze many energy changes that occur during a reaction


• A: energy held by the activated complex

• B: energy of the reactants

• C: energy of the products

• F: heat of reaction• I: activation

energy


• This is an exothermic reaction – the reactants have more energy than the products, so energy has been released


• This is an endothermic reaction – the reactants have less energy than the products, so energy has been absorbed

Energy During Reaction

• Transition state energy: the energy held by the activated complex during a reaction

• Activated complex: a molecule that temporarily exists during the process of a reactions (i.e., it is created in the reaction mechanism)

• Activation energy: the energy required to start a reaction (the spark from a match when lighting a gas burner, for example)

Changing the Energy Requirements

• Catalyst: a substance that speeds a reaction up by decreasing the activation energy required– Homogeneous catalyst: in the same state

as the reactants– Heterogeneous catalyst: in a different state

than the reactants

• Inhibitor: a substance that slows a reaction down by increasing the activation energy required


• D: catalyzed reaction

• E: inhibited reaction

Reaction Energy and Reaction Kinetics Chapter 17 Notes.

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