DMA 4/11/11
Write the balanced chemical
+ equation for the reaction of
aluminum and hydrochloric
acid (HCl) to produce
aluminum chloride and
hydrogen gas.
+
DMA - Monday, May 04, 2009
 List
three ways you might increase the rate
of a chemical reaction? Give some
examples.
+
Tuesday, April 12, 2011
 Class
Notes - Rates of Reactions
 Homework
Notes and Section Review Chapter 22
DMA 4/12/11
Get your notebook and a pencil
and get ready to take some
+ notes!
DMA 4/12/11
Name 2 things you think
+ could increase the rate at
which a chemical reaction
happens.
+
Chemical Kinetics
The area of chemistry that concerns reaction rates.
+
Chemical Kinetics
 The
area of chemistry concerned with the speed at
which reactions occur is called chemical kinetics.
 Reaction
rate is the change in concentration of
reactants and products in a certain amount of time.
Average rate of reaction = Δ [reactant or product]
Δ time
+
Collision Theory
Key Idea: Molecules must collide to react.
•In order to react molecules and atoms must come in
contact with each other.
•They must hit each other hard enough to react.
•Anything that increase these things will make the
reaction faster.
Energy
Reactants
Products
Reaction coordinate
Energy
Activation Energy Minimum energy to
make the reaction
happen
Reactants
Products
Reaction coordinate
Energy
Reactants
Overall energy
change
Products
Reaction coordinate
+ Endothermic Reactions
+ Exothermic Reactions
+
Catalysts
H H

Hydrogen bonds to surface of metal.

Break H-H bonds
H
H H
H H
H
Pt surface
+
Catalysts
H
H
H
C
C
H
H H
H
H
Pt surface
+

Catalysts
The double bond breaks and bonds to the catalyst.
H
H
H
C
H
C
H
H
Pt surface
H H
+

Catalysts
The hydrogen atoms bond with the carbon
H
H
H
C
H
C
H
H
Pt surface
H H
+
Catalysts
H
H
H
H
C
C
H
H
H
Pt surface
H
+
Catalysts
+ Exothermic Reaction with a Catalyst
+ Endothermic Reaction with
a Catalyst
DMA 4/13/11
+ What is a catalyst? How
does it affect a chemical
reaction?
Remember to turn in your homework!
+
Your To-Do List for Today
 Work
 In
independently and quietly!
your books:
 Read
& take notes on “Reaction Rates” on pg.
722-724
 Answer Question #1 on pg 731
 Read & take notes on 22-2 on pg. 732-737
 Answer Questions #1,2,5 on pg. 737
 Read & take notes on 22-3 on pg. 738-743
 Answer Questions #1-3
 DUE TOMORROW!!
DMA 4/14/11
+ Explain the collision theory.
+
To Do Today
 Lab
 Do
activity
the lab as written
 Then, rerun
the lab and time it
 Calculate

the rate of the 2 reactions
reactant/ time
 Due
at the end of Class today
 Yesterday’s
assignment is due tomorrow.
DMA 4/15/11
Name 2 reasons why a scientist
+ would want to change the rate of
a reaction. How could that
scientist accomplish this?
DMA 4/18/11
Chlorine gas reacts with solid
+ sodium to produce solid sodium
chloride. If 4.85 L of chlorine gas at
308 K and 1.05 atm react with
sodium, how many grams of sodium
chloride are produced?
DMA 4/19/11
+ Get out your pre-lab
Sit with your lab group
To Do Before you Start your lab:
1.Exchange your pre-lab with one person.
2.Grade the other person’s lab according to
+ the rubric.
3.Get your pre-lab back from who graded
yours and fix anything that is missing.
4.Bring your pre-labs to me in your lab
groups and I will approve you to begin the
lab.
DMA 4/20/11
If you were approved to begin
yesterday:
Get your lab equipment and try
to get a couple of trials done.
+
If you did not get your notebook
stamped yesterday:
Make sure your pre-lab is done
and get it to me for approval so
you can begin running your lab.
DMA 4/21/11
Define the following:
activation energy
+
catalyst
+
In the Lab today
 If
you were already approved to begin:
 Get
your lab equipment and get all 10 of your
trials done.
 Be sure to keep careful track of your data
 You need the change in products(gas) and
change in time to calculate the rate of the
reaction.
 If
you did not get your notebook stamped
yet:
 Make
sure your pre-lab is done and get it to me
for approval so you can begin running your lab.
DMA #1 4/22/11
What are the seven diatomic
Why is it important
+ molecules?
to know what these are?
+ Post Lab
Lab Due Tuesday

From your “Composition Book Data Collection Format”

Discussion of data collection

Graph-I want you to create a graph and attach it into your lab
notebook.
 The graph should have your manipulated variable on the xaxis and your responding variable (rate of reaction) on the yaxis
 This can be done by hand or computer

Calculations of rates of reaction should be present

Lab question: answer this as part of your conclusion (the
question asked at the beginning of the lab)

Conclusion: Claim, Evidence (data), Explain, Conclude
+
Rate of Reaction Experiments
You will design, conduct,
analyze, and report on how the
rate of a chemical reaction is
affected by the factors listed in
the next column
1.
Type of reactant
2.
Temperature
You are expected to design
your own experiments, not
copy those of other lab groups.
3.
Concentration of
reactant
You will need to collect 5 data
points for each experiment.
+
How can we measure the rate of
a reaction?

In math: rate = distance/time

In life: rate = anything changing/time

How can we measure the rate of the reaction:

Zn(s) + 2HCl(aq)  ZnCl2(aq) + H2(g)

Measure the zinc consumption or gas production versus time.

Draw & observe the standard set up for this lab.
+
Rubber tubing
One-hole stopper
Glass tube
10ml graduated cylinder
Test-tube
beaker
Standard Reaction Rate Set Up
+
Data & Observations
Keep Track of?
Mass of Metal Used for Each Trial (1.0g max)
Amount of Acid Used for Each Trial (10 ml max)
Temperature (50C max, don’t try to be exact)
Time Elapsed for Each Trial (5 minute max)
Volume of Gas Produced for Each Trial (10ml max)
Observations for Each Trial
+ Pre Lab Assignment:
 Question: “What
is the effect of _____ on the
reaction rate of hydrochloric acid and a particular
metal?”
 Hypothesis: (If, then, because)
 Manipulated Variable:
 Responding Variable:
 Controlled Variables:
 Materials:
 Diagram:
 Validity Measure: (Calibrate? Prevent
Contamination?)
 Procedure:
 (at least 5 levels, repeated at least 2x)
+
Experimental Controls?

Amount of Acid?

Strength of acid?

Amount of Metal?

When will you stop the trial?
+
+ Factors Affecting Rate
 Temperature
Increasing temperature always
increases the rate of a reaction.
 Surface Area
Increasing surface area increases the
rate of a reaction
 Concentration
Increasing concentration USUALLY
increases the rate of a reaction
 Presence of Catalysts
+
Factors Affecting Rate
What
are some real world examples
for
 temperature?
 surface Area?
concentration?
presence of Catalysts?
+
Catalysis
Catalyst: A
substance that speeds up a
reaction without being consumed
Enzyme:
A large molecule (usually a
protein) that catalyzes biological reactions.
Homogeneous
catalyst: Present in the same
phase as the reacting molecules.
Heterogeneous
catalyst: Present in a
different phase than the reacting molecules.
+
Catalysts Increase the Number of
Effective Collisions
+ Energy Diagrams
1.
Draw and label an energy diagram.
Calculate the activation energy and change
in energy of the reaction
• Reactants – 78.9 kJ
• Production – 125.3 kJ
• Activated complex – 300 kJ
•
•
•
Reactants - 25.1 kJ
Products – 35.2 kJ
Activated complex – 49.0 kJ
+
Collision Model
Collisions must have
enough energy to produce
the reaction (must equal or
exceed the activation
energy).
Orientation of reactants
must allow formation of
new bonds.
+
2NO2(g)  2NO(g) + O2(g)
Reaction Rates:
1. Can measure
disappearance of
reactants
2. Can measure
appearance of
products
3. Are proportional
stoichiometrically
+ Rate of Reaction Lab
Rate of Reaction Lab
Introduction

explain the collision theory and factors affect the
rate of the reaction

describe how you will investigate your condition

how the collision theory and rate of reaction relates
to the real world.

chemical reaction
Hypothesis
Variables
Procedure
+ DMA - Test Thursday
 What
is the collision theory?
 What
is an energy diagram?
 What
is an endothermic reaction?
 What
is an exothermic reaction?
 What
is activation energy?
 What
is a catalyst?
 How
do catalysis affect energy diagrams?
 What
factors affect the rate of the reaction?
+
Reversible Reactions
In some chemical reactions, it is possible for the
products to react and regenerate the reactants.
[Co(H2O)6]2+(aq) + 4Cl-(aq) 
pink
[CoCl4]2-(aq) + 6H2O(l)
blue
+
Use rate of reaction to explain how
you can tell if a chemical reaction has
reached chemical equilibria.
Equilibria
 Equilibrium
exists when
two opposing processes
occur at the same rate.
 Even
though changes
are occuring, there is no
overall, or net, change.
 Equilibria
exists when
the rate of forward
reaction is equal to the
rate of reverse reaction.
+
Le Châtelier’s Principle
 In
1888, a French mining
engineer named Henri Louis le
Châtelier discovered something
about what happens when a
reaction at equilibrium is
disturbed.
 Le
Châteleier’s Principle: if a
change in conditions is imposed
on a system at equilibrium, the
equilibrium position will shift in
the direction that tends to reduce
that change in conditions.
Le Châtelier’s Principle
If you introduce a change into a system at
equilibrium, the system will shift away from
the change to reestablish equilibrium.
What are two ways
to reestablish
equilibrium on this
teeter totter?
+
Le Châtelier’s Principle
Reactions that
reestablish
equilibrium shift either
to the right (make
more products) or to
the left (lose products,
form reactants).
Adding reactants shift
to the right.
Adding products shift
to the left.
+
Le Châtelier’s
Principle
Removing chemicals
also causes a reaction
to have to shift to
reestablish equilibrium.
Remove products,
shift right.
Remove reactants,
shift left.
Le
Châtelier’s
Principle
+
There are other ways to force an reaction to
reestablish equilibrium.
Increasing pressure on one side of a reaction
is analogous to increasing concentration of chemicals.
Decreasing pressure on one side of a reaction is
analogous to decreasing concentration of chemicals.
Exothermic/Endothermic
Reactions
+
 Exothermic
reactions
produce heat as a
product in addition to the
chemical products.
 Endothermic
reactions
absorb heat as a reactant
in addition to the
chemical reactants.
 Adding
or removing heat
from some reactions can
force the reaction to shift
and reestablish
equilibrium.
The Haber Process
 In
1913, the German scientist
Fritz Haber utilized Le
Châtelier’s Principle in a
process for making ammonia
from nitrogen and hydrogen.
 Ammonia
was used in the
manufacture of explosives.
 The
Haber process allowed the
Germans to continue
manufacturing explosives
despite a blockade on
necessary materials by
German opponents in WWI.
Haber
Einstein
Berlin 1914
Which way will equilibrium shift if the
highlighted chemical is increased.
1.
NO2(g)
N2O4(g)
2.
N2(g) + H2(g)
NH3(g)
3.
COCl2(g)
CO(g) + Cl2(g)
4.
H2(g) + I2(g)
HI(g)
5.
C(s) + H2O(g)
CO(g) + H2(g)
6.
ZnCO3(s)
ZnO(s) + CO2 (g)
7.
NH4Cl(s)
NH3(g) + HCl(g)
8.
SO2(g) + O2(g)
SO3(g)
9.
CO(g) + H2(g)
CH4(g) + H2O(g)
10. NaHCO3(s) + HCl(aq)
H2O(l) + CO2(g) +NaCl(aq)
Which way will equilibrium shift if
the highlighted chemical is
•NO
NO
decreased.
•N
+ H
NH
2(g)
2(g)
2
4(g)
2(g)
3(g)
•COCl2(g)
CO(g) + Cl2(g)
•H2(g) + I2(g)
HI(g)
•C(s) + H2O(g)
CO(g) + H2(g)
•ZnCO3(s)
ZnO(s) + CO2 (g)
•NH4Cl(s)
NH3(g) + HCl(g)
•SO2(g) + O2(g)
SO3(g)
•CO(g) + H2(g)
CH4(g) + H2O(g)
•NaHCO3(s) + HCl(aq)
H2O(l) + CO2(g) +NaCl(aq)
+
Chemistry DMA 5/7
List three things that you could do to a chemical reaction at
equilibrium to disrupt the system.
The Law of Chemical Equilibria
 In
1864, Norwegian
chemists Cato Maximillian
Guldberg and Peter
Waage formulated the Law
of Mass Action that
describes equilibrium.
 They
established an
Equilibrium Expression
and an Equilibrium
Constant,
Guldberg
Waage
+
aA
+ bB
cC + dD
Keq =
[C]c[D]d
[A]a[B]b
+
CO(g) + 3H2(g)
Keq =
[C]c[D]d
[A]a[B]b
Keq = [CH4][H2O]
[CO][H2]3
CH4(g) + H2O(g)
+
Homogenous and Heterogeneous
Equilibria

Homogenous equilibria occurs when all reactants and
products are in the same physical state.

Heterogeneous equilibria occurs when reactants and
products occurs in different physical states.
+
Heterogeneous Equilibria
NH4NO3(s)
Keq
= [N2O][H2O]2
N2O(g) + 2H2O(g)
+
Equilibrium Constant

Keq >> 1 at equilibrium, the system consists mostly of
products.

Keq = 1 at equilibrium, the system contains roughly equal
portions of products and reactants.

Keq << 1 at equilibrium, the system consists mostly of
reactants.
Write
equilibrium
expressions
for
each
+
of the following equations:
1.
NO2(g)
N2O4(g)
2.
N2(g) + H2(g)
3.
COCl2(g)
4.
H2(g) + I2(g)
HI(g)
5.
C(s) + H2O(g)
CO(g) + H2(g)
6.
ZnCO3(s)
ZnO(s) + CO2 (g)
7.
NH4Cl(s)
NH3(g) + HCl(g)
8.
SO2(g) + O2(g)
SO3(g)
9.
CO(g) + H2(g)
CH4(g) + H2O(g)
10.
NaHCO3(s) + HCl(aq)
NH3(g)
CO(g) + Cl2(g)
H2O(l) + CO2(g) +NaCl(aq)
DMA 4/25/11
+ Name 4 things that can
change the rate of a
reaction.
+
To Do Today
 Turn
in Ch 22 notes & questions
 Work
on 22-3 review (tomorrow’s quiz will be
similar to this)
 Create
 14-4
a graph- “Heating Curves”
read, take notes, answer questions #1-5
 Remember,
tomorrow
Rates of Reactions Lab is due
DMA 4/26/11
Turn in your homework:
+
1.Ch 22 notes & questions
2.Rates of Reactions Lab
3.22-3 Review
+ What are molecules doing in each
phase?—Sketch the containers and add
a description of the molecular motion
of each.
What happens to each when heated?
+
Phase Changes
Phases
at the molecular level
Change
 the
of phase-
conversion of a substance from one
state of matter into another
 Always involves a change in energy
+ Phase Changes
+
Phase Changes
Solid → Liquid →Gas
 Heat
of Fusion: Heat required to convert a solid at its
melting point to a liquid
 Heat
of Vaporization: Heat required to convert a liquid
at its boiling point to a gas.
 Heat
of Sublimation: Heat required to take a
substance directly from its solid to its gas state
5 Types of Energy
+
 Kinetic

The energy of motion
 Potential

Stored energy
 Chemical

energy inherent in the chemical bonds which hold molecules
together. Examples are coal and oil, which have energy potential
that is released upon combustion.
 Thermal/heat

the kinetic energy of molecular motion. Measured as temperature
and perceived as heat.
 Nuclear

The energy stored in the nucleus of an atom.
+
To Do
Complete
the “Phase Change” graphing
activity
2
graphs-water and ammonia
Do
not write on the assignment sheet-only
on your graph paper
Return
the data sheet at the end of class
+ Phase Change Diagram - Water
Note that the temperature of a substance does not change as it is going
through a phase change.
+ Calculating
Energy Changes - Heating
Curve for Water
140
120
DH = mol x DHfus
DH = mol x DHvap
Temperature (oC)
100
80
Heat = mass x Dt x Cp, gas
60
40
20
0
Heat = mass x Dt x Cp, liquid
-20
-40
-60
-80
Heat = mass x Dt x Cp, solid
-100
Time
+
Definitions
 Temperature
is a measure of the average
kinetic energy of the particles in a sample
of matter.
A
joule is the SI unit of heat as well as all
other forms of energy.
 Heat
can be thought of as the energy
transferred between sample of matter
because of a difference in their
temperatures.
Units
for
Measuring
Heat
+
The Joule is the SI system unit for
measuring heat:
kg  m
1 J  N m 
2
s
2
The calorie is the heat required to raise the
temperature of 1 gram of water by 1 Celsius
degree
1 calorie  4.18 Joules
+ Latent Heat of Phase Change
Molar Heat of Fusion
The energy that must be absorbed in order
to convert one mole of solid to liquid at its
melting point.
Molar Heat of Solidification
The energy that must be removed in order
to convert one mole of liquid to solid at its
freezing point.
+ Latent Heat of Phase Change #2
Molar Heat of Vaporization
The energy that must be absorbed in order
to convert one mole of liquid to gas at its
boiling point.
Molar Heat of Condensation
The energy that must be removed in order
to convert one mole of gas to liquid at its
condensation point.
+ Latent Heat – Sample Problem
Problem: The molar heat of fusion of water is
6.009 kJ/mol. How much energy is needed to convert 60 grams of ice
at 0C to liquid water at 0C?
60 g H 2O 1 mol H 2O 6.009 kJ
 20.00 kiloJoules
18.02 g H 2O 1 mol
Mass
of ice
Molar
Mass of
water
Heat
of
fusion
+
Calculating Energy Changes Heating Curve for Water
140
120
DH = mol x DHfus
DH = mol x DHvap
Temperature (oC)
100
80
Heat = mass x Dt x Cp, gas
60
40
20
0
Heat = mass x Dt x Cp, liquid
-20
-40
-60
-80
Heat = mass x Dt x Cp, solid
-100
Time
+
Specific Heat
The amount of heat required to
raise the temperature of one
gram of substance by one
degree Celsius.
+ Calculations involving Specific Heat
cp = ___q__
m x DT
OR
q = cp x m x DT
cp = Specific Heat
q = Heat lost or gained
DT = Temperature change
+

Specific heat of ice = 2.092 J/g°C

Specific heat of water = 4.184 J/g°C

Specific heat of steam = 2.013 J/g°C

Heat of vaporization = 2260 J/g

Heat of fusion = 334 J/g
+
Practice Problems

Calculate the energy released if 10 grams of steam at 110℃ is
cooled to -20 ℃.
+
Table of Specific Heats
+
A calorimeter is a
instrument that
measures the energy
absorbed or released
as heat in a chemical
or physical reaction.
+
Specific Heat
The amount of heat required to
raise the temperature of one
gram of substance by one
degree Celsius.
+ Calculations involving Specific Heat
cp = ___q__
m x DT
OR
q = cp x m x DT
cp = Specific Heat
q = Heat lost or gained
DT = Temperature change
+
Table of Specific Heats