Transcript 1Equilibrium

Welcome to 1C.
To do list week 1
Read the syllabus
If desired download powerpoint slides.
Register for MasteringChemistry
Complete extra credit chapter 16 pre- assignment
Complete intro to mastering assignment (due next week, but just get it done
now).
Places to get help
Office hours:
See course calendar.
Facebook. (https://www.facebook.com/groups/21521053832/)
Survey Tool (weekly answer summaries will be provided)
In Class: (TAs will be monitoring and answering during class)
Online eee chatroom.
Road map for doing your best with the least amount of
studying
Before class.
Skim lecture slides.
Skim book material on this topic. Special attention to figures.
Try to remember/connect it to things you have seen before.
Should take about 10 minutes per class period.
InClass
Stay engaged, take good notes, use chat function to ask questions when
needed.
After class.
After class. Do homework WITHOUT solutions. Think of problems as
puzzles with lots of alternate versions, NOT as many individual things to
memorize.
Procrastination:
“But cramming makes me do better on tests, I don’t care if it kills
me, I want an A”
No it doesn’t!
Howell, A. J., Watson, D. C., Powell, R. A., Buro, K. (2006) Academic procrastination: The
pattern and correlates of behavioural postponement
Tice, D. M., & Baumeister, R. F. (1997) Longitudinal Study of Procrastination, Performance,
Stress, and Health: The Costs and Benefits of Dawdling [2]
So how do I stop procrastinating?
Use easy to handle time gaps.
Immediately after class: (preferably within 10 minutes, but at absolute latest, before
sleep.
2+ questions you have (bonus: post these on the survey tool online)
3+ sentence summary of class from notes
This should not take more than about 10 minutes. The sooner you do it, the faster it will
go.
Realistic goals for increased studying.
I will study, do problems, review chemistry at least ____ minutes per day.
Suggested 20 minutes. No interruptions (turn off your
phone/messager/email/facebook, lock your doors).
Alternative: I will do ____ number of problems per day.
Suggested, do mastering problems related to each class immediately afterwards. On
“off” days pick 2 book problems.
Chemical Equilibrium
Pre-requisites
Naming: Ionic, molecular and acids.
Balancing Reactions
Basic Stoichiometry
Gas Law Materials
Places you’ll see this material again.
Next three chapters
All molecular biology and biochemistry classes
Anatomy & Physiology Classes
Pharmacy and medical school classes.
Learning Outcomes:
Express equilibrium constants for chemical equations.
Manipulate equilibrium constants to reflect changes in
the chemical equation.
Relate Kp and Kc
Write equilibrium expressions for reactions involving a
solid or liquid.
Finding equilibrium concentrations from
initial concentrations and equilibrium
constant.
Calculating equilibrium partial pressures from
the equilibrium constant and initial partial
pressures
Finding equilibrium constants from experimental data.
Finding equilibrium concentrations from
initial concentrations in cases with a small
equilibrium constant
Predicting the direction of a reaction by comparing Q
and K
Determine the effect of a concentration
change on equilibrium.
Calculating equilibrium concentrations from the
equilibrium constant and one or more equilibrium
concentrations.
Determining the effect of a concentration
change on equilibrium.
Finding equilibrium concentrations from initial
concentrations and equilibrium constant. Calculating
equilibrium concentrations from the equilibrium
constant and one or more equilibrium concentrations.
Determining the effect of a volume change on
equilibrium.
Determining the effect of a temperature
change on equilibrium.
Dynamic Equilibrium
Learning Outcomes:
Define Dynamic equilibrium.
On a time course graph, identify when equilibrium is
reached.
Throwing Contest Analogy
The Avengers face off against me and my friends:
Rules of the game:
My team starts with 10,000 balls. The Avengers start with none.
Each team must throw balls to the other side as quickly as possible
Initial Count: 10,000
Initial Count: 0
Equilibrium is reached when the number of balls
on each side isn’t changing
Throwing Contest Analogy
Initial Count: 10,000
Final Count: 9997
Initial Count: 0
Final Count: 3
Important similarities to chemical equilibrium:
Even though the number of balls on each side is staying the same, they are
still being exchanged.
At equilibrium the speed of the forward and the reverse reaction are the same.
What is dynamic equilibrium?
A reaction (or phase change)
doesn’t simply go in one direction.
Physical
Chemical
Chemical
Example
© Pearson
Review
Dynamic equilibrium occurs when the rate of the forward process
equals the rate of the reverse process.
When dynamic equilibrium is reached their will be no change in
amount of products or reactants, even though the forward and
reverse reactions are still occurring.
© Pearson
Equilibrium Constant in terms of
concentration (Kc)
Learning Outcomes:
Write equilibrium constant (Kc) for a given reaction.
Use data to determine the equilibrium constant (Kc) for
a given reaction.
Determine the extent of the forward and reverse
reactions (are products or reactants favored) when given
the Kc
Chemical Example
Equilibrium
Equilibrium
Equilibrium
© McGraw Hill
Starting
Conditions
No N2O4
All NO2
No NO2
All N2O4
Some of Each
Equilibrium Constant (Kc)
Mathematical relationship that relates reactants and products.
products
reactants
concentrations
Writing Kc Example:
Write the Kc equation for the following reaction.
Kc Examples:
Write the expression for the equilibrium constant for the following
reactions:
Pure phosgene gas (COCl2), 3.00x10-2 mol, was placed in a 1.50L
container. It was heated to 800K and the pressure of CO was found to
be 0.497 atm. Calculate the equilibrium constant Kp for the reaction.
Why Kc?
© Pearson
Questions:
If Kc is significantly greater than 1 what does that
mean about the concentrations of products to
reactants?
If Kc is significantly less than 1 what does that mean
about the concentrations of products to reactants?
What does the size of Kc say about the
reaction?
Products
Favored
K>>1
© Pearson
Reactants
Favored
K<<1
© Pearson
Review
Kc is equal to the products over the reactants raised to the power of
the coefficient
Kc is used because the final ratios of products to reactants change
based on initial concentraions and stoichiometric coefficents but Kc
does not.
If Kc is much greater than one then products are favored
If Kc is much less than one, then reactants are favored.
The last two will be true of all the equilibrium constants we’ll cover
over the next several chapters.
Kp Equilibrium Constant in Terms of
Pressure
Learning Outcomes:
Write equilibrium constant in terms of pressure (Kp)
for a given reaction.
Use data to determine the equilibrium constant (Kp)
for a given reaction.
Determine the relation between Kc and Kp
Convert between Kc and Kp
Equilibrium Constant (Kp)
pressures
Kp Example
Write the Kc equation for the following reaction.
Relation between Kc and Kp
Why?
Relation between Kc and Kp
Remember
from 1A/1B
Rearranging
Rearranging
And again:
mol/L
Kp Examples:
Find Kp for the decomposition of phosphorus pentachloride into
phosphorus trichloride and chlorine gas given that the
equilibrium partial pressures are 0.875atm, 0.463atm and 1.98
atm respectively at 250oC. (review naming if needed)
Then find its Kc.
When is Kc equal to Kp?
Review Slide:
The magnitude of K tells us whether reactants or products are favored
K>>1 Products favored
K<< 1 Reactants are favored
Only gas and aqueous species are included in the equilibrium
constant expressions
Concentrations at equilibrium vary depending on initial conditions
This is why we use K rather than ratios.
Kc does not change given a constant temperature
Announcements
To do list week 1
Read the syllabus
If desired download powerpoint slides.
Register for MasteringChemistry
Complete extra credit chapter 16 pre- assignment
Complete intro to mastering assignment (due next week, but just get it done now).
Remember
Chat room is going and being monitored by Allison
Weekly Surveys are available.
Two quick clarifications, review
Question:
Only gas and aqueous species are included in the equilibrium constant
expression: Why?
Orange= CaCO3
Green= CaO
Red+Black=CO2
The amount of CaO and CaCO3 doesn’t matter, so
long as there is enough of each that there is leftover
at equilibrium. Concentration is effectively constant
although the mass changes.
Quick note about K units
Equilibrium constants are unitless
This is because we aren’t really using the concentrations and
pressures, but are actually using “activities”
We are actually filling in
Where Po is 1atm and co is 1 mol/L leaving the activity unitless.
Review
Only gas and aqueous species are included in equilibrium constant
expressions
K does not have units. All concentrations and pressures are filled into K as
“activities”.
Medical and Pharmacological
applications.
Learning Outcomes:
Discuss a medical/pharmalogical term that is related to our
topics.
Use Kd definition to compare binding efficiencies and doses
of two theoretical drugs (don’t worry, no bio required).
Related Medical/Pharmalogical Term:
Kc is related to the amount of drug needed to give a specific amount of
the drug/protein complex.
Use to decide dosage of a drug (also used in other protein binding medical applications.)
Kd Example
Drug A has a much higher Kd than Drug B. Both work on the same receptor, and both
have similar cellular responses. Answer the following questions. (Note: no biology background
is required to do the problem.)
Which binds more tightly (aka has a
high affinity).
Drug B
If binding is required for the drug
effect, which would you expect to be
more effective at lower doses.
Drug B.
Review
In medical applications Kd is often used instead of Kc.
The lower the Kd the higher the binding affinity.
Manipulating Chemical Equations and
Effects on Kc and Kp
Learning Outcomes:
Determine Kc when a reaction is reversed.
Determine Kc when a reaction is multiplied by a number.
Manipulating K Example
Write the equation for the Kc of the reverse reaction of the
previous example. What is the relation between them?
Previous example:
Reverse Reaction:
They are inverses of each other!
Manipulating K Example
If you multiply the decomposition of N2O4 equation by 2,
what is the Kc? What is the relation between the Kc of each?
Previous example:
Multiplied by 2 Reaction:
If you multiply by 2 you square it!
Review
If you reverse the equation. K is inverted.
If you multiple the equation by a number K is raised to
that power.
Motivational Moment of the Week
Combining Equations to Find New Kcs
Learning Outcomes:
Reverse and multiply equations and then add them to
create a new chemical equation (this is similar to Hess’s Law Problems).
Use the rules for Kc to determine the Kc of the new
equation (different rules than Hess’s Law)
Combine multiple equations to determine a Kc for a new
reaction. (different rules than Hess’s Law)
Combining multiple Kc
If two equations add to a new equation. The Kc of the new equation
is found by multiplying the component equations.
***Similar to the Hess’s law problems, but be careful its multiplication NOT addition.
Kc= (K’c)(K’’c)
How can we solve for Kc here?
Need to invert
this equation.
How do you think we solve for Kc here?
Need to invert
this equation.
Examples
Given each of the following equilibrium constants, find the
unknown equilibrium constant.
Example 1
Examples
Given each of the following equilibrium constants, find the
unknown equilibrium constant.
Example 2
Review
If you reverse a reaction, take the inverse of the K
If you multiply a reaction by a number, raise the K to that exponent.
If you add reactions to get a new reaction, multiply the Ks.
Reaction Quotient.
Learning Outcomes:
Define the Reaction Quotient (Q)
Identify the difference between Q and K
Determine the direction that a reaction will proceed given
the Q and K.
Reaction Quotient (Q)
Defined the same as K, (products  reactants, raised to
stoichiometric coefficients)
K is at equlibrium
Q can be determined at any point in the reaction.
Compare Q and K to see if the reaction will go “forward” “reverse”
or if it is already at equilibrium
Q<K Reaction moves forward, aka from left to right
Q>K Reaction moves in reverse, aka right to left
Q=K the reaction is already at equilibrium and stays the same
Lets look at why this is.
Q compared to K
Q<K
Products are smaller/reactants are
bigger than K, so must shift to adjust.
Reaction moves forward, to create
more products, and less reactants
Q>K
Reactants are smaller/products are
bigger than K, so must shift to adjust
Reaction moves to left to create more
reactants and less products.
Q=K the reaction is already at
equilibrium and stays the same
Example: Using Q to predict reaction direction
For the synthesis of ammonia the Kc at 375oC is 1.2. The initial concentrations are
H2=0.76 M N2=0.60 M and NH3 = 0.48 M. Which way will the reaction shift? What
will happen to the concentration of each gas?
Find Q and compare it to K
Shifts forward,
or to the right.
Examples: Using Q to predict reaction direction
The Kp for the reaction below is 5.60x104 at 350oC the initial pressures are SO2=
0.350 and O2=0.762. Is the total pressure at the end, less, greater or the same as
intial?
No product. So the reaction MUST shift forward.
Reactants = 3 moles.
Products = 2 moles.
Total moles are lower on side of reaction it is shifting toward,
So the ideal gas law says that pressure goes down.
Review
The reaction quotient (Q) is products over reactants, raised to
their stoichiometric coefficents.
It differs from K, because K is the value at equilibrium while Q is
the value at any given point.
You can determine which way the reaction will shift by comparing
Q and K
If Q is less than K the reaction shifts right
If Q is greater than K the reaction shifts left
Deep Thought for the Week
Calculating Equilibrium Constants
Learning Outcomes
Calculate equilibrium concentrations from the initial
concentrations and Kc.
Calculate Kc from experimental equilibrium
concentrations.
Pre-requisite Requirements
Using Q to tell which way the reaction will shift.
Defining K, writing equation for K.
Calculating Equilibrium Concentrations
For you to use as guidelines, sometimes need to be altered based on the situation, follow
along as we do problems.
Step 1: Use initial concentrations to calculate Q
Step 2: Decide which way the reaction shifts
Step 3: Recommended make an “ICE” chart
Initial, Change, Equilibrium.
Step 4: Fill in initial concentrations and changes (in variable form
if need be).
Step 5: Fill in equilibrium concentrations
Step 6: Fill into K equation and solve for the variable.
Example:
Calculate the number of moles of H2 that are present at equilibrium if a mixture of
0.300 mol of CO and 0.300 mol of H2O is heated to 700oC (kc=0.534) in a 10.0 L
container.
Example:
At a certain temperature, the equilibrium constant, Kc, for this reaction is 53.3. At
this temperature, 0.300 mol of H2 and 0.300 mol of I2 were placed in a 1.00L
container to react. What concentration of HI is present at equilibrium?
Example:
For the decomposition of phosphorous pentachloride to phosphorous trichloride
and chlorine at 400.K the Kc is 1.1x10-2. Given that 1.0g of phosphorous
pentachloride is added to a 250mL reaction flask, find the final concentrations of
each species and the percent decomposition.
Example
Consider the reaction below. A reaction mixture at 780 oC initially contains [CO]=0.500
M and [H2]= 1.00M. At equilibrium, the CO concentration is 0.15 M. What is the value
of the equilibrium constant?
Review
Using ICE charts, the equation of Kc, and initial concentrations we
can determine the concentration of all compounds.
Using ICE charts, the equation of Kp, and initial pressures we can
determine the concentration of all compounds.
Remember to think of all of these problems as one type, rather than
memorizing each protocol separately.
announcements
Things to do this week:
If you haven’t already, visit the website, read the syllabus, sign up for mastering and
download the slides.
Acid/Base Equilibrium slides are posted if you would like to download them.
Assignments:
Intro to mastering Due Tuesday
Extra credit reading assignment due on Wednesday.
Start doing Chapter 16 part 1. You can complete this as of last Thursday’s lecture. I
moved out the deadline due to sorting out enrollment issues, but you should
complete this ASAP, as the next large assignment is due that Thursday.
Extra Examples:
Lets get some extra practice with problems involving ice charts!!!
Week 1 Review Examples:
16.42 (modified a bit) For the reaction 2A(g)⇌B(g)+2C(g), a reaction vessel initially contains
only A at a pressure of PA=0.296 atm . At equilibrium, PA=0.0724 atm.
Calculate the value of Kp. (Assume no changes in volume or temperature.)
16.45 Consider the following reaction:
H2(g)+I2(g)⇌2HI(g)
A reaction mixture in a 3.67 L flask at a certain temperature initially contains 0.763 g H2 and
96.9 g I2. At equilibrium, the flask contains 90.4 g HI.
Calculate the equilibrium constant (Kc) for the reaction at this temperature.
Motivational Moment of the Week
Le Châtelier’s Principle
Learning Outcomes.
Define Le Chatelier’s Principle.
Identify the effect adding or removing a reactant has on a reaction.
Identify the effect adding or removing a product has on a reaction.
Identify the effect changing volume or pressure has on a reaction.
Identify the effect changing temperature has on a reaction.
Le Chatelier’s Principle
If you apply stress to a system it shifts to relieve the stress.
Reaction shifts left or right to relieve the stress.
Analogy-
Concentration
Change in Concentration
Reaction shifts away from added species
Reaction shifts toward subtracted species
For the generic reaction above:
Adding A shifts the reaction right.
Subtracting A shifts the reaction left.
Adding B shifts the reaction right.
Subtracting B shifts the reaction left.
Adding C shifts the reaction left.
Subtracting C shifts the reaction right
Adding D shifts the reaction left.
Subtracting D shifts the reaction right
Example:
If you add N2 which way does the reaction shift? What
happens to each of the species?
Reaction shifts to the right.
H2 decreases
NH3 increases
Example:
Calculating changed concentrations
In the laboratory studying the extraction of iron metal from iron ore, the following
reaction was carried out at 1270K in a reaction vessel of volume 10.0L. At equilibrium
the partial pressure of CO was 4.24bar and that of CO2 was 1.71 bar. The pressure of
the CO2 was reduced to 0.62 bar by reacting some of it with NaOH and the system was
allowed to reach equilibrium again. What will be the partial pressure of each gas once
equilibrium is re-established?
Graphical Representation.
© Pearson
Volume and Pressure (in gases)
Increase volume, and therefore decrease in pressure: shifts toward the
side with more moles of gas
Decrease volume, and therefore increase in pressure it shifts toward
the side with less moles of gas
4 mols
reactant
Low pressure
Shifted further to the left
2mol
product
High pressure
Shifted further to the right
Example:
If the volume of a sample containing the equilibrium below is decreased, what
will happen to the concentration of each species?
Less moles
More moles
Volume decreased means
pressure is increased, shifts to
side with less moles.
Shifts left.
N2O4 increases, NO2
decreases.
Examples:
Predict the direction in which each of the following equilibrium will
shift if the pressure on the system is decreased by expansion.
Decreasing pressure means it shifts to the side with
more moles of gas
1 mole
4 moles
Shifts left
1 mole
2 moles
1 mole
1 mole
Shifts right
Stays the same.
Temperature Change
Alters Kc/Kp (unlike all other changes)
For an endothermic reaction (DH= positive)
Think of heat as a reactant
raising the temperature shifts the reaction to the products
Lowering the temperature shifts the reaction to the reactants
For an exothermic reaction (DH= negative)
Think of heat as a product
raising the temperature shifts the reaction to the reactants
Lowering the temperature shifts the reaction to the products
Example
If you raise the temperature of the reaction below, what happens to the concentration
of each species?**
Exothermic
-DH means heat is produced so think of it as a “product”,
Increased temperature increases a “product”, so it must shift…..?
left
This means SO2, and O2 are increased, while SO3 is decreased
**Note: unless stated otherwise we assume that it stays at a constant pressure. Otherwise that would
have an effect as well.
Change in Concentration
Review:
Reaction shifts away from added species
Reaction shifts toward subtracted species
Change in Volume
Increase volume it shifts toward the side with more moles of gas
Decrease volume it shifts toward the side with less moles of gas
Change in Pressure
Increase in pressure shifts toward the side with less moles of gas
Decrease in pressure shifts toward the side with more moles of gas
Change in Temperature
Only way to alter the equilibrium constant
For an endothermic reaction
raising the temperature shifts the reaction to the products
Lowering the temperature shifts the reaction to the reactants
For an exothermic reaction
raising the temperature shifts the reaction to the reactants
Lowering the temperature shifts the reaction to the products
Review Example
How will the amount of Ammonia be affected by the following
Removing O2?
Increasing pressure?
Shifts left.
Increases ammonia
Adding N2?
Shifts left
Increases ammonia
Adding water?
Shifts left
Increases ammonia
Shifts to less moles of gas
Shifts right
Decreases ammonia
Increasing the temperature?
Exothermic, heat is a product
Adding a product shifts left
Ammonia increases
Free Energy and Equilibrium
Learning Goals
Identify relationship between Gibbs Free Energy and equilibrium constant.
Use equation to convert between Gibbs Free Energy and equilibrium constant.
Derive Van‘t Hoff Equation.
Use Van’t Hoff Equation to relate two pairs of K and T.
Determine if a reaction is endothermic or exothermic based on the graph of LnK
vs 1/T
Determine DH by graph of LnK vs 1/T
Determine DH by using Van’t Hoff equation
Free Energy and Equilibrium
Equation to relate K or Q with DG.
DG: gibbs free energy (same as from 1B)
R: related to energy, so use 8.31 J/mol*K
T: kelvin
Example
Find Kc at 273K, using the values below.
Species
N2*
H2*
NH3
DGfo
0 kJ/mol
0 kJ/mol
-16.4 kJ/mol
Note: remember anything in its standard state is zero, I wouldn’t HAVE to give you this on an exam.
Relating temperature and K
Van’t Hoff Equation: Graphs
Experimentally you can use this to determine the DH of reaction.
Relating temperature and K
Deriving the Van’t hoff equation.
We’ll do this on the document camera
Example
Use the graph to answer the following:
Is the reaction endothermic or
exothermic?
y= -2.144x105x+2559
Endothermic: 1/T decrease
as ln K Decreases
So K decreases as T increases=enothermic
What is DH
Review
K changes as temperature changes.
Q and K can be related to thermodynamic properties though the
following three equations.
This can be used to determine the DH if we know K at two different
temperatures.