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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 helpOffice 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.
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.
InClassStay 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.
Road map for doing your best with the least amount of studying
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-requisitesNaming: Ionic, molecular and acids.
Balancing Reactions
Basic Stoichiometry
Gas Law MaterialsPlaces 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 constants from experimental data.
Predicting the direction of a reaction by comparing Q and K
Calculating equilibrium concentrations from the equilibrium constant and one or more equilibrium concentrations.
Finding equilibrium concentrations from initial concentrations and equilibrium constant. Calculating equilibrium concentrations from the equilibrium constant and one or more equilibrium concentrations.
Finding equilibrium concentrations from initial concentrations and equilibrium constant.
Calculating equilibrium partial pressures from the equilibrium constant and initial partial pressures
Finding equilibrium concentrations from initial concentrations in cases with a small equilibrium constant
Determine the effect of a concentration change on equilibrium.
Determining the effect of a concentration change on equilibrium.
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 AnalogyThe 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
Equilibrium is reached when the number of balls on each side isn’t
changing
Initial Count: 0Initial Count: 10,000
Initial Count: 0Final Count: 3
Initial Count: 10,000Final Count: 9997
Important similarities to chemical equilibrium:
Even though the number of balls on each side is staying the same, they are still being exchanged.
Throwing Contest Analogy
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
Starting Conditions
No N2O4
All NO2
No NO2
All N2O4
Some of Each
Equilibrium
Equilibrium
Equilibrium
© McGraw Hill
Equilibrium Constant (Kc)
Mathematical relationship that relates reactants and products.
concentrations
products
reactants
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
© Pearson
© Pearson
Reactants Favored
K>>1
K<<1
ReviewKc 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
And again:
mol/L
Rearranging
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
AnnouncementsTo do list week 1
Chat room is going and being monitored by Allison
Weekly Surveys are available.
Remember
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).
Two quick clarifications, review
Question:
Orange= CaCO3
Green= CaORed+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.
Only gas and aqueous species are included in the equilibrium constant expression: Why?
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
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.
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.)
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 ExampleWrite 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 ExampleIf 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!
ReviewIf 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.
ExamplesGiven each of the following equilibrium constants, find the unknown equilibrium constant.
Example 1
ExamplesGiven 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 RequirementsUsing Q to tell which way the reaction will shift.
Defining K, writing equation for K.
Calculating Equilibrium Concentrations
Step 1: Use initial concentrations to calculate Q
Step 2: Decide which way the reaction shifts
Step 3: Recommended make an “ICE” chartInitial, 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.
For you to use as guidelines, sometimes need to be altered based on the situation, follow along as we do problems.
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.
ExampleConsider 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.
announcementsThings 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 PrincipleIf you apply stress to a system it shifts to relieve the stress.
Reaction shifts left or right to relieve the stress.
Analogy-
ConcentrationChange 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
concentrationsIn 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
Low pressureShifted further to the left
High pressureShifted further to the right
4 mols reactant
2mol product
Example: If the volume of a sample containing the equilibrium below is decreased, what will happen to the concentration of each species?
More molesLess 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
4 moles 1 mole
Shifts right1 mole 2 moles
Shifts left
1 mole1 moleStays the same.
Temperature ChangeAlters 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
Exothermic
-DH means heat is produced so think of it as a “product”, Increased temperature increases a “product”, so it must shift…..?
left
If you raise the temperature of the reaction below, what happens to the concentration of each species?**
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.
Review:Change in Concentration
Reaction shifts away from added species
Reaction shifts toward subtracted species
Change in VolumeIncrease 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 PressureIncrease 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 TemperatureOnly 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?Shifts left.
Increases ammonia
Adding N2?Shifts left
Increases ammonia
Adding water?Shifts left
Increases ammonia
Increasing pressure?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
DG: gibbs free energy (same as from 1B)
R: related to energy, so use 8.31 J/mol*K
T: kelvin
Equation to relate K or Q with DG.
ExampleFind Kc at 273K, using the values below.
Species
DGfo
N2* 0 kJ/mol
H2* 0 kJ/mol
NH3-16.4 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 KDeriving the Van’t hoff equation.
We’ll do this on the document camera
ExampleUse the graph to answer the following:
Is the reaction endothermic or exothermic?
What is DH
Endothermic: 1/T decrease as ln K Decreases
So K decreases as T increases=enothermic
y= -2.144x105x+2559
ReviewK 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.