Chemical EquilibriumChapter 16

Pre-requisitesNaming: Ionic, molecular and acids.

Balancing Reactions

Basic Stoichiometry

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 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

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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.

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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

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?

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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

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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 KpRemember

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?

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.

Only gas and aqueous species are included in the equilibrium constant expression: Why?

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 expression

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

Units on K

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.

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.

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

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 less than K the reaction shifts left

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.

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.

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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 less moles of gas

4 moles 1 moleShifts right

1 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 decreased, while SO3 is increased.

**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

Expanding the container at constant 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 KDeriving the Van’t hoff equation.

We’ll do this on the document camera

Van’t Hoff Equation: Graphs

Experimentally you can use this to determine the DH of reaction.

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.