Etiket Arşivleri: Chemical Equilibrium

Chemical Equilibrium Notes

Chemical Equilibrium

Chapter 15

Introduction

Lets consider the catalytic methanation reaction

The Equilibrium Constant

Some equilibrium compositions for the methanation reaction

Example 

The Law of Mass Action

The manipulation rules of equilibrium constants

Heterogeneous Equilibria

Calculating Equilibrium Constants

Example

USING THE EQUILIBRIUM CONSTANT

we described how a chemical reaction reaches equilibrium.

how this equilibrium can be characterised by the equilibrium constant.

Applications of Equilibrium Constants

Predicting the direction of Reaction.

Calculating Equilibrium Concentrations

If Qc > Kc, the reaction will go left

If Qc < Kc, the reaction will go right

If Qc = Kc, the reaction is at equilibrium

Example

Three steps in solving equilibrium concentrations:

Set up a table of concentrations.

Substitute the expressions in x for equilibrium concentrations into the equilibrium constant expression.

Solve the equilibrium constant expression for the values of the equilibrium concentrations.

Changing the Reaction Conditions:

Le Chatelier’s Principle

Change in Reactant or Product Concentrations

Effects of Volume and Pressure Changes

Example

The Effect of Catalyst


Experiment 5: Chemical Equilibrium ( Kinetics measured by Pressure )

Chemistry 102 Laboratory

Experiment 5:  Chemical Equilibrium (Kinetics measured by Pressure)

Laboratory

In this experiment you will measure the rate of decomposition of carbonic acid by measuring the rate of production of carbon dioxide in the following reaction:

H2CO3  <–>  CO2(g)  +  H2O

The equilibrium constant for this reaction, in terms of concentration, is given by the expression:

Keq  =  [CO2]/ [H2CO3]

In this experiment, however, the pressure of carbon dioxide, rather than its concentration, is to be measured.  In particular, the pressure is measured as a function of time, and at two different temperatures.  Since the equilibrium and chemical rates are both functions of temperature, two different rates will be measured at the two different temperatures.  Since this decomposition reaction is known to be first order in the concentration of the carbonic acid, the rate varies as the exponent of the absolute temperature of the solution.  For every 10o increase in temperature, the rate of the reaction will approximately increase by a factor of 2.  Equilibrium is attained when the pressure remains a constant with time.  In this experiment, you will make use of the ideal gas equation and the fact that the equilibrium constant may be stated in terms of the gas pressure of carbon dioxide as well as the concentration (mass) of carbon dioxide.

Instructions

Computer Setup

1. Obtain a Pressure Sensor and plug it into an Analog channel of the interface box.  Make sure the interface box is on, then start Science Workshop.  Use the analog plug button to select your channel and then select the Absolute pressure sensor.  Double click on the pressure sensor icon, and allow the sensor to warm up for 5 minutes.

2. Using the barometer in the Main Lab, obtain the current atmospheric pressure in mm Hg.  Convert to atm.  In the screen for senor calibration, change kPa to atm for the pressure units.  In this same screen change the low pressure known value from 300 to the atmospheric pressure you measured in “1” above.  After the current reading of the pressure stabilizes, click on read to calibrate the pressure.  You are done with this part.

3. Check to make sure the sampling rate is set for Fast at 10 Hz by looking at the Sampling Options.  Bring up a graph window by dragging its icon over that of the pressure sensor.

Getting Set Up

4. Insert the barb of a quick release connector into one end of the plastic tubing that comes with the pressure sensor.

5. Obtain a one hole rubber stopper that fits a 250 mL glass bottle or flask.

6. Place a drop of glycerin on the bottom end of the hole of the stopper.

7. Insert the glass part of an eyedropper, tip up, through the hole in the stopper; the tip must clear the top of the stopper.

8. Carefully fit the glass tip into the open end of the plastic tubing.

9. Align the quick-release connector with the connector on the PRESSURE PORT of the Pressure Sensor.

10. Push the connector onto the port and then turn the connector clockwise until it clicks.

11. DO NOT INSERT THE STOPPER IN THE BOTTLE YET!!!  The pressure sensor should still be open to the atmosphere at this point.

Recording Data

12. Obtain 100 mL of room temperature (a posted value) soda water and place it in the 250 mL flask.

13. When everything is ready, click REC to begin recording data.

14. QUICKLY insert the stopper in the bottle; make sure there is a tight seal. DO NOT shake or stir the contents of the bottle.

15. Observe the changes in pressure as the carbonic acid in the soda water decomposes in the bottle.

16. Record for 6 minutes.

17. Slowly remove the stopper from the bottle and allow the pressure readings to stabilize.  This is to determine whether or not the air pressure or calibration have changed.

18. Stop the recording.  Determine the initial rate of decomposition of carbonic acid in the same manner employed in experiment 4 (i.e. a linear fit of the data).  Also record the final pressure of the gas in the flask.

19. Obtain 100 mL of cold soda water and place it into a 250 mL flask which has been cooled on ice in an ice bath.  Leave the water and the flask in the ice bath and return to your lab station.Repeat the procedure with cold soda water.

Chemical Equilibrium

Equilibrium is a state in which there are no observable changes as time goes by.

  Chemical equilibrium is achieved when:

  1.) the rates of the forward and reverse reactions are equal and

  2.) the concentrations of the reactants and products remain constant

  • Equilibrium

  • There are two types of equilibrium: Physical and Chemical.

Physical Equilibrium

  • H20 (l) ↔ H20 (g)

–Chemical Equilibrium

  • N2O4 (g) ↔ 2NO2

  • Physical Equilibrium

  • Chemical Equilibrium

  • Chemical Equilibrium

  • Law of Mass Action

  • Law of Mass Action- For a reversible reaction at equilibrium and constant temperature, a certain ratio of reactant and product concentrations has a constant value (K).

  • The Equilibrium Constant (K)- A number equal to the ratio of the equilibrium concentrations of products to the equilibrium concentrations of reactants each raised to the power of its stoichiometric coefficient.

  • Law of Mass Action

  • For the general reaction:

  • Equilibrium Constant

N2O4 (g) ↔ 2NO2 (g)

  • Chemical Equilibrium

  • Chemical equilibrium is defined by K.

  • The magnitude of K will tell us if the equilibrium reaction favors the reactants or the products.

  • If K » 1……..favors products

  • If K « 1……..favors reactants

  • Equilibrium Constant Expressions

  • Equilibrium constants can be expressed using Kc or Kp.

  • Kc uses the concentration of reactants and products to calculate the eq. constant.

  • Kp uses the pressure of the gaseous reactants and products to calculate the eq. constant.

  • Equilibrium Constant Expressions

  • Equilibrium Constant Equations

  • Homogeneous Equilibrium

  • Homogeneous Equilibrium- applies to reactions in which all reacting species are in the same phase.

               N2O4 (g) ↔ 2NO2 (g)

  • Equilibrium Constant Expressions

  • Relationship between Kc and Kp

  • Equilibrium Constant Calculations

The equilibrium concentrations for the reaction between carbon monoxide and molecular chlorine to form COCl2 (g) at 740C are [CO] = 0.012 M, [Cl2] = 0.054 M, and [COCl2] = 0.14 M.  Calculate the equilibrium constants Kc and Kp.

  • Equilibrium Constant Calculations

  • The equilibrium constant Kp for the reaction is 158 at 1000K. What is the equilibrium pressure of O2 if the PNO = 0.400 atm and PNO = 0.270 atm?

  • Heterogeneous Equilibrium

  • Heterogeneous Equilibrium- results from a reversible reaction involving reactants and products that are in different phases.

  • Can include liquids, gases and solids as either reactants or products.

  • Equilibrium expression is the same as that for a homogeneous equilibrium.

  • Omit pure liquids and solids from the equilibrium constant expressions.

  • Heterogeneous Equilibrium Constant

  • Heterogeneous Equilibrium Constant

  • Equilibrium Constant Calculations

Consider the following equilibrium at 295 K:

The partial pressure of each gas is 0.265 atm. Calculate Kp and Kc for the reaction.

  • Multiple Equilibria

  • Multiple Equilibria- Product molecules of one equilibrium constant are involved in a second equilibrium process.

  • What does the Equilibrium Constant tell us?

  • We can:

–Predict the direction in which a reaction mixture will proceed to reach equilibrium

–Calculate the concentration of reactants and products once equilibrium has been reached

  • Predicting the Direction of a Reaction

  • The Kc for hydrogen iodide in the following equation is 53.4 at 430ºC. Suppose we add 0.243 mol H2, 0.146 mol I2 and 1.98 mol HI to a 1.00L container at 430ºC. Will there be a net reaction to form more H2 and I2 or HI?

                   H2 (g) + I2 (g) → 2HI (g)

  • Reaction Quotient

The reaction quotient (Qc) is calculated by substituting the initial concentrations of the reactants and products into the equilibrium constant (Kc) expression.

IF

  • Qc > Kc system proceeds from right to left to reach equilibrium

  • Qc = Kc the system is at equilibrium

  • Qc < Kc system proceeds from left to right to reach equilibrium

  • Reaction Quotient

  • Calculating Equilibrium Concentrations

  • If we know the equilibrium constant for a reaction and the initial concentrations, we can calculate the reactant concentrations at equilibrium.

  • ICE method

  Reactants  Products

  Initial (M):

  Change (M):

  Equilibrium (M):

  • Calculating Equilibrium Concentrations

  • At 1280ºC the equilibrium constant (Kc) for the reaction is 1.1 x 10-3. If the initial concentrations are [Br2] = 0.063 M and [Br] = 0.012 M, calculate the concentrations of these species at equilibrium.

  • Calculating Equilibrium Concentrations

  • Express the equilibrium concentrations of all species in terms of the initial concentrations and a single unknown x, which represents the change in concentration.

  • Write the equilibrium constant expression in terms of the equilibrium concentrations. Knowing the value of the equilibrium constant, solve for x.

  • Having solved for x, calculate the equilibrium concentrations of all species.

  • Factors that Affect Chemical Equilibrium

  • Chemical Equilibrium represents a balance between forward and reverse reactions.

  • Changes in the following will alter the direction of a reaction:

–Concentration

–Pressure

–Volume

–Temperature

  • Le Châtlier’s Principle

  • Le Châtlier’s Principle- if an external stress is applied to a system at equilibrium, the system adjusts in such a way that the stress is partially offset as the system reaches a new equilibrium position.

  • Stress???

  • Changes in Concentration

  • Increase in concentration of reactants causes the equilibrium to shift to the ________.

  • Increase in concentration of products causes the equilibrium to shift to the ________.

  • Changes in Concentration

  Change  Shift in Equilibrium

Increase in [Products]  left

Decrease in [Products]  right

Increase in [Reactants]  right

Decrease in [Reactants]  left

  • Changes in Concentration

FeSCN2+(aq) ↔ Fe3+(aq) + SCN-(aq)

a.) Solution at equilibrium

b.) Increase in SCN-(aq)

c.) Increase in Fe3+(aq)

d.) Increase in FeSCN2+(aq)

  • Changes in Volume and Pressure

  • Changes in pressure primarily only concern gases.

  • Concentration of gases are greatly affected by pressure changes and volume changes according to the ideal gas law.

  PV = nRT

  P = (n/V)RT

  • Changes in Pressure and Volume

  Change  Shift in Equilibrium

Increase in Pressure  Side with fewest moles

Decrease in Pressure  Side with most moles

Increase in Volume  Side with most moles

Decrease in Volume  Side with fewest moles

  • Changes in Pressure and Volume

  • Changes in Temperature

  • Equilibrium position vs. Equilibrium constant

  • A temperature increase favors an endothermic reaction and a temperature decrease favors and exothermic reaction.

Change  Endo. Rx  Exo. Rx

Increase T  K decreases  K increases

Decrease T  K increases  K decreases

  • Changes in Temperature

Consider: N2O4(g) ↔ 2NO2(g)

The forward reaction absorbs heat; endothermic

  heat + N2O4(g) ↔ 2NO2(g)

So the reverse reaction releases heat; exothermic

  2NO2(g) ↔ N2O4(g) + heat

Changes in temperature??

  • Effect of a catalyst

  • How would the presence of a catalyst affect the equilibrium position of a reaction?

Principles of Chemical Equilibrium ( Dr. Bülent BELİBAĞLI )

Chemical Equilibrium: The state that is reached when the concentrations of reactantsand products remain constant over time is called the state of chemical equilibrium.

Laboratory‎ > ‎Chemical Equilibrium

One useful application of Beer’s Law is to determine the equilibrium constant of a reaction. Reactants (iron (III) and thiocyanate) are mixed in a ratio such that one reactant is presumed to be nearly 100% converted into a colored product. The product is assumed to follow Beer’s Law, and the relationship between relative absorbance and concentration is determined from the first experiment. After that, the reactant originally in excess is reduced in concentration. From the colorimetric determination of product concentration, the concentrations of reactant remaining at equilibrium is determined and the quantitative relationships among these concentrations are studied. One of these relationships, the mass action expression, is found to be nearly constant.