Equilibrium Constants
Equilibrium is a state of balance between products and reactants. When the rate of the forward reaction is equal to the rate of the reverse reaction, the concentrations of products and reactants remain constant. Every reaction is striving to reach equilibrium.
- A variation of Type 1 problems is when you are given the Keq and all the equilibrium concentrations except one and you are asked to calculate that one. The solution for this type of problem is simply writing out the Keq expression, filling in what you know and solving for the unknown. Read through this example: At 200°C, the Keq for the.
- Calculate Keq for 2 moles of electron(s) in each half-reaction and a standard potential of 0.55 V. Standard Electrode Potential: The measure of the individual potential for a reversible electrode.
Le Chatelier’s Principle : if equilibrium is disturbed (eg. if product or reactant concentrations are altered), the direction of the reaction will shift to restore the state of equilibrium.
Equilibria are reflections of Le Chatelier’s Principle. If we begin a reaction with only the reactants, they will readily form products. As the concentration of the products increases, the rate of the reverse reaction increases. As the concentration of reactants decreases, the rate of the forward reaction decreases. Eventually the reaction rates will cancel each other out, and there will be no net change in the concentration of reagents.
Perform calculations to find the value of Keq and concentrations of substances within an equilibrium system, and use these values to make predictions on the direction in which a reaction may proceed (ACSCH096) qualitatively analyse the effect of temperature on the value of Keq (ACSCH093). Keq equation calculator: calculate the equilibrium constant for the decomposition of water: how to calculate equilibrium constant for a reaction: calculate standard free energy change from equilibrium constant: calculate solubility product constant: calculate the cell potential the equilibrium constant and the free energy change for ba. This chemistry video tutorial provides a basic introduction into how to solve chemical equilibrium problems. It explains how to calculate the equilibrium con.
The Equilibrium Constant(Keq) describes the ratio of products to reactants in a state of equilibrium.
How To Calculate Keq
- Some reactions have a big Keq. When these reactions reach equilibrium, there will be a lot of product, and little reactant left.
- Some reactions have a small Keq. Don’t expect these reactions to produce much product.
- See that the Keq for the reverse reaction will be the reciprocal of the forward reaction.
Strong acids and bases dissociate with large equilibrium constants.
We use a specific type of equilibrium constant when describing the dissociation reactions of acids and bases. Recall the equation for the dissociation of an acid:
The Keq of this reaction is given by. We call it Ka, the Acid Dissociation Constant
Likewise, the equilibrium constant for a base dissociation is given by Kb, its Base Dissociation Constant.
The Dissociation Constant determines the strength of the Acid or Base
Strong vs weak Acids
Here is a list of many acids along with their Ka values. The top three on the list would be considered strong acids – they will essentially dissociate completely. Most acids are weak, though two weak acids can vary greatly in Ka.
Most important difference between strong and weak acids: once the reaction reaches equilibrium, strong acids will be dominantly in the conjugate base form. Weak acids will exist as a mixture of the original acid and the conjugate base form.
For simplicity, the strength of the acid will often be given as pKa. pKa represents the same information as Ka except its logarithmic scale eliminates the need for exponential notation. This simple equation converts between Ka and pKa.
Using these equations, see how a large pKa denotes a very weak acid. Bases have very large pKa
Bases use Kb and pKb as measurement of how favorably they accept protons (basicity). Kb works exactly like ka. A small Kb means equilibrium highly favors conjugate acid formation, indicating a strong base. A larger Kb indicates a weak base that will partially form conjugate acid and partially remain in its base form.
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General Note: A stronger acid or base will produce a weaker conjugate. A weaker acid or base will produce a stronger conjugate.
Summary
Here is an example of an acid base reaction. The Keq is 1.8e-5
Do we start with an acid or a base?
NH3 (ammonia) is a base. This is clear because the products include its protonated form NH4 (ammonium). We see it is acting like a base and accepting a proton from water.
The conjugate acid is ammonium. In the reverse reaction, ammonium acts as an acid and hydroxide acts as a base. Ammonium donates a proton to hydroxide.
Find Kb for ammonia: This would simply be the given Keq for the reaction. Kb is just the Keq of a specific type of reaction between a base and water. Kb = 1.8e-5. Ammonia is a weak base.
Find Ka of conjugate acid. This is would equal the Keq of the reverse reaction. The Keq of the reverse reaction is the reciprocal of the the forward reaction. Ka = 1/Kb = 5.6e-10. Since ammonia is a base of moderate strength (though defined as ‘weak’), see how its conjugate acid is is very very weak. This makes sense because if the conjugate acid were stronger, the reverse reaction would be more favorable, causing the forward base reaction to become less favorable.
Convert Ka and Kb to pKa and pKb.
pKb = 4.75 pKb = 9.25
Note: pKa + pKb = 14.0
After reading this you should be able to:
- make a prediction about the properties of a molecule by looking at its pKa value.
- label the acid, base, conjugate acid, and conjugate base in a reaction
- Use Keq to determine concentrations of reagents.
How To Find The Equilibrium Constant
Calculating Equilibrium Constants
We need to know two things in order to calculate the numeric value of the equilibrium constant:
- the balanced equation for the reaction system, including the physical states of each species. From this the equilibrium expression for calculating Kc or Kp is derived.
- the equilibrium concentrations or pressures of each species that occurs in the equilibrium expression, or enough information to determine them. These values are substitued into the equilibrium expression and the value of the equilibrium constant is then calculated.
- Write the equilibrium expression for the reaction.
- Determine the molar concentrations or partial pressures of each species involved.
- Subsititute into the equilibrium expression and solve for K.
c , for the system shown, if 0.1908 moles of CO2
, 0.0908 moles of H2, 0.0092 moles of CO, and 0.0092 moles of H2O vapor were present in a 2.00 L reaction vessel were present at equilibrium.
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- Write the equlibrium expression for the reaction system.
- Since Kc is being determined, check to see if the given equilibrium amounts are expressed in moles per liter (molarity). In this example they are not; conversion of each is requried.
[CO2] = 0.1908 mol CO2/2.00 L = 0.0954 M
[H2] = 0.0454 M
[CO] = 0.0046 M
[H2O] = 0.0046 M - Substitute each concentration into the equilibrium expression and calculate the value of the equilibrium constant.
Calculating K from Initial Amounts and One Known Equilibrium Amount
- Write the equilibrium expression for the reaction.
- Determine the molar concentrations or partial pressures of each species involved.
- Determine all equilibrium concentrations or partial pressures using an ICE chart.
- Substitute into the equilibrium expression and solve for K
N2 + 2 H2O' height=32 src='EquilibriumArt/CalcKexample2.gif' width=326 >
- Write the equilibrium expression for the reaction.
- Check to see if the amounts are expressed in moles per liter (molarity) since Kc is being . In this example they are.
- Create an ICE chart that expresses the initial concentration, the change in concentration, and the equilibrium concentration for each species in the reaction. From the chart you can determine the changes in the concentrations of each species and the equilibrium concentrations. From the example, we start with the folowing information.
Initial Concentration (M) | ||||
Change in Concentration (M) | - 2 x | - 2 x | + x | + 2 x |
Equilibrium Concentration (M) |
The change in concentration of the NO was (0.062 M - 0.100M) = - 0.038 M. Thus -2 x = - 0.038 and x = 0.019. Note: the negative sign indicates a decreasing concentration, not a negative concentration. The changes in the other species must agree with the stoichiometry dictated by the balance equation. The hydrogen will also change by - 0.038 M, while the nitrogen will increase by + 0.019 M and the water will increase by + 0.038 M. From these changes we can complete the chart to find the equilibrium concentrations for each species.
Initial Concentration (M) | ||||
Change in Concentration (M) | ||||
Equilibrium Concentration (M) |
- Substitute the equilibrium concentrations into the equilibrium expression and solve for Kc.
Calculating K from Known Initial Amounts and the Known Change in Amount of One of the Species
- Write the equilibrium expression for the reaction.
- Determine the molar concentrations or partial pressures of each species involved.
- Determine all equilibrium concentrations or partial pressures using an ICE chart.
- Substitute into the equilibrium expression and solve for K.
2 NO2' height =29 src='EquilibriumArt/CalcKexample3.gif' width=181 NOSAVE>
- Write the equilibrium expression to find Kp.
- Check to see that the given amounts are measured in appropriate pressure units since Kp is to be . In this example they are (atmospheres).
- Create an ICE chart and calculate the changes in pressure and equilibrium pressures for each species.
Initial Pressure (atm) | ||
Change in Pressure (atm) | ||
Equilibrium Pressure (atm) |
- Substitute the equilibrium pressures into the expression for Kp and solve for Kp.