# Spontaneity of a chemical reaction based on the free energy change, Delta G

Submitted by rulzdapool01 on Fri, 03/30/2007 - 13:07

a) Write the balanced net ionic equation for the reaction which happens when aqueous silver nitrate is mixed with aqueous potassium chloride

b) Using the appropriate values from the attached thermodynamics tables (do not use others).  Calculate deltaH, DeltaG, and deltaS for the reaction in part a.

c) Is this reaction spontaneous as written?  Explain based on the value of deltag as well as what you remember from what you saw in lab when the two were mixed.

d) Does your value (especially sign) of  deltaS make sense?  Explain your reasoning. (i.e. does delta S agree with your net ionic equation indicates about changes in randomness from reactants to products)

e) In class we talked about driving forces in nature.  Which is more important in driving this reaction, enthalpy or entropy, or are they of equal importance?  Explain your reasoning.

f) This reaction is spontaneous at room temperature.  Are there any temperatures when this reaction is no longer spontaneous?  Explain.

g) Do you expect the equilibrium constant for this reaction to be large or small?  Explain your answer.  Remember, all reactions have equilibrium constants!

h) Using the correct thermodynamic information above, determine the equilibrium constant for this reaction.  Does your value support your answer in part e?  Explain why/why not.

i) Use the net ionic equation to write the equilibrium constant expression Kc for the formation of silver chloride.

j) Using Le Chatelier’s principle, describe the effect each of the following changes has on the equilibrium between silver chloride and its ions.  Your answer should state and explain which way the reaction moves (i.e towards reactants, towards products, or no change)
i. heating the reaction

k) When you mix solutions of Ag+ and Cl-, the precipitate forms immediately (at least, it appears to happen very quickly).  What does this tell you about the activation energy of the reaction?  I don’t want a number here, I am thinking in terms of: is it large or small, and why!

l) Construct a reaction coordinate diagram showing the relative positions of reactants and products. Use the delta g values you calculated earlier for proper placement.  Make sure that your activation energy barrier is reasonable based on your answer to part i.

m) Suppose you mix together 50.0 mL of 0.01M AgNO3 with 50.0 mL of 0.20 M KCl.
i. What are the initial concentrations of Ag + and Cl- (just after mixing but before any reaction begins you have made a total of 100.00 mL of solution so each reactant has been diluted)
ii. Use the value of the equilibrium constant you determined earlier to find the concentrations of both Ag+ and Cl- at equilibrium. ( hint: use an ice table)
iii. At this point, what could you do to form more AgCl?
iv. Do you expect the solution to be warmer, colder, or the same temperature as when you started?  Explain your reasoning.

please refrain from posting just a list of questions. our goal is not to do you homework for you but to help you when you get stuck. We are more than happy to help out but show us what you have first.

ok. how would you go about figureing out the degrees of formation for delta h, delta g, and delta s for the reaction? (part b)

You can look up values for S (entropy), H (enthalpy) and G (Gibbs free energy) usually in the appendix of the book you have.

These are listed as S and H of formation.

To get a change in S or change in H for a reaction you would simply add the values of the products (multiplied by the coefficient in the reaction) and subtract the sum of the values for the reactants (multiplie by the coeeficient in the reaction)

For instance,

aA + bB ---> cC

you would look up the value for enthalpy , H , in the appendix you would then do the following

change in H =  c x H of formation for C - (a x H of formation of A + b x H of formation of B)

you can also do this for S and G

I'd agree.  You should try and work as much as you can before you post because it'll help you come exam time.

However,

a)

Ag(NO3) (aq) + KCl (aq) AgCl(s) + K(NO3) (aq)

You get this by using your solubility rules:  Halides (Cl, Br...basically anything that's a negatively charged ion in solution) are insoluble (thus won't form a precipitate) UNLESS in combination with a silver, mercury, or lead ion.
Fact: Memorize your solubility tables; it's normally not given to you on an exam.

For your products, using the solubility rules, you will find that Nitrates (no3) have no exceptions and will generally be insoluble (thus, stay in aqueous solution) with any other element. (Other ions like nitrates are chlorates, bicarbonates (HCO3) and Group1A alkali metals.

j)  I have always found Le Chatlier's principle fascinating.  In a nutshell, and as a general rule, the principle states that when you have a system at equilibrium and you apply some type of stress on the system, the system will adjust itself to a new equilibrium.  So when you apply heat to one side of the equation above, you will may get more "stuff" on the other side of the stress.  Adding more AgCl will cause more of this (notice the arrow where it is pointing):

Ag(No3) + KCl AgCl + K(no3)

The term "Stress" is defined as a change in pressure, temperature, concentration, or volume to a system.

Woops didn't read the "net ionic equation part"...lol

net ionic eqn: Ag+(aq)  + Cl-(aq)  ----->AgCl(s)

It's probably easier working with your net ionic equation when you are dealing with Le Chatlier rather than your balanced chemical equation.