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Characteristics of Chemical Equilibrium The purpose of this lab is to investigate how changing various factors affect an equilibrium system. This lab consists of four independent activities. Each activity changes a component (concentration, temperature, and pressure) in an equilibrium system. Students observe how the changes affect the equilibrium system and analyze the changes using Le Chatelier's Principle. Please explain your answers to the following questions using Le Chatelier's Principle. Activity A: Nature of Equilibrium HIn (aq) ⇌ H+(aq) + In-(aq) …… (A) Colorless Pink H+(aq) + OH-(aq) → H2O (l) If HCl solution is added to the solution in equilibrium A, would the solution become more pink or less pink? If NaOH solution is added to the solution, which direction would the equilibrium shift to? Would the color become more pink or less pink when NaOH solution is added to the equilibrium? Activity B: Effect of Concentration Fe3+(aq) + SCN-(aq) ⇌ FeSCN2+(aq) ……… (B) Yellow Colorless Red Fe3+(aq) + H2PO4-(aq) → FeH2PO42+(aq) Yellow Colorless Colorless Which direction would the equilibrium shift to if Fe(NO3)3 solution is added to the equilibrium B? Which direction would the equilibrium shift to if Na3PO4 is added to the equilibrium B? Activity C: Effect of Temperature Fe3+(aq) + SCN-(aq) ⇌ FeSCN2+(aq) is an exothermic reaction. Yellow Colorless Red Would the color of the solution become more red or less red if the temperature is increased? What happens if the temperature is decreased? Co(H2O)62+(aq) + 4Cl-(aq) ⇌ CoCl42-(aq) + 6H2O (l) Pink Blue When the solution is heated to boiling, it turns from pink to blue. Is this reaction endothermic or exothermic? Bonus: if AgNO3 is added to the above blue solution, would the solution change color back to pink? If HCl is then added to the solution, would the color change again? Activity D: Effect of Pressure 2 CO2(g) + H2O(l) ⇌ CO2 (aq) + H+(aq) + HCO3-(aq) The reaction forms equilibrium in a large syringe with changeable volume. Would the solution become more acidic or less acidic when the volume is increased? What happens if the volume is decreased?

          Characteristics of Chemical Equilibrium

The purpose of this lab is to investigate how changing various factors affect an equilibrium system. This lab consists of four independent activities. Each activity changes a component (concentration, temperature, and pressure) in an equilibrium system. Students observe how the changes affect the equilibrium system and analyze the changes using Le Chatelier's Principle. Please explain your answers to the following questions using Le Chatelier's Principle.

Activity A: Nature of Equilibrium

HIn (aq) ⇌ H+(aq) + In-(aq) …… (A)
Colorless   Pink

H+(aq) + OH-(aq) → H2O (l)

If HCl solution is added to the solution in equilibrium A, would the solution become more pink or less pink?
If NaOH solution is added to the solution, which direction would the equilibrium shift to? Would the color become more pink or less pink when NaOH solution is added to the equilibrium?

Activity B: Effect of Concentration

Fe3+(aq) + SCN-(aq) ⇌ FeSCN2+(aq) ……… (B)
Yellow    Colorless  Red

Fe3+(aq) + H2PO4-(aq) → FeH2PO42+(aq)
Yellow    Colorless    Colorless

Which direction would the equilibrium shift to if Fe(NO3)3 solution is added to the equilibrium B?
Which direction would the equilibrium shift to if Na3PO4 is added to the equilibrium B?

Activity C: Effect of Temperature

Fe3+(aq) + SCN-(aq) ⇌ FeSCN2+(aq) is an exothermic reaction.
Yellow    Colorless  Red

Would the color of the solution become more red or less red if the temperature is increased? What happens if the temperature is decreased?

Co(H2O)62+(aq) + 4Cl-(aq) ⇌ CoCl42-(aq) + 6H2O (l)
Pink                    Blue

When the solution is heated to boiling, it turns from pink to blue.

Is this reaction endothermic or exothermic?

Bonus: if AgNO3 is added to the above blue solution, would the solution change color back to pink? If HCl is then added to the solution, would the color change again?

Activity D: Effect of Pressure

2 CO2(g) + H2O(l) ⇌ CO2 (aq) + H+(aq) + HCO3-(aq)

The reaction forms equilibrium in a large syringe with changeable volume.

Would the solution become more acidic or less acidic when the volume is increased? What happens if the volume is decreased?

Added by Alexis W.

Title: Equilibrium Shifts and Le Chatelier's Principle in Acid-Base Reactions (i) Observation when H2SO4 is added to the K2CrO4 solution Use net ionic equation and write out the pertinent equilibrium that illustrates what happened when H2SO4 was added to the K2CrO4 solution. Indicate the stress applied, the shift in equilibrium, and the concentration changes of all reactants and products in the resulting equilibrium. Equation: _____________+______________<--->___________+_________ yellow orange Stress: ______________________ Shift: _______________________ New Equilibrium: _______+_________<--->_______+_______ {Increase (I), Decrease (D)} (ii) Observation when NaOH is added to the K2CrO4 solution. _______________________________ _________________________________________________________________________________ (iii) Briefly explain how OH- exerts an effect on the CrO42-/Cr2O72- equilibrium? ________________ ________________________________________________________________________________ Provide a balanced net ionic equation to support your argument. _______________ + _______________ → _______________ (iv) Interpret your observations in terms of the equilibrium described in part (i) and Le Chatelier’s principle. Weak Acid–Base Indicator Equilibria (i) Describe the effects of adding... HCl to methyl orange ____RED________ HCl to phenolphthalein ___CLEAR_________ NaOH to methyl orange ____YELLOW________ NaOH to phenolphthalein ____PINK________ Interpret these effects in terms of the indicator equilibrium equation given below and Le Chatelier’s principle. HIn + H2O <--> H3O+ + In- Weak Acid–Weak Base Equilibria (i) Complete the equation for the dissociation equilibrium of acetic acid in water. CH3COOH + H2O<--->_____________ Predict the direction of the shift in the above equilibrium when 1 M sodium acetate is added to 0.1 M acetic acid containing methyl orange indicator. __________________________________________ Are the observed changes to the methyl orange indicator consistent with Le Chatelier’s principle upon the addition of 1 M sodium acetate? Briefly explain. (ii) Complete the equation for the reaction of ammonia with water: NH3 + H2O <-->_________________ Describe any observed odor or color changes when you add the following to 0.1 M NH3(aq) containing phenolphthalein indicator: NH4Cl - ________________________________________ HCl - ________________________________________ Indicate in which direction, left (←) or right (→), does each reagent above shift the equilibrium for the reaction of NH3 with water? Adding NH4Cl = adding HCl = Explain clearly the shift indicated in the ammonia with water equilibrium using Le Chatelier’s principle and the observed color change in the phenolphthalein indicator. When NH4Cl is added - ___________________________________________________________________________ ___________________________________________________________________________ ___________________________________________________________________________. When HCl is added – Start with the net ionic equation between HCl and NH3. ___________________________________________________________________________ ___________________________________________________________________________ ___________________________________________________________________________. Le Chatelier’s Principle 2. Complex ion equilibria (a) The Thiocyanatoiron(III) Complex Ion Write the equation for this equilibrium {Eq’n (5)}. _____+_____<-->______+______ The complex ion, [Fe(H2O)5(SCN)]2+, is responsible for the observed red color. Compare the intensity of the red color upon addition of each of the following with the intensity of the color in tube 4. Interpret your observations in terms of the equation describing the equilibrium and Le Chatelier’s principle. Fe(NO3)3 - LIGHTER / DARKER -red; equilibrium shifts- LEFT OR RIGHT {circle your choices} KSCN - LIGHTER / DARKER -red; equilibrium shifts- LEFT OR RIGHT {circle your choices} NaOH – Red color LIGHTER /DARKER/ DISAPPEARS ; equilibrium shifts- LEFT OR RIGHT {circle your choices} The Temperature-Dependent Equilibrium of Co(II) Complex Ions Write the equation for this equilibrium [Equation (6)]. What happens when you add 12 M HCl (or Cl-) to the pink (aquo) complex? solution turns blue-purple / pink; equilibrium shifts- left right {circle your choices} What happens when you heat the pink (aquo) complex and when you cool it? heat solution turns blue-purple / pink; cool solution turns blue-purple / pink {circle} Explain why these results occur in terms of the equation describing the equilibrium and Le Chatelier’s principle. 3. The equilibria of saturated solutions Saturated Sodium Chloride What happens when 12 M HCl is added to saturated (5.4 M) NaCl? nothing /solution turns transparent/ a precipitate forms {circle your choice} The [Cl-] in the original saturated solution is 5.4 M. What is the Cl- concentration (i) just after the addition of the 12 M HCl (before anything occurs)? Explain this behavior in terms of the [Cl-] in the two solutions and in terms of the equilibrium equation NaCl(s) <--> Na+(aq) + Cl- (aq). Saturated Barium Chromate Write a balanced equation for the reaction that occurs when K2CrO4 and BaCl2 solutions are mixed. Write a balanced equation describing the equilibrium that is established upon mixing K2CrO4 and BaCl2 solutions. Explain the changes observed when this equilibrium (above) is treated with 6 M HCl. 4. Application of the law of chemical equilibrium to solubility equilibria Le Chatelier’s Principle Solutions mixed Relative amounts of precipitate, if any (1) CaCl2(aq) + H2C2O4(aq) (2) CaCl2(aq) + K2C2O4(aq) (3) Results of (1) + 6 M HCl (4) Results of (3) + 6 M NH3 (5) CaCl2(aq) + 6 M NH3 1 Ca2+ + C2O42- <--> CaC2O4(s) + 3 H3O+ + NH3 <--> NH4+ + H2O 2^ 4 HC2O4- + H3O+ -> H2C2O4 + H2O + H2O Answer questions (a) and (d), relating these data on the relative amounts of CaC2O4(s) formed, or dissolved, under various conditions to the following coupled equilibrium equations. The equilibria are said to be coupled because each equilibrium shares at least one species in common with another equilibrium. (a) Account for the difference in behavior of H2C2O4 and K2C2O4 in tubes 1 and 2 toward CaCl2. Specifically, why does one solution form so much more CaC2O4(s)? {Hint: Consider the nature of the electrolytes.} (b) From this array of coupled equilibria, explain the effect of adding HCl on these equilibria. {Note that H3O+ participates in equilibria 2, 3, and 4.} Why does the addition of HCl change the amount of CaC2O4(s)? (c) Consider solutions 4 and 5, where NH3 was added. How does ammonia affect the amount of CaC2O4(s) present in solution 4? (i) Does Ca(OH)2(s) form when NH3 is added to a Ca2+ solution (solution 5)? ____________ (d) Why does the H3O+ concentration have more influence on the precipitation of the salt of a weak acid, such as CaC2O4(s), than on the precipitation of the salt of a strong acid, such as AgCl(s) or CaSO4(s)?

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