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This is the lab document. It has all the information for you to complete the lab at home. Download this document and perform the lab according to its procedures. Fill in the Report Sheet on your computer. Answer any questions and write a conclusion. Save the Report Sheet as a doc or docx file (do not save as any other file.) Please ONLY save and upload the Report Sheet, not the entire lab.Online CHEM 1412 Systems at Equilibrium Lab
Objectives:
1. To study three equilibrium systems.
2. To observe Le Châtelier’s principle at work: To see how a change will affect an
equilibrium system, i.e. how an equilibrium will shift in response to a perturbation.
3. To identify whether a reaction is exothermic or endothermic based on equilibrium shifts
in response to temperature changes.
Materials:
• 2 small test tubes
• test tube holder
• stir rod
• sterno burner and stand, or stove
• 10 mL graduated cylinder
• 100mL graduated cylinder
• plastic pipet
• 100 mL glass beaker*
• balance
• KCl (s) (30 g)
• 3M HCl (1 mL)
• Ammonia solution (1mL)
• NH4Cl (s) (1g)
• 0.10 M Cu(NO3)2 (1 mL)**
• Phenolphthalein (3 drops)
• Distilled water
• Ice
*If your beaker is not glass use the glass Erlenmeyer flask.
**Please see procedure for making this solution from the solid copper(II)nitrate
Introduction:
A reaction is at equilibrium if the rate of the forward reaction is the same as the rate of the
reverse reaction. At equilibrium, the concentrations of reactants and products remain constant.
A reversible reaction at equilibrium can be disturbed if a stress is applied to it. Examples of
stresses include increasing or decreasing chemical concentrations, or temperature changes. If
such a stress is applied, the reversible reaction will undergo a shift in order to re-establish its
equilibrium. This phenomenon was discovered in the 1800s by Frenchman Henri Louis Le
Châtelier (1850-1936). This is known as Le Chatelier’s Principle.
Effect of Change in Concentration
LeChâtelier’s principle predicts that when additional reactant or product is added to the system
the equilibrium will shift to the opposite side. Adding additional reactants is a stress that must be
counteracted by shifting away from the reactant side and producing more products. Similarly,
adding more of a product must be balanced by an equilibrium shift away from the product side
and make more reactants. In short, if the concentration of a reactant or product is increased, the
system moves to the side of the equation that consumes the excess.
When the concentration of a substance is decreased, the system shifts to the side of the reaction
that will replace the removed substance.
A substance can be decreased by removing it physically (if it is a solid or a gas) or chemically.
Sometimes, a chemical can be added which will react with one of the reactants or products and
remove the reactant/product from the system. Addition of this chemical serves to decrease the
species it reacts with. An example of this will be seen when we study and perturb the copper
equilibrium.
Temperature Changes
Temperature changes also exert a stress on the system. Heating or cooling either adds energy to
or removes energy from the system. The effect of a change in temperature on a system at
equilibrium depends on whether the reaction is endothermic (∆Hrxn>0) or exothermic (∆Hrxn < 0). H can be written as a term in the chemical equation. H will be a reactant in an endothermic reaction and a product in the exothermic reaction. Endothermic H + R →P R→P + H Exothermic If an equilibrium system experiences a temperature increase, we can view this as increasing the H term. For an exothermic reaction a temperature increase will shift the system to the left or reactant side. A decrease in temperature will shift the reaction to the right or product side. The following table shows the relationship between Heat of Reaction and temperature change. Exothermic “ Endothermic “ Temperature Equilibrium shifts ↑ ↓ ↑ ↓ ← → → ← Pressure Changes When the pressure inside the reaction container in a system at equilibrium is increased (a decrease in volume), the system shifts to the side of the reaction with the least number of gaseous moles. The reverse response occurs when the pressure is decreased. To illustrate Le Châtelier’s principle lets consider the Haber-Bosch process (an artificial nitrogen fixation process used to prepare 176 million metric tons of ammonia each year worldwide.) N2 (g) + 3 H2(g) ⇌ 2 NH3(g) North Lake College ∆H = - 92.38 kJ CHEM 1412 online Systems at Equilibrium Lab p. 2 • • • • Removing NH3 will cause the reaction to shift to the right to produce more NH3. Removing H2 will cause the reaction to shift to the left to produce more H2. A decrease in volume (or increase in pressure) will shift the reaction to the right since there are 4 moles of gaseous reactant and 2 moles of gaseous product. A temperature increase results in less ammonia being produced because the system shifts to the left. In this lab experiment, you will study 3 systems at equilibrium. For each system you will observe what happens after perturbing the system in various ways. The observed changes reveal the direction of the shift in response to the perturbation. Such shifts can be explained in terms of LeChatelier’s Principle. The three systems are: 1. Copper (II) ion Cupric nitrate in the presence of ammonia forms an equilibrium as shown. Dissolved copper (II) ions are in the aqua blue form of Cu(H2O)4 2+. This aqua blue complex reacts with ammonia to give a deep blue ion, Cu(NH3)4 2+. Cu(H2O)42+(aq) + 4NH3(aq) ⇌ Cu(NH3)42+(aq) + 4H2O(l) aqua blue cobalt blue 2. Aqueous Ammonia Ammonia is a weak base and partially ionizes in water to set up the following equilibrium. The hydroxide ions present will turn pink/magenta in the presence of the indicator phenolphthalein. NH3 (aq) + H2O (l) ⇌ NH4+1(aq) + OH-1(aq) clear pink (with phenolphthalein) 3. KCl dissolution Potassium chloride dissolves in water and will form an equilibrium in a saturated solution. KCl (s) ⇌ K+ (aq) + Cl–(aq) Procedure: This lab calls for thorough mixing of test tubes. The following video shows how to do this: mixing the contents of a test tube You may heat beakers on the alcohol burner provided in the lab kit. You may also heat beakers on your stovetop if you find this more convenient. North Lake College CHEM 1412 online Systems at Equilibrium Lab p. 3 An ice bath is a container of mostly ice with some water. The water is to ensure a good surface contact with the test tube that is put into the ice bath. To make 0.1M Cu(NO3)2 (aq) weigh 1.88 g of the solid and dissolve in 10 mL distilled water. Cu(II) system 1. 2. 3. 4. 5. 6. 7. 8. 9. Put 1.0 mL of 0.1M copper (II) salt solution in a clean test tube. Record observations. Use a pipet to add 3 drops of ammonia solution. Mix thoroughly. Record observations. Use a pipet to add 3 M HCl drop by drop, mixing after each added drop of HCl. Stop adding the drops of HCl when a color change occurs. Record observations. Fill in the table in the Report sheet to explain the observations in terms of equilibrium shift and Le Chatelier’s principle. Aqueous Ammonia System 1. Prepare a solution of aqueous ammonia by adding 1 mL of ammonia solution and 3 drops of phenolphthalein to a 100mL beaker, add 50.0 mL distilled water and mix with a stirring rod. 2. Place 3 mL of this solution into a small test tube. 3. Record observations. 4. Add a small amount (about ¼ teaspoon) of NH4Cl powder to the same test tube. 5. Mix thoroughly. 6. Record your observations. 7. Fill in the table in the Report sheet to explain the observations in terms of equilibrium shift and Le Chatelier’s principle. KCl system 1. Put 20 mL of distilled water in a 100-mL glass beaker (or Erlenmeyer flask if your beaker is not glass.) 2. Make a saturated solution by adding 30 g KI and dissolving as much salt as possible. Not all of this will dissolve since this amount is beyond the solubility of KI in water at room temperature. 3. Stir occasionally over a 15 minute time period to ensure all of the KI that will dissolve, dissolves. Record Observations. 4. Heat the beaker at low heat on your alcohol burner or stovetop until more (if not all) of the salt is dissolved. Avoid splashing and loss of water. Record your observations. 5. Let the beaker cool and then place it in an ice bath. 6. Record your observations. 7. Fill in the table in the Report sheet to explain the observations in terms of equilibrium shift and Le Chatelier’s principle. North Lake College CHEM 1412 online Systems at Equilibrium Lab p. 4 WASTE DISPOSAL: Flush all leftover solutions down the drain with lots of water. North Lake College CHEM 1412 online Systems at Equilibrium Lab p. 5 KCl Write the Equilibrium Equation: Step Process 3 Saturated KCl solution 4 + heat 6 - heat (Put on ice) → Observation Eq. shift Explanation n/a n/a From your observations, is the dissolution of KCl in water exothermic or endothermic? Justify your answer using Le Châtelier’s principle. Conclusion North Lake College CHEM 1412 online Systems at Equilibrium Lab p. 6 Online CHEM1412 Qualitative Analysis of Ions Objectives: 1. To become acquainted with qualitative analysis of cations and anions. 2. To determine the identity of ions in an unknown solution. Materials: • • • • • • • • • • • Small test tubes (6) large test tubes (3) Test tube rack 100 mL beaker 250 mL beaker Erlenmeyer flask Sterno burner and stand/stove Disposable pipettes (10) Water for bath Litmus paper (red or neutral) Distilled water • • • • • • • • • • • • • • • • • 0.1 M BaCl2* 1mL 0.1M Na3PO4 1mL 0.1 M Zn(C2H3O2)2 2mL 1.0 M Ca(C2H3O2)2 1mL 1.0 M Copper (II) gluconate 1mL 0.25g Epsom Salts, MgSO4..7H2O (to make 10 mL 0.1M) 0.05% NaF (or 0.01M) 1ml NaHCO3 (s) 1g NaOH (s) 1g Bleach 1mL 3M HCl 3mL (NH4)2CO3 (s) 0.5g 0.1M NaI 1mL Ammonia solution 1mL Mineral oil – 2 mL Unknown G – cation unknown Unknown H – anion unknown *Caution: Do not breath BaCl2 vapors. Do not get in eyes, on skin, or on clothing. Discard empty containers. Keep the container tightly closed in a cool, well-ventilated place. Introduction: Qualitative Analysis is a technique used to identify chemical species (e.g. molecules or ions.) The presence of chemical species can be confirmed by using various reagents that react selectively with them. Qualitative analysis is a technique that answers the basic question: ‘’What is it?” It differs from quantitative analysis that answers the question "How much do I have of it?” Titration and gravimetric analysis are examples of quantitative analysis. In this lab we will study qualitative analyisis techniques: • • Precipitation reaction - Some ions can be selectively precipitated and can thus be identified. For example if you have a solution that has either sodium or calcium ions, you can add fluoride ions. If a precipitate forms then your solution has calcium ions. If a precipitate does not form, the solution has sodium ions. Dissolution of a formed precipitate - Some cations are further distinguished by reactions that dissolve their precipitates. o Some metal cations form insoluble hydroxides that will redissolve when more hydroxide is added. This occurs when the metal cations are also • • Lewis acids (see lecture notes.) The zinc ion and copper (II) ion are both Lewis acids and will redissolve and form soluble complex ions. Cations that behave in this manner are amphoteric. In this lab, the zinc ion and copper (II) ion are first precipitated as insoluble hydroxides in alkaline media. Then, the solid hydroxides are converted to the soluble complex ions. Evolution of a gas – Some anions form a gas in the presence of acid or base. Color – Some chemical species are colored or will react to form a product with a color. Knowledge of these reactions and their colors helps to identify the chemical species. We will study the following ions in this at-home lab:* Type of Ion Cation Anion Formula Calcium ion Ca2+, zinc ion Zn2+, copper (II) ion Cu2+, ammonium ion NH4+ sulfate SO42-,carbonate CO32-, phosphate PO43-, bromide Br−, iodide I− First, we will perform “qual tests” on known solutions in order to observe what a positive test for that ion looks like. Then, we will perform the same tests on unknown ions to determine their identity. The most common qual test in this lab is mixing solutions together and seeing if a precipitate forms. A precipitate is a solid and it seems like it should be easy to identify but sometimes it is not. Some solids are so fine that they do not “settle” quickly and are dispersed throughout the liquid. In this case the liquid will look cloudy. If the liquid looks cloudy, then a precipitate has formed. If you let the test tube sit long enough the solid will settle and you will see the solid at the bottom and a transparent liquid above it. If the liquid is transparent, then a precipitate has not formed. Sometimes solutions are colored and this adds to the difficulty of identifying whether or not a precipitate forms. One thing to do is to shine a light (from your phone’s flashlight) through the test tube. If you put a piece of paper with letters between the test tube and the light you can identify cloudy vs. transparent by whether or not you can see the letters. See pictures below. The top picture shows a test tube that has a colored liquid which is transparent and therefore does not contain a precipitate. The bottom picure shows a test tube where the colored liquid is not transparent and therefore contains a precipitate. * Note: We study Fe3+, Ag+, Ni2+, and Cr3+ on campus. These give very beautiful precipitates with various reagents but we do not want you working with heavy metals and these reagents at home. Not only can they be toxic but they must not be flushed down the drain. They require special disposal protocols that are not possible at home. Qualitative Analysis of Cations Ammonium ion, NH4+ Ammoniuim ion will liberate ammonia gas in the presence of base according to the following molecular and NIE: (NH4)2CO3 (aq) + 2NaOH (aq) → 2NH3 (g) + 2H2O (l) + Na2CO3 NH4+ (aq) + OH- (aq) → NH3 (g) + H2O (l) Additionally, the odor of ammonia will provide further evidence for presence of ammonium ion. Calcium ion, Ca2+ Calcium ion will precipitate in the presence of fluoride ions according to the following net ionic equation (NIE): Ca2+ (aq) + 2F- (aq) → CaF2 (s) Zinc ion, Zn2+ Zinc ion will precipitate upon addition of a small amount of hydroxide ion to form insoluble zinc hydroxide according to the following NIE: Zn2+ (aq) + 2 OH- (aq) → Zn(OH)2 (s) When more hydroxide ions are added, the precipitate disappears as the soluble complex ion, zincate (Zn(OH)42-), forms: Zn(OH)2 (s) + 2 OH- (aq) → Zn(OH)42- (aq) Zincate clear-no color The zinc hydroxide is amphoteric since the zinc reacts as a lewis acid and the hydroxide reacts as a lewis (and Bronsted-Lowry) base. In the presence of either acid or base the insoluble precipitate disappears. Copper (II) ion, Cu2+ The copper (II) ion is also amphoteric. Cu(II) ion will precipitate upon addition of a small amount of hydroxide ion according to the following NIE: Cu2+ (aq) + 2OH- (aq) → Cu(OH)2 (s) When aqueous ammonia is added the precipitate disappears as the soluble complex ion, tetraamine copper (II) ion (Cu(NH3)4 2+) forms. This ion has a deep blue color. Cu(OH)2 (s) + 4 NH3 (aq) → Cu(NH3)4 2+ (aq) + 2 OH- (aq) tetraamine copper (II) ion blue Complex ions such as this form on the surface of metals and are partly the reason for the rich patinas metals form. See Appendix B for examples of this. Qualitative Analysis of Anions Sulfate Ion Sulfate will react with barium ion to produce a precipitate according to the following Molecular and NIE: BaCl2(aq) + MgSO4(aq) → BaSO4(s) + MgCl2 (aq) Ba2+(aq) + SO42-(aq) → BaSO4(s) Unlike the copper (II) and zinc precipitates, this precipitate will not dissolve when acid or base is added. Carbonate Ion Carbonates and bicarbonates react with acids to give odorless carbon dioxide gas and water according to the following molecular and NIE. NaHCO3(aq) + HCl(aq) → CO2(g) + H2O(l) + NaCl (aq) HCO3-(aq) + H+(aq) → CO2(g) + H2O(l) Phosphate Ion Phosphate will react with the zinc ion to produce a precipitate according to the following molecular and NIE: 2 Na3PO4 (aq) + 3 Zn(C2H3O2)2 (aq) → Zn3(PO4)2 (s) + 6 NaC2H3O2 (aq) 2 PO43- (aq) + 3 Zn2+ (aq) → Zn3(PO4)2 (s) Iodide Ion The iodide ion is oxidized by chlorine to produce elemental iodine. When iodide solutions are treated with chlorine water a reddish-brown color is observed due to formation of aqueous elemental iodine, I2 (aq). Oil is added to extract the iodine from the water into the oil layer. When in the oil, iodine shows a violet color. 2NaI (aq) + Cl2(aq) → I2 (aq) + 2NaCl (aq) 2I−(aq) + Cl2(aq) → I2 (aq) + 2Cl−(aq) reddish brown I2(aq) → reddish brown I2(oil) violet There are additional tests that confirm presence of these anions but due to their toxicity we will not perform these at home. Two such tests are the following: 1. Precipitation with silver ion - Reactions of halides (Cl-, Br-, I-) with silver ion gives silver halide precipitates of various colors. 2. Reaction of the solid anion with concentrated sulfuric acid - adding 3 drops of 18 M sulfuric acid to a small spatula tip (pea size) amount of solid salt containing the anion gives the following results: Anion SO42CO32 Cl − Br −. I− Observation of reaction of solid with 18M H2SO4 No Reaction A colorless odorless gas forms. CO32- (s) + 2H+ (aq) → H2O (l) + CO2 (g) A colorless gas forms with a sharp-pungent odor, gives an acidic test result with litmus, and fumes in moist air. Cl− (s) + H+ (aq) → HCl (g) A brownish red gas forms with a sharp odor, gives an acidic test result with litmus, and fumes in moist air. 2Br− (s) + 4 H+(aq) + SO42- (aq) → Br2 (g) + SO2 (g) + 2H2O (l) Solid turns dark brown immediately, with slight formation of violet fumes. Gas has odor of rotten eggs. Gives acidic test result with litmus, and fumes in moist air. 2I− (s) + 4H+ (aq) + SO42- (aq) → I2 (g) + SO2 (g) + 2H2O (l) (HI is also liberated) Observations Write what you see when you perform the qualitative tests. You are describing what you see to someone who is not seeing it so be specific and clear. Do not just say “it is blue.” Distinguish between the color of the solution and the color of any precipitate that forms. It is also not sufficient to say “it is blue” since the solution could have started off blue. Instead say “the solution color changed from clear to blue.” If there is no reaction, write “NR.” Know the distinction between an observation and a conclusion. For example it is an observation to see that the liquid in the test tube turned cloudy and white. It is a conclusion to say that a calciuim fluoride precipitate formed. It would be incorrect to put in the Observation column “carbon dioxide formed.” Instead you would say “I saw bubbles or fizziness.” Sometimes the precipitate may take time to settle or fall to the bottom of the test tube. If the solution looks cloudy that is evidence of a precipitate. Carefully observe your solutions to see if they are clear or hazy. Remember that a true solution is clear and transparent. Any haze at all indicates a precipitate albeit a sparse or fine one. Results This lab will require you to analyze your observations and write results. That is, for each qualitative test you perform you will write your interpretation of your observations. For example if you mix two solutions together and observe a cloudy mixture you can say “cloudiness indicates that a precipitate formed.” To help illustrate this I filled in the box on the Report Sheet for one of the cations and one of the anions. GLOSSARY Confirmation test - the final experiment done on a supernatant or a precipitate in order to prove or disprove evidence of the presence of a given cation or anion. Decantate - liquid obtained from a decantation process. Decantation - separating a l... Purchase answer to see full attachment

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