General, Organic and Biological Chemistry Kit
Materials Included In Kit
Bromcresol green indicator, 0.04%, 35 mL
Iron(III) nitrate solution, Fe(NO3)3, 0.1 M, 75 mL
Hydrochloric acid, HCl, 0.1 M, 75 mL
Potassium thiocyanate solution, KSCN, 0.1 M, 75 mL
Sodium hydroxide solution, NaOH, 0.1 M, 75 mL
Sodium phosphate (monobasic) solution, NaH2PO4, 0.1 M, 40 mL
Pipets, Beral-type, graduated, 96*
*May be dispensed using dropper bottles to reduce disposable pipets needed.
Additional Materials Required
Beakers, 50-mL, 12
Beakers, 250- or 400-mL, 6*
Hot plates, 3*
Stirring rods, 12
Test tubes, 13 x 100 mm, 72
Test tube racks, 12
Wash bottles, 12
Water, distilled or deionized
*May be shared.
Hydrochloric acid and sodium hydroxide solutions may cause skin and eye burns. Potassium thiocyanate solution is irritating to skin and eyes. Iron(III) nitrate solution is also a possible skin and body tissue irritant; it will stain clothes and skin. Avoid contact of all chemicals with eyes and skin. Clean up all chemical spills immediately. Wear chemical splash goggles, chemical-resistant gloves and a lab coat or chemical-resistant apron. Remind students to wash hands thoroughly with soap and water before leaving the lab. Please review current Safety Data Sheets for additional safety, handling and disposal information.
Please consult your current Flinn Scientific Catalog/Reference Manual for general guidelines and specific procedures, and review all federal, state and local regulations that may apply, before proceeding. The leftover reaction mixtures may be rinsed down the drain with excess water according to Flinn Suggested Disposal Method #26b. Save all remaining solutions and reagents for future use.
- The laboratory work for this experiment can reasonably be completed within a typical 2-hour lab period. The Prelab Assignment should be completed by students as preparation for lab, and should be reviewed during prelab discussion.
- Enough disposable pipets are recommended in the Materials section so that each pair of students has a separate pipet for each solution. Encourage students to label their pipets to avoid contamination and waste. The pipets may be colorcoded, for example, using colored tape. If the reagents are placed in dropping bottles, the number of pipets may be reduced. To avoid congestion in the materials dispensing area, stagger the starting points for the two parts of the experiment. (Half the student groups start with the complex-ion equilibrium, the other half with the acid–base equilibrium.)
- Reaction of iron(III) nitrate with potassium thiocyanate may also be viewed as a double replacement reaction. The products are a series of complex ions having the general formula Fe(SCN)n3–n, where n = 1– 4. Although the red product is generally represented as FeSCN2+, all of the products are possible and all are deep red. The neutral compound Fe(SCN)3 can be extracted from aqueous solution using ether.
- The effect of temperature on the equilibrium in the Fe3+/FeSCN2+ system may be counterintuitive. Most students rightly expect that the rate of a reaction always increases with temperature. This is true, but the rates of the forward and the reverse reactions will increase by different amounts because they have different activation energies. In the Fe3+/FeSCN2+ example, the rate of the reverse reaction must increase more than the rate of the forward reaction when the temperature is raised.
- In contrast to the effect of adding or removing reactants and products on the position of equilibrium, the effect of temperature cannot be generalized without further information. The equilibrium constant for a reaction depends on temperature. When Le Chatelier’s principle is introduced, students can determine whether a reaction is exothermic or endothermic as written based on the effect of temperature on the position of equilibrium. In this exploratory introduction to equilibrium, the temperature effect may be discussed in terms of the rates of the forward and reverse reactions.
- The green transition color (step 16) in the reversible reaction of bromcresol green is easy to overshoot. Students should carefully add HCl drop by drop and gently swirl the solution between drops. Have students try to add half a drop at a time. (Squeeze out a small amount from the pipet, press the pipet tip against the side of the test tube to dislodge the halfdrop, and then swirl the test tube contents to mix the half-drop into solution.) The pH range for the color transition is 3.8–5.4.
- Moisture-sensitive test paper coated with cobalt chloride is available under the trade name Hydrion Humidicator Paper (see Flinn Catalog No. AP4656). The test strips are coated with anhydrous cobalt chloride that changes from blue to pink when it is hydrated. A color chart is available to estimate relative humidity levels between 20 and 80%.
- Equilibrium is one of the most challenging topics in the chemistry curriculum. The dynamic nature of equilibrium, in particular, seems to be prone to misunderstanding. The purpose of this experiment is to allow students to discover the key principles of equilibrium based on direct observations and logical reasoning. The reactions A and B can occur in both directions. Why then do the reactions not proceed to completion under conditions where reactants and products are clearly present? Is it possible that some molecules react, but that others do not? This is the hard part! Some students will probably not be able to “make the leap” and conclude that both reactions are occurring. Those students who accept the idea that both the forward and reverse reactions must take place simultaneously will then easily conclude that they must occur at the same rate (otherwise the color would change).
- The activity of hemoglobin, the main oxygen-binding protein in red blood cells, illustrates a biomedical application of complex-ion equilibrium. Hemoglobin (Hb) contains four iron(II) ions that bind to oxygen molecules. This is a reversible reaction, since hemoglobin must release the bound oxygen molecules in cells and body tissues (Equation 5).
Students should be able to apply what they have learned in this experiment to explain the effects of high altitudes on humans. At high altitudes, where the concentration of oxygen is lower, the equilibrium shown in Equation 5 is shifted in the reverse direction. Less oxygen is therefore available in the bloodstream to be transported to the cells. The physical symptoms of reduced oxygen availability are fatigue and dizziness. The human body, however, is marvelous in its adaptability. People who live or train at high altitudes compensate for the reduced oxygen supply by synthesizing more red blood cells. Increasing the concentration of hemoglobin increases the rate of the forward reaction and thus increases the amount of available oxygen.
- An interesting application of acid–base equilibrium is found in the carbonic acid–bicarbonate buffering system in the blood (Equation 6). The “blood buffer” is able to maintain blood pH within a very narrow range despite the acidifying effect of carbon dioxide produced by metabolism (approximately 350 g of CO2 are produced per 1000 Calories of food burned). People with impaired lung function, due to emphysema, for instance, are not able to exchange CO2 efficiently between the lungs and air as they exhale (Equations 7 and 8). The result is an increase in the amount of carbonic acid in the blood, making the blood more acidic. This condition is called respiratory acidosis.
- A particulate model may help students visualize the presence of both reactants and products at equilibrium. Ask students to draw pictures of what they imagine the Fe3+–SCN– system looks like at equilibrium. Then ask them to add more reactants or products to their picture and predict what a new equilibrium position would look like.
Correlation to Next Generation Science Standards (NGSS)†
Science & Engineering Practices
Asking questions and defining problems
Developing and using models
Planning and carrying out investigations
Analyzing and interpreting data
Engaging in argument from evidence
Obtaining, evaluation, and communicating information
Disciplinary Core Ideas
MS-PS1.A: Structure and Properties of Matter
MS-PS1.B: Chemical Reactions
HS-PS1.A: Structure and Properties of Matter
HS-PS1.B: Chemical Reactions
Cause and effect
Scale, proportion, and quantity
Systems and system models
MS-PS1-1. Develop models to describe the atomic composition of simple molecules and extended structures.
MS-PS1-2. Analyze and interpret data on the properties of substances before and after the substances interact to determine if a chemical reaction has occurred.
HS-PS1-5. Apply scientific principles and evidence to provide an explanation about the effects of changing the temperature or concentration of the reacting particles on the rate at which a reaction occurs.
HS-PS1-1. Use the periodic table as a model to predict the relative properties of elements based on the patterns of electrons in the outermost energy level of atoms.
HS-PS1-2. Construct and revise an explanation for the outcome of a simple chemical reaction based on the outermost electron states of atoms, trends in the periodic table, and knowledge of the patterns of chemical properties.
HS-PS1-6. Refine the design of a chemical system by specifying a change in conditions that would produce increased amounts of products at equilibrium.
Answers to Prelab Questions
- True or False: At equilibrium, no more reactants are transformed into products. If false, rewrite the statement so that it correctly describes the nature of chemical equilibrium.
False. At equilibrium, the rate at which reactants are transformed into products is equal to the rate at which products are transformed back into reactants.
- True or False: At equilibrium, the concentrations of reactants and products are equal. If false, rewrite the statement so that it correctly describes the nature of chemical equilibrium.
False. At equilibrium, the concentrations of reactants and products are constant. (Alternatively: At equilibrium, the forward and reverse reaction rates are equal.) Note: The concentrations of reactants and products may be equal, but that would be a special case.
- Paper coated with cobalt chloride is sold commercially as moisture-sensitive test strips to estimate relative humidity levels between 20 and 80 percent in air. The following reversible reaction takes place with water:
- What color do you think the paper will be when the humidity is low (20%)? What color will it be when the humidity is high (80%)? Explain.
The paper should be blue when the relative humidity is low (20%), because a low concentration of the reactant H2O means that the rate of the forward reaction will also be low. When the relative humidity is high (80%), and the concentration of water needed to react with CoCl2 is greater, the paper should turn pink.
- The test strips come with a color chart to estimate intermediate humidity levels. Predict the intermediate color that might be observed when the humidity is about 50%. State your reasoning.
At intermediate humidity levels, appreciable amounts of both reactants and products should be present at equilibrium, and an intermediate or transition color should be observed that is midway between blue and pink—lavender.
- The effect of adding more reactant to a system at equilibrium is often explained or predicted using Le Chatelier’s principle. Look up this principle in a textbook or online and restate it here to explain the color change predicted in Question 3a at high humidity.
Le Chatelier’s principle states that if an equilibrium system is subjected to a stress, the system will react in such a way as to reduce the stress. In the context of Question 3a, adding or increasing the concentration of water vapor at high humidity is a stress. The cobalt chloride system reacts with the additional water vapor, thereby reducing its amount (tbe stress), forming the hydrate and turning pink.
Complex-Ion Equilibrium of Iron(III) and Thiocyanate Ions
Acid–Base Equilibrium of Bromcresol Green
Answers to Questions
- Write the chemical equation for the reversible reaction of iron(III) ions with thiocyanate ions. Label this Equation A. Use the information in the data table to write the color of each reactant and product underneath its formula.
- How did the color of the solution change when additional reactant—either Fe(NO3)3 in step 6 or KSCN in step 7—was added? Explain the observed color changes: Adding more reactant to an equilibrium mixture of reactants and products increases the rate of the (forward/reverse) reaction and thus (increases/decreases) the amount of product.
Adding extra Fe(NO3)3 or extra KSCN produced the same effect—the color of the solution changed from orange to dark red. The dark red color indicates that more product was formed under these conditions. Adding more reactant to an equilibrium mixture of reactants and products increases the rate of the forward reaction and thus increases the amount of product.
- How do the results obtained in steps 6 and 7 demonstrate that both reactants and products are present at equilibrium?
Adding either reactant alone increased the amount of product. Since additional product formed when either reactant was added, the other reactant must already be present in solution. Both reactants and products are present at equilibrium.
- How did the color of the solution change when it was cooled (step 8) or heated (step 9)? How do these results demonstrate that the reaction shown in Equation A does indeed occur in both the forward and reverse directions?
Opposite color changes were observed when the control solution was cooled or heated. The original orange solution turned red-orange when it was cooled (step 8), yellow when it was heated (step 9). These results indicate that the reaction can indeed “go both ways.” The stock solution must contain equilibrium concentrations of reactants and products. When the solution was cooled, the concentration of the red product increased (net reaction in forward direction). When the solution was heated, the concentration of the red product decreased and more of the reactants were formed (net reaction in reverse direction).
- In step 10, H2PO4– ions combined with iron(III) ions and removed them from solution. How did the color of the solution change when NaH2PO4 was added? Explain the observed color change: Removing one of the reactants from an equilibrium mixture of reactants and products decreases the rate of the (forward/reverse) reaction and thus (increases/decreases) the amount of product.
Adding sodium phosphate decolorized the solution—the red color disappeared and the solution turned light yellow and cloudy. The amount of product decreased. Removing one of the reactants from an equilibrium mixture of reactants and products decreases the rate of the forward reaction and thus decreases the amount of product.
- After observing the effect of NaH2PO4 on the equilibrium mixture in step 10, a student was skeptical that both Fe3+ and SCN– ions were still present. Suggest additional experiments that could be done to prove that both reactants are still present at this point.
Try adding more of either reactant separately to the reaction mixture at this point. If Fe3+ is still present in solution, it will react with the added SCN–. Similarly, if SCN– is still present in solution, it will react with the added Fe3+. In either case, the solution should turn red again. Note: Try it! It works.
- Write the chemical equation for the reversible reaction of bromcresol green with water. Label this Equation B. Hint: Refer to Equation 3 in the Background section.
- Use the color changes observed for the indicator before and after adding HCl (steps 12 and 13) to predict the colors of the HIn and In– forms of bromcresol green. Write the colors of HIn and In– underneath their formulas in Equation B. Explain your reasoning. Hint: Adding HCl increases the concentration of H+ ions. Which reaction, forward or reverse, would that increase?
The colors of HIn and In– are shown in Equation B. The colors can be inferred based on the color change observed when HCl was added to the initial indicator solution. The initial indicator color was blue; when HCl was added, the indicator turned yellow. Adding H+ ions (in the form of HCl) should increase the rate of the reverse reaction. When a new equilibrium is re-established, there will be a greater concentation of HIn. The yellow color must be due to HIn, the blue color to In–.
- Explain the observed color change: Adding more product to an equilibrium mixture of reactants and products increases the rate of the (forward/reverse) reaction and thus (increases/decreases) the amount of product.
Adding more product to an equilibrium mixture of reactants and products increases the rate of the reverse reaction and thus decreases the amount of product.
- In step 14, hydroxide ions reacted with and removed H+ ions from solution (see Equation 4 in the Background section). What color change was observed when NaOH was added? Explain the observed color change: Removing one of the products from an equilibrium mixture of reactants and products decreases the rate of the (forward/reverse) reaction and thus (increases/decreases) the amount of product.
The indicator color changed from yellow to blue when NaOH was added. Adding NaOH increased the amount of product present at equilibrium. Removing one of the products from an equilibrium mixture of reactants and products decreases the rate of the reverse reaction and thus increases the amount of product.
- What form(s) of the indicator were most likely present when the transition color was observed in step 16? How does this observation provide visual proof that not all reactions “go to completion?”
The transition color of bromcresol green is green. The green color, midway between yellow and blue, suggests that approximately half of the available indicator molecules are present in the uncharged form HIn (yellow) and half in the ionic form In– (blue). The green color offers visual proof that both reactants and products are present.