THIS LAB IS CHAOS! 2. In liquids or gases? Explain.

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1 THIS LAB IS CHAOS! PRELAB INTRODUCTION Part 1 We are already familiar with the Enthalpy (H) for a reaction. We know that if a reaction gives off heat, that it is considered exothermic and has a negative ΔH value, and that if a reaction absorbs heat, it is endothermic and has a positive ΔH value. Another important concept is Entropy. A simplified definition of Entropy (S) is a measure of the disorder, chaos, or randomness of a system or process. Natural processes tend to go toward greater states of disorder or chaos. This is an increase in entropy (S). Thinking Questions (you need to consider the particulate level) 1. Is entropy greater in solids or liquids? Explain. 2. In liquids or gases? Explain. 3. a. Examine the following processes and determine if entropy is increasing or decreasing. b. Label ΔS with a (+) if entropy is increasing or ( -) if entropy is decreasing Process ΔS ΔH Ice melting at room temperature Water freezing at room temperature Steam condensing at room temperature Water boiling at room temperature Salt dissolving (feels cold) Add a metal to an acid and generate a gas (gets hot) c. Discuss your choices and the reasons for those choices with your lab partner. d. Can you identify those processes above that are exothermic and those that are endothermic? Discuss and label with your lab partner ΔH as (-) or (+). Part 2: How can we determine when a process occurs spontaneously? J. Willard Gibbs, a professor at Yale, developed an equation to determine when a process occurs spontaneously. But first what is a spontaneous process? A spontaneous process is one that occurs if a system left to itself with no external action needed to make it happen. Whether or not a process is spontaneous Processes that are exothermic tend to occur spontaneously, but not always. Which of the above processes do you think are spontaneous at room temperature? Discuss with your lab partner and mark with a check. Is there a pattern? What factors do you think will determine if a process is spontaneous? 1

2 LAB ACTIVITY MATERIALS LIST (per lab group) CHEMICAL LIST 2 - Test tubes Potassium nitrate Insulating holder- Styrofoam Calcium chloride Thermometer Distilled water 10 ml Graduated cylinder Electronic Balance SAFETY CONCERNS Potassium nitrate is a very strong oxidizer. There is fire and explosion risk when it is heated or in contact with organic material. It is a skin irritant. It is moderately toxic by ingestion. LD mg/kg. Do not heat or touch this powder. Rinse immediately if spilled on the skin. Wash your hands after using LAB PROBLEM You are part of a research team at JB Labs trying to develop instant hot and cold packs. You are experimenting with dissolving various salts in water. With your lab partner, devise a procedure to experimentally measure the enthalpy change for the dissolving of each salt in water. Use only about one gram of these chemicals. Record your procedure and results on your Student Lab Report. Do three trials and average your results. In your conclusions, calculate the enthalpy change using the formulas: q = mcδt and ΔH = q/# moles In addition to measuring the change in enthalpy, for each case: 1. Observe the dissolving process and determine if it is spontaneous. 2. Determine whether the entropy of the salt increases or decreases as it dissolves in water. Justify your answer. Do these two chemicals have potential for use in instant hot or cold packs? DISPOSAL The solutions may be rinsed down the drain. 2

3 CHAOS STUDENT LAB REPORT Name: Lab Partner: Lab Procedure: (Write as a series of steps) Lab Safety Issues: Lab Data and Observations: (Include a table when appropriate and provide qualitative observations) 3

4 Calculations: Results and Conclusions: (Discuss what happened and why. Try to explain results.) Questions for further study: (What questions do you still have or what would you like to know more about concerning this investigation?) 4

5 Graphing/Equations/Interpretations: Exothermic processes start with high potential energy and go to low potential energy. The loss of potential energy is usually released as heat. Exothermic reactions feel warm or hot. Endothermic reactions absorb heat, feel cool, and increase in potential energy. The dissolving of a salt can absorb or release energy. The following is the dissociation or dissolving process for sodium chloride. NaCl(s) + heat Na + (aq) + Cl - (aq), ΔH= +3.9kJ/mole (endothermic) The following potential energy diagram represents the above process: Potential Energy Na + (aq) + Cl - (aq) NaCl(s) Using the NaCl dissociation equation as a guide, write the dissociation equations for the two salts used in this lab. Remember to include the heat term on one side of the reaction. Label phases. (Use KNO 3 in equation #1) Reaction Now draw a sketch of the potential energy diagrams for each process on the graphs below, label the reaction coordinates and label each process as endothermic or exothermic: Equation #1 Equation #2 Questions: Reaction Coordinates Reaction Coordinates 5

6 Questions: 1. Is the salt dissociation process a physical or chemical change? 2. Does entropy increase or decrease for the dissolving process? 3. Using the relationship ΔG = ΔH - TΔS, and the information from your experimental calculations and question 2, is it possible for one or both of the dissolving processes to change from being spontaneous to non-spontaneous? Explain your answers. 6

7 GOING FURTHER J. Willard Gibbs discovered how to combine the two driving forces in nature: enthalpy and entropy. His equation is ΔG = ΔH - TΔS, where G is the free energy released or absorbed by the system. The sign of ΔG determines if the reaction is spontaneous. T is the temperature in Kelvin. ΔG is negative ΔG is positive ΔG is 0 Reaction is spontaneous Reaction is not spontaneous Reaction is in equilibrium In many cases, enthalpy (ΔH ), entropy (S ), and Gibbs free energy (ΔG ) values for elements, ions, and compounds have been determined. We can use these values to calculate enthalpy, entropy, or Gibbs free energy changes for reactions. To calculate these values use the following formulae: ΔG o reaction = ΣΔG o products ΣΔG o reactants ΔS= ΣS o products ΣS o reactants ΔH o = ΣΔH o products ΣΔH o reactants Table of Thermodynamic Values at Standard State Substance ΔH o (kj/mole) ΔG o (kj/mole) S o (J/mole K) Ca 2+ (aq) Cl - (aq) Cu 2+ (aq) H + (aq) K + (aq) Mg 2+ (aq) NO - 3 (aq) Cu(s) CaCl 2 (s) H 2 (g) H 2 O(l) H 2 O(g) KNO 3 (s) Mg(s) (Plambeck, 1995) Note: The units for all values are kj or J per mole. The o in ΔG o, ΔH o and S o means these values are for standard state conditions (298.15K, 1.00 atm, 1 M solutions). Use the Thermodynamic table to determine the ΔG o for the condensing of water. (Show your work) H 2 O(g) H 2 O(l) Is this a spontaneous process at 25 o C? ΔG o reaction =? 7

8 1. Use the table to calculate ΔH o for dissolving KNO 3 and CaCl 2 in water. Show your work: formula, substitution, answer. a. KNO 3 : b. CaCl 2 : 2. Label the above equations as endothermic or exothermic. Does this agree with your experimental results? Explain your answer: b. Calculate your percent error in each case 3. Calculate the standard entropy change (ΔS o ) associated with each equation from the table: a. KNO 3 b. CaCl 2 4. Do these results agree with your observations and beliefs about the dissolving process? Explain your answer. 5. What are standard conditions and how close were you to these conditions in your experiment? 6. Calculate the standard Gibbs free energy change (ΔG o ) for each process from the table: a. KNO 3 b. CaCl 2 7. What do the ΔG values tell you about the spontaneity of the dissolving processes? Do these results agree with your experimental results? Explain. 8