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INTRODUCTION During a chemical reaction, chemical bonds are typically broken and formed. For example, examine the reaction between hydrogen and oxygen to form water: 2H₂ + O₂ → 2H₂O (Eqn. 2-1) For this reaction to occur, one O-O bond and two H-H bonds are broken while four H-O bonds are formed. It takes energy to break the chemical bonds of the H₂ and O₂ reactants and, conversely, energy is released when the H-O bonds are formed. This particular reaction is highly exothermic, that is, it gives off energy. Therefore, another product of this reaction, which is not included in the equation above, is energy in the form of heat. 2H₂ + O₂ → 2H₂O + Heat (Eqn. 2-2) In addition to the laws of conservation of mass and charge, chemical reactions such as that shown above are also subject to the law of conservation of energy. One way to state this law is that the energy released (or consumed) into the surroundings during a chemical reaction is equal to the difference in the energy spent breaking chemical bonds and the energy given off forming chemical bonds during the course of the reaction. The energy exchanged with the surroundings during a chemical reaction can take many forms; in practice, such energy is most commonly observed as heat. Reactions releasing energy in the form of heat are called exothermic and those absorbing energy in the form of heat are called endothermic. Exothermic reactions result when the energy released by the formation of chemical bonds is greater than the energy needed to break the bonds of the reactants. The opposite is true of endothermic reactions. The heat energy released by a reaction under constant pressure conditions is called enthalpy. The heat transferred to the surroundings during the course of any reaction is called the enthalpy of reaction and is denoted by AH₁. By convention, AH, is positive for endothermic reactions and negative for exothermic reactions. AH, values are typically molar quantities, in other words, they tell you how much energy is released or consumed per mole of product formed. Some reactions are so common that they have their own symbols for enthalpy.

Table 2-1: Some symbols for Enthalpy Name Enthalpy of combustion Enthalpy of formation Enthalpy of neutralization Enthalpy of solution Symbol AHcomb AH SAFETY AHneut . AHsol In this experiment, you will use a Styrofoam cup calorimeter to measure the heat released by three reactions: Reaction #1: Dissolving solid sodium hydroxide in water: • Description Enthalpy of combustion of a substance in O₂ Enthalpy of formation of one mole of a substance from its constituent elements. NaOH(s) → NaOH(aq) (Eqn. 2-3) Reaction #2: The reaction between aqueous sodium hydroxide and aqueous hydrochloric acid: Enthalpy of reaction between one equivalent of acid with one equivalent of base. Enthalpy of dissolution of one mole of solute. NaOH(aq) + HCl(aq) → H₂O(l) + NaCl(aq) (Eqn. 2-4) Reaction #3: The reaction between solid sodium hydroxide and aqueous hydrochloric acid: NaOH(s) + HCl(aq) → H₂0 (1) + NaCl(aq) (Eqn. 2-5) Notice that reaction #3 is the same as the sum of reactions #1 and #2. Applying Hess's law, AH, for #3 should simply be the sum of AH, for reaction #1 and AH, for reaction #2. For purposes of this experiment, you may assume that the heat lost to the calorimeter and the surrounding air is negligible. (Heat lost to either of these is a fairly constant factor in each part of the experiment, and has little effect on the final results.) NaOH (Sodium Hydroxide) is an irritant. Avoid contact with the eyes. HCI (Hydrochloric Acid) is corrosive and will cause burns. Handle carefully. Avoid skin contact and inhalation of vapors. PROCEDURE Reaction #1: NaOH(s) → NaOH(aq) 1. Prepare for data collection by connecting the Vernier temperature probe to the LabQuest device. The LabQuest device should automatically detect the temperature probe and open the correct data acquisition window. Make sure the vertical axis is Temperature (°C) and the

horizontal axis is Time (s). Adjust the time acquisition to 400s by clicking Duration on the meter screen. Click "OK". 2. Weigh the empty "double" Styrofoam cup. Measure out 100.0 mL of water into this Styrofoam cup. 3. Place the filled "double" Styrofoam cup into a 250 ml beaker to stabilize. 4. Lower the temperature probe into the solution. The setup is shown below in Figure 2-1. 5. Place the lid on the bottom "double" Styrofoam cups. Have the stirrer ready. 6. Click the green arrow (bottom left corner) to begin data collection and obtain initial temperature, T₁. Collect temperature readings for several seconds (min. 10 sec.). It may take several seconds for the temperature probe to equilibrate to the temperature of the solution. Immediately proceed with Step 7. iting rod thermometer styrofoam cup with the rim trimmed off 2 stacked styrofoam cups water and dissolved. reactants Beaker Figure 2-1: Setup of "double" Styrofoam cup calorimeter with temperature probe Weigh out about 2 g of solid sodium hydroxide and record the mass to the nearest 0.001 g.. Since sodium hydroxide readily picks up moisture from the air, it is necessary to weigh it and proceed to the next step without delay. AUTION: Handle the sodium hydroxide with care. Be sure to replace the lid immediately as aOH will pick up water from the air. Clean up any excess reagent. Open the lid of your calorimeter and add the solid sodium hydroxide to the Styrofoam cup (quickly!) and cover with lid. Avoid dumping the solid onto the probe. Using a stirring rod

(or straw), stir continuously for the remainder of the 400 seconds or until the temperature maximizes and starts to level out/very slowly drops. As soon as the temperature has begun to drop for 30 seconds, you may terminate the trial by clicking the red square button (bottom left). 9. Disassemble the Styrofoam calorimeter. Take the bottom "double" Styrofoam cup that still contains the water with the dissolved NaOH out of the glass beaker and weigh this filled cup. Record its mass. 10. Examine the initial readings in the table window to determine the initial temperature, T₁. To determine the final temperature, T₂, click the Analyze → Statistics → Temperature. The maximum temperature (T₂) is listed in the statistics box to the right of the graph (T₁ is also listed). Record T: and T₂ in your lab notebook. 11. Rinse and dry the temperature probe, Styrofoam cups, and stirring rod. Dispose of the solution as directed. This solution is basic (0.5 M NaOH, pH > 13) and needs to be handled with care. Reaction #2: NaOH(aq) + HCl(aq) →H₂0 (1) + NaCl(aq) 12. Repeat Steps 2-11, initially measuring out 50.0 mL of 1.0 M HCI (instead of water) into the Styrofoam calorimeter. In Step 8, instead of solid NaOH, measure 50.0 mL of 1.0 M NaOH. solution into a graduated cylinder. After T₁ has been determined for the 1.0 M HCI, add the 1.0 M NaOH solution to the calorimeter and record the resulting temperature changes as before. Remember to record the masses (like in Step 9) and the exact concentrations. Reaction #3: NaOH(s) + HCl(aq) →H₂0(1) + NaCl(aq) 13. Repeat Steps 2-11 using 100.0 mL of 0.50 M HCI (instead of water) and 2.0 g of solid NaOH. Record the resulting temperature changes as before. Again, remember to record the masses (like in Step 9). Clean up o Rinse and dry your Styrofoam cups. o Wipe down your work space. o Make sure shared spaces are clean and clear of debris. o Have your instructor initial your worksheet before leaving.

(include UNITS for each entry) Mass of Dry Calorimeter Mass of Solid NaOH Total Mass of Filled Calorimeter Mass of Solution Initial Temperature, Ti Final Temperature, T Change in Temperature, AT Enthalpy of Reaction, Moles of NaOH Molar Enthalpy of Reaction, AH, Reaction #3 Reaction #1 7,2669 7.2669 1.855 g 1,9999 N/A 104-9959 57.7309 106.5169 106.8069 108.83369 108.181 g | 19.90 19.2 C 29.0c° 9.8 19.9 24.3 4.4 Reaction #2 7.7665 26.2c 6.3