The baseline data was acquired during the last semester of the standard approach (Y1) the last two years (Y2 and Y3) were based on students taking the atoms-first curriculum. Student data was evaluated during the second semester portion of the course in three academic years (Y1, Y2, and Y3). The experiment was implemented concurrently with the curricular change from a classical approach with one textbook and homework system to an atoms-first approach which used a different textbook and homework system. The laboratory is broken into sections of a maximum of 60 students each and are not aligned with a specific lecture section, with the individual lab sections under the direction of a faculty instructor or graduate teaching assistant. The course has multiple lecture sections with approximately 100 students each, with different lecturers for the individual sections. The experiment was developed and implemented as part of a second-semester STEM majors General Chemistry course at a public university in rural Tennessee. Ice from a laboratory ice machine was used to cool the samples. pH buffers (4 and 7) were used to calibrate pH probes.
Sodium hydroxide solutions of equal concentration were provided to students to neutralize the acetic acid solutions so the spent solutions could be disposed of safely down the drains in accordance with local environmental regulations. The analyte, 1 M acetic acid, was prepared by dilution of glacial acetic acid. The experiment was developed and implemented to answer the following research question: Can the incorporation of this activity improve student success rates on typical calculations presented as part of summative assessments at the conclusion of a second semester course in General Chemistry, specifically acid-base equilibrium calculations? The procedure and material requirements are light enough to allow this experiment to be performed in high schools (AP Chemistry curricula) or at colleges/universities where resources are limited. In the interest of bringing this concept to a broader group of students, an alternate methodology was developed to perform the analysis based on just the natural dissociation of the acid as determined by pH and temperature measurement. 5, 6, 7, 8, 9, 10, 11, 12, 13 However, many of these analyses require instrumentation that may not be widely available to all laboratories, or require more sophisticated procedures that are unsuitable for an introductory audience. Laboratory investigations of this linkage with different parameters have been published using a variety of techniques such as titration, flame emission spectroscopy, emission & absorption spectroscopy, and calorimetry. This relationship allows the determination of thermodynamic quantities from the value of the equilibrium constant over a range of temperatures. The ∆G° itself can be calculated from two other properties, enthalpy (∆H°) and entropy (∆S°), which are related through the equation:Įquations (1) and (2) can be combined to give the van’t Hoff equation:
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1 A well-known fundamental equation relates the Gibbs free energy (∆G°) and equilibrium constant ( K) via the gas constant (R) and the absolute temperature (T): 2, 3, 4 The concepts of equilibrium, thermodynamics, and the connection between the two are cornerstone concepts covered in a typical general chemistry curriculum. Calculated values of ∆G° 298 are consistent with literature values, indicating that the experiment is suitably robust to be performed successfully by students of a wide range of skill levels. This experiment outlines the simple measurement of the equilibrium constant and temperature, with aspects of graphical analysis, allowing students to link the concepts of equilibrium and thermodynamics conceptually and mathematically. Measuring the pH of a weak acid solution while varying the temperature allows for this analysis to be conducted simply, making it more accessible to broader range of academic laboratories.
Previously published procedures require sophisticated technology or methodology, such as simultaneously measuring temperature and absorbance using a spectrophotometer, which may be unavailable to small and/or rural colleges and universities. The thermodynamic properties ∆G°, ∆H°, and ∆S° are difficult to measure directly in a laboratory setting, but can be determined by monitoring the temperature dependence of the equilibrium constant, K. The concepts of equilibrium and thermodynamics are among the most important topics covered in a general chemistry course.
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