CHEM 101 EXAM 2 IFORMATION SPRING 2016 GENERAL INFORMATION Exam Time, Place, Format, and Rules The exam will take place on Wednesday, 05/25, in room SCI 314 during the regular class time. (We will go over a couple of topics from Chapter 11 and will do the lab check out before the exam.) The exam will have two parts: multiple choice questions and open-end questions. You will need to bring a Scantron form No. 882-E. Partial credit will be given for open-end questions only. Open-end questions will be of two categories: problems and short essay questions. They will be similar to those from your homework assignments, worksheets, and quizzes. You are expected to give a logical, well organized, and complete solution for each problem. The answer only responses to problems will not receive much credit. You do not need to show your work for conversions involving metric prefixes. Your responses to essay question should be short, but complete; they should demonstrate your knowledge of specific information as it was presented and discussed in lecture and/or the textbook. Your random personal thoughts on a subject will generally get little credit. During the exam, you are not allowed to use any dictionaries and you are not allowed to use any electronic devices except for your personal electronic calculator. You cannot share a calculator with a classmate or use your cell phone in place of a calculator. Your cell phone cannot be in your hands at any time during the exam. General Tips on Preparation 1. Review your own lecture notes or, in case your are not a good note taker or missed a lecture, borrow them from a classmate. On the exam you should expect questions on any material covered in lecture and lab including instructor’s demonstrations. You will not be tested on any topics that are covered in the current chapters of the textbook, but have not been discussed in lecture or lab. 2. Complete all the worksheets from lecture and lab (even if you have not submitted them for grading) and review your graded quizzes. Problems on the exam will be similar to problems from the lecture and lab worksheets, homework assignments, and quizzes. You should be able to redo all the problems on MasteringChemistry for practice even if you already submitted the answers. 3. Review the exam outline below and identify areas that you need to work on. Review the related textbook material. Solve a few problems from each chapter that are similar to the problems we solved in class or to the problems you solved on MasteringChemistry. A problem is considered to be solved if your answer matches closely the answer from the textbook or a worksheet. The answers to the textbook in-chapter problems are in Appendix IV; the answers to the textbook end-of-chapter odd-numbered problems are in Appendix III. Textbook Chapters (look through the exam outline for specific sections) Chapter 5 (summary on pages 235 – 236); Chapter 6 (summary on pages 284 – 286); Chapter 7 (summary on pages 328 – 329); Chapter 8 (summary on pages 372 – 373); Chapter 9 (summary on pages 416 – 417); Chapter 10 (summary on pages 472 – 473); Chapter 11 (summary on pages 533 - 534); Chapter 18: section 18.2 only (balancing redox reactions). Look through/study the sections to avoid surprises on the day of the exam! Solve as many recommended problems from the textbook as possible! Laboratory Work to Know Exp. 6 “Analysis of Al/Zn alloy”; Exp. 7 “Enthalpy of Reaction”; Exp. 10 “Periodic Properties”; Exp. 9/16 “Conductometric Titration / Gravimetric Analysis”; and Exp. 18 “Redox Titrations” EXAM OUTLINE GASES Textbook sections: 5.1 – 5.7 (concepts and calculations); 5.8 – 5.10 (no calculations except for Graham’s Law); 11.5 (vapor pressure; omit the Clausius-Clapeyron). How do properties of gases differ from properties of liquids and solids on macroscopic level (appearance, density, compressibility) and microscopic level (distances between molecules and molecular motion)? Be familiar with basic ideas of kinetic molecular theory. Be able to distinguish between ideal gas and real gas. Know how to interpret temperature and pressure of gasses in terms molecular motion. Remember that T is directly proportional to the average kinetic energy (K.E.avg) of gas molecules, not to the average velocity 2 mvavg (vavg) ! T ~ K.E.avg = , where m is the mass of one molecule of gas. At the same temperature larger 2 molecules on average move slower than smaller ones. Pressure (force per unit area). Common units pressure that you should know how to use: atm, mmHg, and torr. (The SI unit of pressure is pascal that is abbreviated as Pa, but you are not responsible for remembering its relation to other common units.) Gas laws: mathematical relationships between macroscopic parameters (amount of substance, mol; temperature, K; pressure; volume) describing the ideal gas. You are encouraged, but do not have to remember each law by name as long as you can write the appropriate relationship when solving gas law problems. Remember how the ideal gas law formula PV=nRT allows you to deduce equations such as P1×V1 = P2×V2 (n=const, T=const). atm ⋅ L Remember the value of R with 4 significant digits and the appropriate units: R = 0.08206 . As long mol ⋅ K as you remember that 1 atm = 760 torr or 760 mmHg (exactly), you can multiply the value of R given above torr torr ⋅ L by 760 to obtain R = 62.36 . atm mol ⋅ K Be able to distinguish between densities and molar volumes of liquids, solids and gases. Equal numbers of molecules of gases occupy the same volume under the same conditions of temperature and pressure (the Avogadro’s law). The Avogadro’s law is a gas law and it is not applicable to liquids and solids. The molar volume of 22.4 L/mol at STP (0°C and 1 atm) is a good approximation for all gases, but cannot be applied to liquids and solids. The molar volume of water at 25ºC and 1 atm is about 0.018 L/mol (the volume of the ideal gas under the same conditions is 24.5 L/mol). Know how the densities of gases depend on temperature, pressure and identity of a particular gas (its molar mass). Know how to use the ideal gas law (PV=nRT) to convert between the amount of gas in moles and the volume (an essential skill in solving stoichiometry problems where a reactant or a product is a gas). In solving stoichiometry problems always rely on a properly written and properly balanced chemical equation. (Do not assume that all substances react in one-to-one mole ratio. Do not simply equate the moles of gas with the moles of a nongaseous reactant or product: use the mole ratio from the balanced chemical equation!!!) If three out of the four parameters describing the ideal gas (P, V, n, T) are given, know how to use the ideal gas law formula (PV=nRT) to find the forth. Be familiar with the concept of partial pressure of gases in a mixture and the law of summation of partial pressures of gases (Dalton’s) law. Know how to use the Dalton’s law to calculate the partial pressure of the gas collected over water if the vapor pressure of water is given in the problem or a reference table “Vapor Pressure of Water vs. Temperature” is available). Liquid-vapor equilibrium: vapor pressure depends on temperature, but not on the volume of the system. What is common and what is different between boiling and vaporization? Vapor pressure vs. temperature tables (a blue laminated handout for water) and graphs. Boiling point as a function of pressure. Normal boiling point. THERMOCHEMISTRY Textbook sections: 6.2 – 6.9 (no PV work calculations such as in Example 6.4); 9.10 (bond energies); 11.7 (∆Hfus and ∆Hvap). Have a basic understanding of the subject and significance of thermochemistry and chemical thermodynamics. Have a basic understanding of the concepts of energy, internal energy, heat and work. J (a derived SI unit) and cal as units of energy, work, and heat. Know the new definition of calorie (1 cal is exactly 4.184 J) and the old one (the amount of energy required to increase the temperature of one gram of liquid water by one degree Celsius). Distinguish between heat at constant volume (qv = ΔE) and heat at constant pressure (qp = ΔH). Concept of enthalpy. Why do chemists prefer to use enthalpy rather that internal energy? Know how calorimetry can be used to determine specific heats of substances and heats of reactions. What are the advantages and limitations of a coffee-cup calorimeter? What are other types calorimeters that you know? Be able to compare and contrast the terms in each of the following pairs: system and surroundings, temperature and heat, heat and specific heat, heat capacity and molar heat capacity, specific heat and molar heat capacity, heat and work, internal energy and enthalpy, chemical equation and thermochemical equation, endothermic and exothermic. Be able to understand and use the formula: q = c×m×ΔT. You are recommended to memorize the specific heat of H2O(l): 1.00 cal/(g∙°C) and 4.184 J/(g∙°C). Know what thermochemical equations and heats of reactions are. The ΔH that accompanies a thermochemical equations has a unit kJ (per reaction as described by the equation with the coefficients representing moles of each reactant and product), not kJ per mole of a specific substance. Be able to give a written statement of Hess’s law of heat summation and know how to apply in solving problems.