introduction - lee college · 1. define and give examples of rate equations, reaction order,...

Introduction

The main focus of any chemistry course is for the students to learn and understand the

nature laws and theories that pertain to matter and the changes that matter can undergo.

For students to achieve this they must learn to “think like a chemist”. When I say “think

like a chemist” I mean students must learn to solve complicated problems by using logic,

trial and error, intuition, and inferences. A chemist is used to being wrong or making

mistakes. The students must learn that mistakes are normal and that they can learn from

mistakes. They must learn to reevaluate their assumptions, inferences, and conclusions.

This is all based on the “scientific method” which involves using observations to

formulate laws and theories then testing these laws and theories using experiments and

again observations. The overall goal is to solve problems or puzzles, in other words to

“think critically”.

My chemistry courses are comprised of three components laboratory activities, lectures,

and exams. I have included an example of all three of these components with responses

from students in bold for the laboratory activity and lecture. For the exam I have included

the unit objective each problem is focused on.

The lab activity is inquiry based and consists of three basic components exploration,

invention, and application. The exploration is the activity such as a chemical or physical

reaction and observation. The observations may be qualitative or quantitative

(measurements). The inventions are designed for students to explain what their

observation might tell them about the activity or results. The application requires the

student to expand their results to other but similar processes.

The lectures are designed to introduce and explain chemical and physical concepts. I try

to make the lectures inquiry based also. Questions, often leading questions, are asked and

the student responses lead them to the concept being covered. The idea is to get the

student to develop solutions to particular problems and then generalize their solutions.

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This is to engage the student in active learning and develop a method of thinking like a

scientist.

The exams are designed to measure the student’s understanding of the objectives and

concepts they should master. The problems give the student information or data that they

must separate into what is needed and what is not needed in order to get the solution

asked to find. The must then use the laws or theories that will lead them to a correct

solution.

COURSE OUTLINE/SYLLABUS CHEM 1412 - GENERAL CHEMISTRY II

Instructor: Dr. Brian Hale Office: Science Building 142 Phone: 425-6330 (office) 837-8652 (home) Office Hours: MW 8:00 am - 9:00 am or by Appointment Division Chairman: Carolyn Foster 214 Science Building 425-6333 Required Materials: General Chemistry, 5th ed. Whitten, Davis, and Peck General Chemistry Laboratory Series, Wodeski and Zubkowski Safety glasses or goggles Calculator Compbook or other hard bound notebook

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Course Prerequisite: Math 330 or higher and Chemistry 1411 This course is intended for the student majoring in one to the physical sciences (chemistry, physics, etc.), engineering, or mathematics, who need general college level chemistry in preparation for higher level science courses in their respective curricula. Chemical concepts are emphasized from a mathematical approach. This course consists of three hours per week of lecture and three hours per week of lab. The material will be split into four units for testing purposes. Specific learning objectives for each unit are at the end of this syllabus. These objectives should be used to prepare for the unit tests and the comprehensive final exam. In addition, this course has a laboratory component of four hours per week. A single course grade will be awarded as described in the grading procedures. Please inform the instructor if you are a student with a disability and need any accommodations for this class. GENERAL COURSE OBJECTIVES Instructional Goals and Purposes: Lee College’s instructional goals include (1) creating an academic atmosphere in which students may develop their intellects and skills and (2) providing courses so students may receive a certificate/an associate degree or transfer to a senior institution that offers baccalaureate degrees. GENERAL COURSE OBJECTIVES Successful completion of this course will promote the general student learning outcomes listed below. The student will be able 1. To develop an understanding general laboratory techniques. 2. To develop an understanding of chemical thermodynamics including

the first and second laws, calorimetry, and bond energies. 3. To develop an understanding of chemical kinetics including rate laws,

collision theory, and mechanisms. 4. To develop an understanding of chemical equilibrium and the law a

mass action. 5. To develop an understanding of ionic equilibrium as it pertains to acids

and bases.

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6. To develop an understanding of buffer solutions and pH. 7. To develop an understanding of ionic equilibrium as it pertains to

solubility. 8. To develop an understanding of electrochemistry including voltaic

cells, standard potentials, and the Nernst equation. 9. To develop an understanding of nuclear chemistry including decay

reactions and half-lifes. 10. To develop an basic understanding of organic chemistry including

formulas, names, and properties. SPECIFICE COURES OBJECTIVES

Unit Objectives:

Unit 1 Chapter 15. 1. Define and give examples of the following terms: internal energy, endothermic,

exothermic, spontaneous, enthalpy, heat, standard enthalpy of formation, free energy, calorimeter, and Hess' Law.

2. Calculate enthalpy, entropy, and free energy using tables and/or formulas. 3. Use specific heat or heat capacity to calculate enthalpy and vice versa as in a

calorimetry experiment. Chapter 16. 1. Define and give examples of rate equations, reaction order, half-life, activation

energy, catalyst, reaction energy diagrams, and activated complex. 2. Find the rate law from simple concentration change data. 3. Give the rate law for a simple, multistep reaction. 4. Use the integrated rate equations. 5. Compare rate equations with proposed mechanisms.

Unit 2

Chapter 17.

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1. Define and give examples of the Law of Mass Action, equilibrium constants Kp

and Kc, homogeneous equilibrium, heterogeneous equilibrium, LeChatelier's Principle.

2. Write equilibrium constant expressions. 3. Calculate equilibrium constants. Convert Kc and Kp. 4. Predict reaction direction using Q and/or LeChatelier's Principle. 5. Calculate K from G. Chapter 18. 1. Define and give examples of the following: solute, solvent, molarity, molality,

normality, strong acid/base, weak acid/base, amphoteric, amphoprotic, and titration.

2. Discuss each acid/base theory: Arrhenius, Bronsted-Lowry, Lewis. Define acid/base for each theory.

3. Identify conjugate acid/base pairs.

Unit 3 Chapter 19. 1. Define and give examples of pH, pOH, buffers, Kw, Ka, Kb, common ion effect

and polyprotic acids. 2. Calculate pH, pOH given concentration of unionized acids and bases. Find the

percent ionization. 3. Work common ion problems for acid/base neutralizations and buffers. 4. Work common ion problems for salts.

Chapter 20. 1. Define and give examples of solubility, Ksp, complex ion formation. 2. Write Ksp expressions. 3. Calculate Ksp from solubility data. 4. Work common ion problems from Ksp.

Unit 4 Chapter 21. 1. Identify redox reactions. 2. Understand the activation series. 3. Assign oxidation numbers to atoms in molecules and ions. 4. Balance aqueous redox reactions. 5. Define and give examples of electrolytic cells, and Faraday's Law. 6. Diagram and label an electrochemical cell. Identify the half-reactions, cathode,

and anode.

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7. Calculate the standard potential for a cell. 8. Calculate the nonstandard potential using the Nernst equation. 9. Find G and K for a cell. Calculate time, mass, etc. for electrolysis reactions. Chapter 26. 1. Define and describe each of the types of radiation and subatomic particles. 2. Describe the importance of the neutron to proton ratio. 3. Balance nuclear equations. 4. Work half-life problems. Chapter 27: 1. Define and give examples of alkanes, cyclic, saturated, unsaturated aliphatic,

aromatic, olefin, alkyne primary, secondary, and tertiary positions. 2. Name alkanes, alkenes, alkynes and benzene derivatives. 3. Identify functional groups. 4. Discuss properties of functional groups. Chemistry 1412 Assigned Problems Chapter 15 6, 7, 9, 10, 13, 19, 21, 22, 23, 25, 27, 29, 31, 33, 37, 39, 41, 43, 51, 53, 55, 59, 63,

71, 73, 75, 79, 81, 83, 87, 89, 93, 97, 99 Chapter 16 7, 9, 13, 15, 19, 21, 23, 29, 31, 33, 35, 37, 39, 41, 45, 47, 49, 51, 53, 56, 57, 58,

59, 61, 63, 64, 65, 67 Chapter 17 6, 7, 11, 13, 15, 17, 19, 21, 27, 29, 31, 33, 35, 37, 47, 49, 59, 61, 63, 65, 71, 77, 79 Chapter 18 8, 13, 15, 19, 21, 25, 27, 31, 35, 37, 39, 41, 43, 47, 51, 59, 61, 63, 65, 69, 71, 73,

75, 77, 81, 83, 85, 88, 89, 92 Chapter 19 9, 13, 15, 17, 19, 21, 27, 31, 33, 37, 39, 43, 47, 52, 53, 54, 55 Chapter 20

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5, 7, 11, 13, 15, 17, 19, 21, 23, 27, 29, 31, 33 Chapter 21 9, 11, 17, 19, 21, 23, 27, 31, 35, 45, 47, 51, 53, 55, 57, 61, 69, 71, 73, 75, 77, 81,

85, 87, 107, 109 Chapter 26 25, 29, 33, 45, 47, 49, 51, 67 Chapter 27 23, 25, 27, 33, 37, 41, 43 Schedule: Week of Lecture Laboratory Aug. 29 Chapter 15 Exp. 9 Sept.5 Chapter 15 Exp. 6 Sept.12 Chapter 15, 16 Exp. 15 Sept.19 Chapter 16 Exp. 18 Sept.26 Chapter 16, Exam 1 Oct. 3 Chapter 17 Exp. 16 Oct. 10 Chapter 17,18 Exp. 17 Oct. 17 Chapter 18, Exam 2 Oct. 24 Chapter 19 Exp. 19 Oct. 31 Chapter 19, 20 Handout Nov. 7 Chapter 20 Handout Nov. 14 Chapter 20, Exam 3 Nov. 21 Chapter 21 Handout Nov. 28 Chapter 21, 26 Exp. 22 Dec. 5 Chapter 26, 27, Exam 4 Check -out Dec. 12 Final Exam Subject to change if needed Course Grading System:

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Total Labs 25% 25% 4 Exams 10% each 40% Homework Quizzes 10% 10% Final Exam 25% 25% 100% Make-up exams will only be given with prior approval of instructor. Students’ final grades are determined by the following grading scheme: 100-90 A 89-80 B 79-70 C 69-60 D 59 or below F EXPERIMENT 9: PERIODIC PROPERTIES OF ELEMENTS Name ___________________________________ Date ______________________________ Lab Partner ______________________________ Instructor __________________________ (If applicable) In this experiment you will study the properties of some elements to determine whether patterns of reactivity can be found in the Periodic Table. We want to determine whether all the elements in a Group or a Period have the same properties or whether there is a trend (increase or decrease) in a particular property within a Group or Period. The properties which you will explore are:

(1) the reactivity of elements with water

(2) reactions of solutions of cations of the alkaline earth metals (Group 2) with selected anions in solution.

Information which will help in identifying some reactions is given below. Learn this information before coming to lab.

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(1) Two colorless gases often produced in chemical reactions are hydrogen and oxygen. These gases can be distinguished by their properties. When ignited by a flame, hydrogen gas will burn with a characteristic “pop” or “squeak”. If the gas is oxygen, a glowing splint will burst into flame when placed in the gas.

(2) The presence of an acid or base in water can be tested by taking a drop of the

solution on a stirring rod and touching it to litmus paper. An acid produces H+ in solution and turns blue litmus red. A base produces OH- in solution and turns red litmus blue.

(3) A salt is an ionic compound containing a cation and an anion (e.g. NaCl and

CaCO3). All nitrate sales are soluble. All sodium salts are soluble. When two soluble salts of the general formula, MX and NY are mixed, a reaction will occur if either MY or NX is insoluble. For example, when solutions of NaCl and AgNO3 (both soluble) are mixed, a precipitate is formed because AgCl is insoluble.

EXPLORATION 1. Reactions of elements with water.

1. Fill a 400 mL beaker about three-fourths full with distilled water. Test the water with litmus paper to determine the presence of either acid or base. Fill a test tube with water and turn it over into the beaker of water so that the test tube remains completely filled with water. CAUTION: Be sure to handle Na metal with tongs. Hold a watch glass under the Na sample while you take it to your lab table. Obtain a small piece of Na metal with tongs and wrap it in a piece of aluminum foil which has been perforated. Using tongs, quickly place the wrapped sodium at the open end of the inverted test tube. You want to collect the gas that is produced in the test tube by displacement of water. When the reaction has ceased and the test tube is filled with gas, take the test tube out of the water, open end down. Keep the tube inverted and immediately test the gas by placing a burning wooden splint at the open end of the tube. Also test the water with litmus paper. Describe your observations on the next page.

Record your observations from 1. below. A reaction occurred that produced a gas. When ignited the gas made “pop” indicating it was H2(g). The water tested positive for a basic solution thus indicating the presence of NaOH.

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What two substances reacted? ____________Na & H2O __________ What two products were formed? _____________NaOH & H2________________ Write a balanced chemical equation for the reaction you have observed.

2Na(s) + 2H2O(l) -----> 2NaOH(aq) + H2g 2. Compare the reactivity of several elements with water. The elements which are

available are Na, K, Ca, Mg, Al and S. CAUTION: Be particularly careful in using the samples of Na and K. Use a 400 or 600 mL beaker for the Na and K reactions. Add a small sample of each element to distilled water in a beaker. Handle the samples with tongs or your spatula.

In table 1 below, record your observations and indicate whether a chemical reaction took place.

TABLE 1

Element

Observations

Was there evidence of a chemical reation?

Na

Reacted quickly dissolving the Na and forming a gas.

Yes. Na dissolved and a gas formed.

K

Reacted very rapidly dissolving the K and forming a gas.

Yes. K dissolved and a gas formed.

Ca

Reacted quickly but slower that Na forming a gas and dissolving the Ca.

Yes. Ca dissolved and a gas formed.

Mg

Reacted slowly forming a gas and dissolving some of the Mg.

Yes. Mg dissolved to some extent and a gas formed

Al

No reaction. No.

S

No reaction. No.

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INVENTION 1 3. You tested the reactivity of several elements in the same Group of the Periodic

Table. Describe any observed trend (increase or decrease) in the reactivity of the elements in a Group of the Periodic Table. Was the same trend observed for more than one Group? Explain.

K reacted faster that Na indicating that in group I the activity increase as you go down the periodic table. Both Na and K reacted to produce H2(g) and a basic solution indicating the react in a similar manner with water. 4. Describe any observed trend in the reactivity of the elements in a Period of the

Periodic Table. K was more reactive than Ca. Ca was more reactive that Mg. Mg was more reactive that Al which did not react as did S. This would indicate that the reactivity decreases as you go across the periodic table. 5. Predict what would be observed in each case if the elements, Ba Cs and Si were

added to water. Give the basis for your predictions. Cs would react very fast. Ba would react fairly fast. Si would not react. EXPLORATION 2 1. Reactions of Cations in Solution with Anions in Solution; Solubility

In this experiment you will test the solubilities of salts formed by the Group 2 cations, Mg2+, Ca2+, Ba2+, Sr2+, with various anions. Then you will study the results to see if there is any trend in solubility that corresponds to the positions of the cations in the Group.

Add about 1 mL (20 drops) of 0.1 M solution of the nitrate salts of barium, calcium, magnesium and strontium to separate small test tubes. To each test tube add 1 mL of 1 M H2SO4 and stir with your glass stirring rod. Rinse our stirring rod in a beaker of water between tests. Note whether a precipitate forms, as well as any distinct characteristics (color, etc.) of the precipitate in each test. Record your results in the first row of Table 2.

TABLE 2

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Sr2+ Mg2+ Ca2+ Ba2+ SO4

2- A white ppt formed No ppt formed No ppt formed A white ppt formed

CO32- A white ppt formed No ppt formed A white ppt formed A white ppt formed

C2O42- A white ppt formed A white ppt formed A white ppt formed A white ppt formed

CrO42- No ppt formed No ppt formed No ppt formed A yellow ppt formed

2. Write balanced chemical equations for any reactions which occurred when the SO4

2- ion was added to the solutions of the cations.

Sr2+ + SO4

2- ------> SrSO4(s) Mg2+ + SO4

2- ------> No Reaction Ca2+ + SO4

2- ------> No Reaction

Ba2+ + SO42- ------> BaSO4(s)

3. Repeat the experiment testing 1 mL of each alkaline earth cation with 1 mL of 1 M

Na2CO3. Record all observations in the Table.

4. Next, test for the solubilities of the oxalate, C2O42-, salts of the cations by adding 1

mL of 0.25 M (NH4)2C2O4 to 1 mL of each cation. Record all observations. 5. Test for the solubilities fo the chromate, CrO4

2-, salts of the cations by using 1 mL of 1 M K2CrO4 plus 1 mL of 1 M acetic acid with 1 mL of each cation. Record your results in the Table.

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INVENTION 3 6. Study Table 2 for an patterns of solubility. List the cations in order, starting with the

one that forms the largest number of soluble salts with the anions and ending with the cation that forms least number of soluble salts.

Ba2+, Sr2+, Ca2+, Mg2+ 7. Was there any trend in the ability of the cations of the alkaline earth metals to form

insoluble salts as you proceed up or down the Group? If so, describe this trend. As you go down the periodic table the more insoluble salts are formed. APPLICATION 1. In this part of the experiment you will test a solution containing a nitrate salt of one of

the alkaline earth cations to determine which cation the solution contains. This should require only two tests. Before you do any tests, study Table 2 and devise a procedure that will require only two tests to determine the identity of your unknown. Describe the procedure below.

Add SO4

2- to a solution containing an unknown cation. If a precipitate forms then the cation is either Sr2+ or Ba2+. Next add CrO4

2-. If a precipitate forms the unknown is Ba2+. If a precipitate does not form the unknown is Sr2+.

If a precipitate did not form when SO4

2- was added the unknown was either Mg2+ or Ca2+. If a precipitate forms when CO3

2- is added then the unknown is Ca2+. If a precipitate did not form the unknown is Mg2+.

2. Now test your unknown solution according to the procedure which you described

above. In the space below describe the tests you do and the observations you make.

My unknown number is ____#4____

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When SO4

2- was added a ppt formed. When CrO42- was added a ppt formed.

3. Draw a conclusion about what cation is present in your solution and state your reasoning. My unknown was Ba2+

Lecture Buffer Solutions What does the terms “buffer”, “buffer solution”, or “pH balance” mean to you? Answers will vary depending on their exposure to these terms. A buffer solution is a solution that does not change its pH with a small addition of an acid or base. An acidic buffer solution will contain a weak acid and the salt of that weak acid. A basic buffer solution will contain a weak base and the salt of that weak base. The action of a buffer solution is governed by the common ion effect, a special example of Le Chatelier’s principle. What is the Le Chatelier’s principle? Answer will vary but something along the line: When a stress is put on a system at equilibrium the system will shift to relieve the stress. Where might you find a buffer solution? Answers will vary: Some solutions or systems that are buffers are swimming pools, soft drinks, deodorant, shampoo, blood, etc.

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In order to understand what a buffer solution is and how it works we will construct a solution using the weak acid HF (Ka = 7.2 x 10-4) and its salt NaF. Consider a solution that contains 0.10 M HF and 0.30 M NaF. This solution contains two species and therefore will involve two reactions. What are these reactions? Students must recall NaF is a soluble salt and that all compounds that contain Na are soluble. NaF(aq) -------------------> Na+

(aq) + F-(aq)

Also the HF is a weak acid. HF(aq) <==========> H+

(aq) + F-(aq)

What are the initial and final concentrations for the first reaction? NaF(aq) -------------------> Na+

(aq) + F-(aq)

Initial 0.30 M 0.00 M 0.00 M After 0.00 M 0.30 M 0.30 M What are the initial and equilibrium concentrations for the second reaction (remember we now have another source of F -)? Students use what they learned in previous chapters. They must recall how they treated reactions that reach equilibrium. They must also recall how they treated weak acids. HF(aq) <==========> H+

(aq) + F-(aq)

Initial 0.10 M 0.00 M 0.30 M Equilibrium (0.10 – x)M x M (0.30 + x) M How would you solve for the equilibrium concentrations? Again students must use what they have already learned about equilibrium and weak acids. Ka = x(0.30 + x) (0.10 – x) Can we assume x is small? Yes making (0.30 + x) = 0.30 and (0.10 – x) = 0.10 or you can use the quadratic formula. Ka = x0.30 0.10 x = Ka0.10 0.30

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x = (7.2 x 10-4)(0.10) 0.30 x = 2.4 x 10-4 pH = -log(2.4 x 10-4) = 3.62 General Chemistry II Name____________________ Exam #1 1. [10 pts] The rate constant of the first-order reaction 2N2O(g) ----> 2N2(g) +

O2(g) is 0.38 s-1 at 1000. K and 0.87 s-1 1030. K. Calculate the activation energy of the reaction.

To develop an understanding of chemical kinetics including rate laws, collision theory, and mechanisms.

2. [10 pts] Calculate the reaction enthalpy for the formation of anhydrous aluminum

chloride 2 Al(s) + 3 Cl2(g) ----> 2 AlCl3(g)

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from the following data: 2 Al(s) + 6 HCl(aq) ----> 2 AlCl3(aq) + 3 H2(g) ΔHo = -1049 kJ HCl(g) ----> HCl(aq) ΔHo = -74.8 kJ H2(g) + Cl2(g) ----> 2 HCl(g) ΔHo

= -185 kJ AlCl3(g) ----> AlCl3(aq) ΔHo

= -323kJ

To develop an understanding of chemical thermodynamics including the first and second laws, calorimetry, and bond energies.

3. [10 pts] The following kinetic data were obtained for the reaction NO2(g) + O3(g) ----> NO3(g) + O2(g) Initial concentration, mol/L Initial rate, Experiment [NO2]0 [O3]0 mol/L.s 1 0.21 0.70 6.3 2 0.21 1.39 12.5 3 0.42 0.70 12.7 4 0.66 0.18 ? (a) Write the rate law for the reaction. (b) What is the order of the

reaction? (c) From the data, determine the value of the rate constant. (d) Use the data to predict the reaction rate for Experiment 4.

To develop an understanding general laboratory techniques.

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4. [10 pts] A certain first-order reaction has rate constant of 3.7 x 10-2 s-1. a. What is the half-life? b. If the initial concentration of the reactant is 2.33 x 10-2 mol/L, What will

be the concentration after 23 s?


5. [10 pts] The heat capacity of a certain calorimeter is 265.1 J/oC. When 25.0 mL of 0.700 M NaOH(aq) was mixed in that calorimeter with 25.0 mL of 0.700 M HCl(aq) , both initially at 20.0 oC, the temperature increased to 22.1 oC. Calculate the enthalpy of neutralization in kilojoules per mole of HCl. (SSOLUTION = 4.20 J/g OC and density of solution = 1.00 g/mL)

To develop an understanding general laboratory techniques. To develop an understanding of chemical thermodynamics

including the first and second laws, calorimetry, and bond energies.

6. [10 pts] The following is a proposed mechanism:

Cl2 <====> 2Cl (fast, equilibrium) CHCl3 + Cl ----> CCl3 + HCl (slow) CCl3 + Cl ----> CCl4 (fast)

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The overall reaction is; Cl2 + CHCl3 ----> CCl4 + HCl determine the rate-law expression for this reaction.


7. [10 pts] The rate constant for the following reaction at various temperatures. Find the activation energy of this reaction using the graphical method.

N2O5 ----> 2NO2 + 1/2 O2 Temperature (K) k (1/s) 1/T(K) ln k 195 1. 08 x 109 230 2.95 x 109 260 5.42 x 109 298 12.0 x 109 369 35.5 X 109

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To develop an understanding general laboratory techniques. To develop an understanding of chemical kinetics including rate laws, collision theory, and mechanisms.

8. [10 pts] Calculate the free energy change for the combustion of methane from the standard free energies of formation.

ΔGo

f (CH4(g)) = -50.7 kJ/mol ΔGo

f (H2O(g)) = -228.6 kJ/mol ΔGo

f (CO2(g)) = -394.4 kJ/mol


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9. [10 pts] Determine whether the following reaction will occur spontaneous at 25 oC

as written. 2SO2(g) + O2(g) ----> 2SO3(g) SO2(g) ΔHof = -297 kJ/mole So = 248 J/mole K O2(g) ΔHof = 0 kJ/mole So = 205.0 J/mole K SO3(g) ΔHof = -396 kJ/mole So = 257 J/mole K


10. [10 pts] The following reaction was studied, and the following data were obtained

at a particular temperature: 2N2O(g) ----> 2N2(g) + O2(g) Time (s) [N2O] (Mol/L) ln[N2O] 15.0 0.0835 30.0 0.0680 80.0 0.0350 120.0 0.0220 Using graphical techniques, determine if this reaction is first order in [N2O],

the rate constant and the value of [N2O] after 60 s.

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To develop an understanding general laboratory techniques. To develop an understanding of chemical kinetics including rate laws, collision theory, and mechanisms.

As a member of the FLC I have enjoyed the sharing of ideas about teaching and learning

in general as well as the different methodologies use by my colleagues. It has forced me

to evaluate my teaching methods as well as what I consider important for the student to

learn and understand. Although this does not effect the content of my courses or more

importantly the concepts I expect my students to understand I believe I now articulate

these concepts more meaningful manner. For example, along with the laws of nature and

the theories that model nature I expect my students to be able to use these laws and

theories to solve problems. I also expect my students to observe nature and the process

that occur in nature. From their observations using the laws and theories they should be

able to make certain predictions. I call this “thinking like a scientist.” This is the same

logic that is used in critical thinking only the terminology is different. This semester I

have incorporated this terminology along with the “scientific terminology” in my courses.

My thought is that students will be hearing about critical thinking (and that terminology)

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in other courses and this will reinforce the points I’m trying to make. I also hope they see

the connectivity between science and other disciplines. I believe this has enabled me to

better articulate the ideas and concepts I am teaching.

introduction - lee college · 1. define and give examples of rate equations, reaction order,...

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