If you look at any chemistry textbook, you will see Lewis structures introduced long before electronic and molecular geometries. This makes sense since you need Lewis structures to determine molecular geometry. Unfortunately, research has shown that students often do not recognize that the purpose of drawing Lewis structures is not to create the structure itself but to use it as tool to understand the properties of the molecule (Cooper, Grove, Underwood & Klymkowsky, 2010). 

While chatting with my colleague across the hall about teaching Lewis structures and geometries, he showed me an intriguing activity he uses to introduce the VSEPR model before ever having students draw a Lewis structure. I had my doubts about this model, to which my colleague replied, “I’m not saying this is the best way to teach this topic. Actually, I guess I am saying it is the best way.” As if that argument was not convincing enough, I did a little digging into the issue. I found that research (cited above) supports the practice of connecting 3-dimensional representations to 2-dimensional Lewis structures and not the other way around. I tried this approach and was amazed with how quickly my students were able to grasp the complex concepts of the VSEPR model, Lewis structures and polarity.  

I started by reviewing basic atomic structure with my students. Atoms are made of protons, neutrons and electrons with electrons on the outside. I then asked students what the natural tendency of electrons is when near each other. They quickly replied, “they repel each other.” With that idea in mind, I issued a challenge to my students: “construct 3-dimensional models of molecules with 2, 3, 4, 5 and 6 atoms around a central atom with styrofoam spheres. Keep in mind that electrons repel each other and therefore will be most stable in an arrangement that minimizes repulsions.” I tried to avoid anthropomorphic language like “the electrons do not want to be near each other” because it is just as easy and more chemically correct to relate the phenomenon to stability. Students checked their geometries with me before filling out this worksheet (you may download below). Students gave each geometry their own name to help them remember what it looks like and I gave them the actual names later.

The first two geometries were pretty easy for students to come up with because they are planar. When they got to “central atom + 3” (tetrahedral), every group first constructed this planar geometry:

screen_shot_2016-04-12_at_7.23.51_pm.png

and were flabbergasted when I told them that it was incorrect. In every class, one group would eventually create a 3-dimensional geometry and the rest of the class would follow. As groups finished their worksheet, I asked them to rebuild one of their configurations so I could use it as my example piece. By the time every group finished, I had a full set of molecular models.

geometric-approach-lewis-structures.png

We then went through each model as a class, identified the bond angles and gave each geometry a real name. I also showed students how to draw 3-dimensional structures with wedge-dash notation (essentially a Lewis structure). The amount of information my students were able to recall the next day about their geometries was pleasantly surprising.

I then made the transition from electronic geometry to molecular geometry through water. I wrote the formula for water on the board and asked students to draw what they thought the molecule looked like. Of course they all drew this:

screen_shot_2016-04-12_at_7.27.15_pm.png

What they drew was completely consistent with their current model so students were surprised to see the actual geometry of water is bent (I showed students a styrofoam sphere model). With the backbone of the VSEPR model established, students were able to determine there must be something repelling the hydrogen atoms to create the bent geometry. At this point, I had students construct a Lewis structure of water with me to introduce the concept of “lone pairs.” I gave students very few “rules” for drawing Lewis structures except that the valence electrons are the only ones that participate in bonding and it is energetically favorable for electrons of opposite spins to pair together (electrons are rarely found unpaired). We drew the Lewis structure with a bent geometry to make sure it represented the actual water molecule.

screen_shot_2016-04-12_at_7.27.39_pm.png

I then compared the geometry of water to the tetrahedral model my students had made the day before. Instead of 4 atoms attached the central atom, 2 atoms and 2 lone pairs were attached to the central atom. Students concluded that both models had 4 “electron domains” attached to the central atom. With the idea of “electron domains” in mind, I set my students off on the “Molecular Shapes” PhET. Students made a table of geometries by replacing bonds with lone pairs (keeping the number of electron domains the same). The table looked like this:

This table (you can download a copy below) became the basis for all future Lewis diagrams. After visualizing the possible 3-dimensional structures of molecules, my students were able to easily construct Lewis structures with correct geometries.

My students now have an understanding of the purpose of Lewis structures and do not see molecular geometries as something to simply be memorized. The two tools are one in the same. Now that we are learning about polarity, it is incredibly simple for my students to sketch the geometry of a molecule and determine if it is polar or non-polar and the implications of such. 

In a dramatic movie trailer voice: “The Boiling Point. Gone without a trace. Or were they? The scene… a mystery. Had they disappeared? Been broken up into unrecognizable pieces? Can our hero find the answer? Or will it be too late?”

Sound like a Hollywood-style film? (Well, maybe not.) It’s the drama of chemistry, with the learner as the star. The April 2016 issue of the Journal of Chemical Education shares a new show that blends entertainment and learning.

The Boiling Point

Kerby et al. describe their theater-based demonstration show The Boiling Point in the article Demonstration Show That Promotes and Assesses Conceptual Understanding Using the Structure of Drama (full article available to JCE subscribers). Although the article title presents it as a demo show, it’s probably not what you expect. There are only three demonstrations, presented near the middle of the show. They are offered without explanation—rather, they are meant to provide information needed to answer the question, “What happens to water when it boils?” Audience members are active learners and part of the show, with a goal of answering the question based on what they observe during the experience.  The show uses “a design where engagement is driven by the audience’s desire to learn the concept, rather than the traditional strategy of piquing interest through spectacle alone.”

The show is aimed at ages 5 through 11 and requires less material than some of the group’s previous shows. It could be an opportunity for a school’s science department or chemistry club to connect with the drama department, with a view toward presenting an outreach show at an elementary school or within the community, perhaps in a library setting. The authors have made the script available online at the Fusion Science Theater (FST) website. The script can be freely used, as long as the authors and FST are acknowledged. There is also a video of a performance of the show at the site.

The Science of Measurement

The second piece that caught my eye in this issue was another performance, but this time with measurement as the main event. Amy Charles shares her review of a BBC series on DVD, in What Are the Units, and Why? Review of The Science of Measurement: Taking the Measure of the World (available to non-subscribers). Just reading the review was entertaining, but it also spurred my interest in the series itself.

It takes the viewer on a “lightning tour” of how we “define amount of what you’re looking at.” It not only brings us to the SI units we want students to know (and remember to include as labels), but also relates the difficulties associated with some of the units, and the future of measurement as researchers continue to strive for increased precision. One of the difficulties highlighted by the series is the weight loss of “Le Grand K,” the piece of metal meant to set the standard for the kilogram. The BBC website has a short clip from one of the episodes, showing a visit to the metal that’s locked up tight, including scientists getting photos of themselves posing in front of the opened vault, with the celebrity kilogram.

What are your thoughts on these two shows? What are some ways you’ve chosen to bring drama to your classroom?

For more from the April 2016 issue, see Mary Saecker’s JCE 93.04 April 2016 Issue Highlights.

Chemists Celebrate Earth Day

The April 2016 issue of the Journal of Chemical Education is now available online to subscribers. In honor of Earth Day 2016, the April issue includes a variety of content that provides ideas and suggestions for bringing environmental chemistry to students on the topics of: water quality; climate science and greenhouse gases; atmospheric chemistry; sustainability, green chemistry, and environmental awareness; and energy storage technology. Also in the issue are articles on: outreach and public understanding; teaching organic chemistry; physical chemistry; exploring biochemistry with proteins; research experiences in the laboratory; educational resources.

Environmental Chemistry

Cover: Earth Day

Chemists Celebrate Earth Day is an ACS annual event that seeks to bring international focus to environmental topics, such as clean air, water, and energy, to illustrate the positive role chemistry plays in understanding and preserving the Earth. Environmental chemistry engages students of all levels in learning new chemistry concepts by relating chemistry to their daily lives as well as heightening awareness of environmental issues. (Image of the Earth straddling the limb of the Moon taken by the Lunar Reconnaissance Orbiter Camera is used with permission: NASA/Goddard Space Flight Center/Arizona State University.)

Water Quality

Editorial ~ Jerald L. Schnoor, professor of engineering at The University of Iowa and former Editor-in-Chief of Environmental Science & Technology, strongly encourages the recognition that aging water infrastructure, particularly lead pipes, solder, and faucets, represents a community health hazard of enduring significance.

A Multidisciplinary Science Summer Camp for Students with Emphasis on Environmental and Analytical Chemistry ~ Gunnar Schwarz, Wolfgang Frenzel, Wolfgang M. Richter, Lothar Täuscher, and Georg Kubsch

Development and Implementation of a Simple, Engaging Acid Rain Neutralization Experiment and Corresponding Animated Instructional Video for Introductory Chemistry Students ~Danielle Rand, Craig J. Yennie, Patrick Lynch, Gregory Lowry, James Budarz, Wenlei Zhu, and Li-Qiong Wang

Introducing Students to Gas Chromatography–Mass Spectrometry Analysis and Determination of Kerosene Components in a Complex Mixture ~ Giselle Mae M. Pacot, Lyn May Lee, Sung-Tong Chin, and Philip J. Marriott

Determination of Arsenic in Sinus Wash and Tap Water by Inductively Coupled Plasma–Mass Spectrometry ~ Anna M. Donnell, Keaton Nahan, Dawone Holloway, and Anne P. Vonderheide

Climate Science & Greenhouse Gases

Designing and Demonstrating a Master Student Project To Explore Carbon Dioxide Capture Technology ~ Florine Asherman, Gilles Cabot, Cyril Crua, Lionel Estel, Charlotte Gagnepain, Thibault Lecerf, Alain Ledoux, Sébastien Leveneur, Marie Lucereau, Sarah Maucorps, Mélanie Ragot, Julie Syrykh, and Manon Vige

CO₂ Investigations: An Open Inquiry Experiment for General Chemistry ~ Roland P. Stout

Ocean Acidification: Investigation and Presentation of the Effects of Elevated Carbon Dioxide Levels on Seawater Chemistry and Calcareous Organisms ~ Jeffrey M. Buth

Updating a Student-Generated Ice-Core Data Plot Exercise for Courses Investigating Climate Change Topics ~ Edward Maslowsky Jr.

Atmospheric Chemistry

Investigating Freezing Point Depression and Cirrus Cloud Nucleation Mechanisms Using a Differential Scanning Calorimeter ~ Kentaro Y. Bodzewski, Ryan L. Caylor, Ashley M. Comstock, Austin T. Hadley, Felisha M. Imholt, Kory D. Kirwan, Kira S. Oyama, and Matthew E. Wise

Measuring Heterogeneous Reaction Rates with ATR-FTIR Spectroscopy To Evaluate Chemical Fates in an Atmospheric Environment: A Physical Chemistry and Environmental Chemistry Laboratory Experiment ~ Jason E. Roberts, Guang Zeng, Marta K Maron, Mindy Mach, Iman Dwebi, and Yong Liu

Sustainability, Green Chemistry, & Environmental Awareness

Incorporating Sustainability and Life Cycle Assessment into First-Year Inorganic Chemistry Major Laboratories ~ Marta Guron, Jared J. Paul, and Margaret H. Roeder

Design of a Dynamic Undergraduate Green Chemistry Course ~ Sarah A. Kennedy

Exploring Green Chemistry Metrics with Interlocking Building Block Molecular Models ~ Reuben Hudson, Daniel Leaman, Kiana E. Kawamura, Kristin N. Esdale, Samuel Glaisher, Alexandra Bishop, and Jeffrey L. Katz

Using a Carbide Lantern To Illustrate General Chemistry Concepts and Introduce Students to Artisanal and Small-Scale Gold Mining ~ Garland L. Crawford and Adam M. Kiefer

Flotation of Mineral and Dyes: A Laboratory Experiment for Separation Method Molecular Hitchhikers ~ Tim Rappon, Jarrett A. Sylvestre, and Manit Rappon

Energy Storage Technology

Experiments To Illustrate the Chemistry and Bouncing Ability of Fresh and Spent Zinc–Manganese Oxide Alkaline Batteries ~ Johnathon M. Hall, John R. Amend, and Thomas S. Kuntzleman

“Can” You Really Make a Battery Out of That? ~ Michael A. Parkes, Thomas Chen, Billy Wu, Vladimir Yufit, and Gregory J. Offer

Discovering Inexpensive, Effective Catalysts for Solar Energy Conversion: An Authentic Research Laboratory Experience ~ Sarah E. Shaner, Paul D. Hooker, Anne-Marie Nickel, Amanda R. Leichtfuss, Carissa S. Adams, Dionisia de la Cerda, Yuqi She, James B. Gerken, Ravi Pokhrel, Nicholas J. Ambrose, David Khaliqi, Shannon S. Stahl, and Jennifer D. Schuttlefield Christus

A Biphasic Ligand Exchange Reaction on CdSe Nanoparticles: Introducing Undergraduates to Functionalizing Nanoparticles for Solar Cells ~ Jennifer M. Zemke and Justin Franz

Outreach and Public Understanding

Demonstration Show That Promotes and Assesses Conceptual Understanding Using the Structure of Drama ~ Holly Walter Kerby, Brittland K. DeKorver, Joanne Cantor, Marcia J. Weiland, and Christopher L. Babiarz

OCTET and BIOTEC: A Model of a Summer Intensive Camp Designed To Cultivate the Future Generation of Young Leaders in STEM ~ Julie Donnelly, Carlos Diaz, and Florencio E. Hernandez

Teaching Organic Chemistry

Chemical Education Research Strategies of Successful Synthesis Solutions: Mapping, Mechanisms, and More ~ Nicholas E. Bodé and Alison B. Flynn (This article is available for free through the ACS Editors’ Choice program.)

Activities

A Simple Card Game To Teach Synthesis in Organic Chemistry Courses ~ Steven C. Farmer and Molly K. Schuman

Introducing 2D NMR Spectroscopy to Second-Year Undergraduate Chemistry Majors Using a Building-Up Approach ~ Amelia M. Anderson-Wile

Laboratories

Aza-Michael Reaction for an Undergraduate Organic Chemistry Laboratory ~ Manisha Nigam, Brittney Rush, Jay Patel, Raul Castillo, and Preeti Dhar

Introducing Aliphatic Substitution with a Discovery Experiment Using Competing Electrophiles ~ Timothy P. Curran, Amelia J. Mostovoy, Margaret E. Curran, and Clara Berger

Measuring Structural and Electronic Effects on Keto–Enol Equilibrium in 1,3-Dicarbonyl Compounds ~ Kyle T. Smith, Sherri C. Young, James W. DeBlasio, and Christian S. Hamann

Physical Chemistry

A Quantum Chemistry Concept Inventory for Physical Chemistry Classes ~ Marilu Dick-Perez, Cynthia J. Luxford, Theresa L. Windus, and Thomas Holme

Solving Simple Kinetics without Integrals ~ Lisandro Hernández de la Peña

Working toward a Paperless Undergraduate Physical Chemistry Teaching Laboratory ~ Jason D. Weibel

Exploring Biochemistry with Proteins

A Two-Week Guided Inquiry Protein Separation and Detection Experiment for Undergraduate Biochemistry ~ James P. Carolan and Kathleen V. Nolta

Utilizing Unnatural Amino Acids To Illustrate Protein Structure–Function Relationships: An Experiment Designed for an Undergraduate Biochemistry Laboratory ~ Johnathan C. Maza, Jordan K. Villa, Lisa M. Landino, and Douglas D. Young

Peptide Mass Fingerprinting of Egg White Proteins ~ Lisa T. Alty and Frederick J. LaRiviere

A Biochemistry Question-Guided Derivation of a Potential Mechanism for HbA1c Formation in Diabetes Mellitus Leading to a Data-Driven Clinical Diagnosis ~ Bomina Park, Robert W. Holman, Tyler Slade, Margaret Murdock, Kenneth J. Rodnick, and Arthur L. M. Swislocki

Research Experiences in the Laboratory

Synthesis of Copper(I) Oxide Particles with Variable Color: Demonstrating Size-Dependent Optical Properties for High School Students ~ Natalia E. Markina, Mikhail V. Pozharov, and Alexey V. Markin

An Exploration of the “Sweet Nanochemistry” Synthesis for Silver and Gold Colloids ~ Paulien Van fraeyenhoven, Natasja Glorie, Natalie Chiaverini, and Tom Mortier

Incorporating Course-Based Undergraduate Research Experiences into Analytical Chemistry Laboratory Curricula ~ Melissa A. Kerr and Fei Yan

Bridging the Gap between Instructional and Research Laboratories: Teaching Data Analysis Software Skills through the Manipulation of Original Research Data ~ Sarah J.R. Hansen, Jieling Zhu, Jessica M. Karch, Cristina M. Sorrento, Joseph C. Ulichny, and Laura J. Kaufman

Educational Resources

Commentary What’s in a Name? Amount of Substance, Chemical Amount, and Stoichiometric Amount ~ Carmen J. Giunta

Using ChemDuino, Excel, and PowerPoint as Tools for Real-Time Measurement Representation in Class ~ Malte Walkowiak and Andreas Nehring

Mining the Archives:Chemistry and Camp

Several articles in this issue discuss the appeal of teaching chemistry in a camp setting and there have been a number of posts on ChemEdX about camps, including Chemistry Summer Camp by Dan Myers. Here are some more “camp” articles from past JCE issues:

Addressing the STEM Gender Gap by Designing and Implementing an Educational Outreach Chemistry Camp for Middle School Girls~ Mindy Levine, Nicole Serio, Bhasker Radaram, Sauradip Chaudhuri, and William Talbert

Graduate Student Outreach: Model of a One-Day “Chemistry Camp” for Elementary School Students~Joseph D. Houck, Natalie K. Machamer, and Karla A. Erickson


Offering a Forensic Science Camp To Introduce and Engage High School Students in Interdisciplinary Science Topics ~ Linda Ahrenkiel and Martin Worm-Leonhard


Canisius College Summer Science Camp: Combining Science and Education Experts To Increase Middle School Students’ Interest in Science ~ Phillip M. Sheridan, Steven H. Szczepankiewicz, Christopher R. Mekelburg, and Kara M. Schwabel

Adventures in Coaching Young Chemists ~ Thomas S. Kuntzleman and Bruce W. Baldwin

JCE Contains Down-to-Earth Content

With 93 volumes of the Journal of Chemical Education to explore, you will always find something informative—including the articles mentioned above, and many more, in the Journal of Chemical Education. Articles that are edited and published online ahead of print (ASAP—As Soon As Publishable) are also available.

If you are considering writing an article for JCE, there are numerous author resources available on JCE’s ACS Web site, including recently updated Author Guidelines, Document Templates, and Reference Guidelines.

Recently, I was contacted on twitter by a concerned teacher who was using Chemistry Modeling Instruction for the first time. He was looking for supplemental materials to help his students understand why two atoms would form a covalent bond. His tweet sparked a conversation which eventually engaged some authors at AMTA (American Modeling Teachers Association). We agreed that our materials needed some work so several people started brainstorming and as a result we now have a new version of the MI teacher’s notes and a few new activities. These updated materials will be available to AMTA members this spring.  

One of the resources we have vetted is an interactive slideshow from PBS on both ionic and covalent bonds.  Teachers using Modeling Instruction will find these resources elucidate a model of electron behavior which adds to the particle story of matter we have been telling throughout the year. The ionic slideshow does a nice job explaining attraction and repulsion of different charges, and how ions arrange in a lattice formation.  The interactive nature of the covalent slideshow helps students "see" how two electrons are shared between two hydrogen nuclei.  Students can move the atoms around and watch how the nuclei repel each other while holding the electrons in the space between them. 

The resources can be found on the PBS website at the links below:

Covalent Bonds

Ionic Bonds

What am I doing to help kids achieve?

How do I know when they are there?

What is the evidence?

  Every teacher I know of has a certain student. It is THAT kid. There is always the one kid that when other students are cruising along or at least trying to cruise along, there is always the one student...THAT student...who for whatever reason does not want to play by the same rules as everyone else. He or she may just text and think you do not notice, not turn in homework or take a test in five minutes. I learned the hard way to try not to judge THAT kid. I have no idea what is happening in his or her life, if they are hungry, if they been bullied or how their home life is. It is really hard to get to know THAT kid especially when I have classes of other kids who are important and have needs also. Stack on top of this teenage hormones, spring, nice weather, prom, AP tests, state testing and trying to sell as hard as I can how fun "stoichiometry" is....I now run the risk of turning a bunch of other kids into THAT kid pretty quickly.

  I was about ready to give up...sort of. I finally tried a "plan B". The Flinn E-Learning series has a great target stoichiometry activity. I adapted it for my students. I basically gave them these instructions..."When you come in, you are going to be provided a test tube filled with baking soda. You need to heat the test tube to decompose the baking soda and predict the mass of the test tube and solid material remaining. I will have the only balance. You will have to give me your work and prediction BEFORE I place the tube and solid on the balance. The closer your prediction is to the correct answer on the balance, the higher your grade. By the way, everyone gets different starting amount of baking soda."

  It quickly changed the dynamics of the situation. THAT kid saw that it did not matter if you were the student who was the smartest or struggling the most. What mattered in this situation is that you would have to actually DO something other then a pen or paper test to get a grade. THAT kid went crazy and took over. Long story short, he predicted the correct answer within 5%. I can't say for sure this will work with every student every time but there is something to be said about doing a lab and searching for answers on a "need to know" basis instead of just saying "Let's prove and existing theory."

  Another idea I would like to toss out is the idea of test reviews. We are currently in the middle of a weird schedule and I know the best thing to do is to review a bit before a unit test but the kids are tired. So here is what I hope to do...I made up a fake test, took it as a student and am going to have the students grade it as their review instead of watching me lecture or go over problems. I am sure that one of the responses I will get is, "I can't read this writing and there are no units...." Really?

  Anyone else having end of year struggles? Got any ideas or clinchers that are a big hit with learning and the kids? Don't be afraid to comment...we are all on the same journey fighting the same battles.....

In the lab, students are given a 1.5 gram samples of copper. The copper is taken through a series of five chemical reactions ending with the precipitation of solid copper. After the five reactions, students are asked to return their 1.5 gram samples of copper to the teacher.

My dear friend and fellow AP chemistry teacher suggested that I use this lab during the first week of AP chemistry to review and reinforce lab techniques, for example, making solutions, diluting solutions, and decanting liquids. The experience prepared students for kinetics, equilibrium, and titration labs to come. 

Here is a brief overview of the lab, and you’ll find my co-worker’s lab instructions attached.

Conversion 1: Convert copper turnings to copper(II) nitrate

Students dilute 15.8 M nitric acid to 8 M. They then react the given copper turnings with the diluted acid. Warning! Nitrogen dioxide, a brown gas, is released, which requires using the fume hood. When the reaction is complete, students will have a beautiful, blue solution of copper(II) nitrate.

Skills: measurement, dilution, selection of appropriate glassware, selection of appropriate tools for measurement, use of the fume hood

Conversion 2: Convert copper(II) nitrate to copper(II) hydroxide

Using pH paper or litmus paper, students test the copper(II) nitrate solution. Then, they make a solution of 6 M NaOH and test the solution with pH paper or litmus paper. The reaction of copper(II) nitrate and NaOH is slowly performed in an ice-water bath. Students use the indicator to test the progress of the reaction. The reaction is complete when the pH paper or litmus paper color matches the NaOH’s test result. The reaction forms a pale blue precipitate, copper(II) hydroxide.

Skills: using indicator paper, measurement, making a solution, selection of appropriate glassware, selection of appropriate tools for measurement, and set up and use of an ice water bath

Conversion 3: Convert copper(II) hydroxide to copper(II) oxide

Students wash the blue precipitate with distilled water and heat the mixture to a gentle boil. The heat initiates the decomposition of the hydroxide compound forming a black, solid copper(II) oxide. When the oxide cools, students decant any remaining liquid and wash twice with distilled water.

Skills: measurement, using hot plates, and proper decanting

Conversion 4: Convert copper(II) oxide to copper(II) chloride

Students dilute a solution of hydrochloric acid to 6 M. The solution is poured over the copper(II) oxide causing an instantaneous change. The black sludge-like solid dissolves, leaving a crystal-clear green solution. Students love this conversion!

Skills: dilution and measurement

Conversion 5: Convert copper(II) chloride to copper

Students add small amounts of aluminum to displace the copper from solution. As the solution returns to a colorless liquid, students remove any unreacted aluminum and decant the fluid. The resulting copper is washed with distilled water and transferred to a watch glass. The product is dried in the drying oven and mass when cool and dry.

Skills: collection of product, use of drying oven, and measurement

A special thanks to Karen Heiges, our fellow AP chemistry teacher, who allowed me to share her work. Do you have a lab activity to share? Please post ideas as we finish this AP year and begin looking ahead. Do you have ideas for adding particle-level modeling to the lab? I look forward to reading your thoughts!

Previously I wrote about taking part in a district-wide high school blended learning pilot. You can read about it here. I received my Chromebook cart near the end of February/beginning of March. A little late but just in time for the periodicity unit I was planning as a blended unit. The following is a breakdown of how I designed the unit.

1. Day 1: Vocab self-awareness followed by a lab to classify unknown elements as metal/nonmetal. I had already done this lab in other semesters and usually I would do the lab after we defined everything. This time we did it as an explore activity that allowed for students to define terms after the fact.

2. Day 2-3: Students worked with their assigned Chromebooks to start on 4 EdPuzzle videos with quiz questions. It was self-paced so some students went back to lab to finish data collection while others got through all 4 videos and started on the next activity. All in all, it went well. Each class ranked the second day as a 8-9 out of 10 for liking the style of learning as opposed to straight up lecture.

3. Day 4: The great thing about blended learning and utilizing technology - students could continue working at their own pace while I was home sick. On the 4th day students continued on and worked through FLINN’s Periodic Table March Madness activity.

4. Day 5: I reviewed and explained concepts and had students work through a Periodicity Graphing Activity (students are given 4 sets of data, including electronegativity, first ionization energy, atomic radius and valence electrons, to graph separately with the periodic trend on y-axis and the atomic number on x-axis. Then, they answer some questions about the trends they see being demonstrated on their graphs.). Some students continued to finish EdPuzzle work on Chromebooks if they were absent.

5. Day 6: Rarely do I use a multiple choice assessment in chemistry (aside from a portion of the final exam) except during the periodicity unit. I took my multiple choice assessment and created a Google Form assessment where the question order was shuffled as was the answer order within the question. Why? To dissuade cheating.

All in all the students preferred testing on Chromebooks. One student stated that it was easier for him to focus on the computer screen during the test since he spends a lot of time outside of class looking at a computer screen. Meet them where they’re at, right?

What was interesting in this pilot experiment was entrusting students to take responsibility for their own learning while I served as a facilitator. I was able to answer questions and was on my feet more walking around helping students than I would have been lecturing from a powerpoint at the front of the room. My students are not used to learning in this way, however. Some students struggled with motivation to get things done, especially in my regular classes. At the same time, many of my students thoroughly enjoyed the freedom they had to go work in the lab next door if they were crunched for time the first day.

In the end, I’d do this again but I’d have to tweak some things. For example, test averages were lower across the board (honors and regular chemistry) with the blended unit and Google Form test than they were with my old method of teaching periodicity/periodic table and utilizing a paper test. The regular chemistry average was down about 20% while honors average was only down about 3%. This is in comparison with last semester's scores. As the averages demonstrate, my regular students really struggled. There are a variety of reasons that I won’t go into here but it certainly is concerning and may require some retooling next year.

Although not a chemistry app, I have been using Classkick in my chemistry class strictly as a formative assessment tool and wanted to share the many benefits I have found with it. Classkick is a free app that is currently available through the itunes store. I use it with the iPads I have in my classroom. Soon, classkick will be available on other devices besides just the iPad.

        One of the things I love about Classkick is its ease of use. What sets it apart from other apps I use for assessment purposes is that it turns my students iPads into a white board that they can use to SHOW ME THEIR WORK! I have used other apps for multiple choice questions but when it comes to assessments when I need to quickly see the math work of a stoichiometry problem or I ask my students to draw particulate models of atoms then I can’t truly assess this by asking a multiple choice question or even a short answer question. For me to truly gain an understanding of what my students are doing, I need to see their work. 

      Another benefit is that I can see all of my students work in real time on my iPad as it shows me each of their screens. The students can virtually raise their hands and I can join in on their screen or I can allow other students to join in on their classmates screens and help one another. I highly recommend that if you are using iPads then you check it out. The capability to post questions with pictures or with videos, then share those questions with other teachers, make this an incredibly useful tool. I look forward to using it next to show my lecture slides in class while having the capability to mark all over them while lecturing.

You Tube: Providing Feedback in High School Chemistry

There has been a fair amount of discussion on ChemEd X within the last two years about how to teach stoichiometry. (See Stoichiometry is Easy, Conceptual Chemistry, and Rethinking Stoichiometry.) I've been teaching chemistry for 22 years now, and while I've made small changes to my method of teaching stoichiometry, I haven't tried anything drastic, until this year. Inspired by the discussion going on here about stoichiometry, I decided to give BCA Tables (a method borrowed from the "Modeling Instruction" movement) a go with my introductory Chemistry class.

I have traditionally taught stoichiometry using what is commonly called the "algorithm method" of using conversions. I have a PowerPoint (attached below) and a Notes Handout (also attached below) for students. This takes them through the concept of what the balanced equation represents first. So on some level, I was attempting to go beyond the mere algorithm with my introduction to stoichiometry. Then I take them through a couple problems to practice the newly found skill of using the balanced equation as a conversion.

After the notes, I give them some practice problems (also attached below) and wander around helping students as they work. Any problems not finished are completed as homework. I give students access to a set of answers (attached below) and a set of worked solutions (attached below) so they can check their work as they go - even from home.

Following this pen-and-pencil attempt at stoichiometry, I run a stoichiometry lab the following class. I've typically used a variation of the lab provided by David in Stoichiometry is Easy but have struggled with this as I had the students evaporate the water to actually SEE the NaCl that is made. This has never given good results. Last fall I searched for an easier one-day lab that wouldn't require the same level of water evaporation. By asking #ChemChat on Twitter, I discovered an alternative: Using the thermal decomposition of calcium carbonate. (This is where I should give credit to the person that gave me this idea. I'd love to do this, but I couldn't find the conversation in a search of my history. Sorry!)

The reaction itself has an interesting color change, from the greenish copper II carbonate basic to a very stark black of copper II oxide. If you'd like to see the reaction, you can watch a video clip from a student lab last fall posted below. The lab handout requires students to go step-by-step through the process of stoichiometry - with the goal that they see the possibilities. Ideally, they will get close to 100% yield (most groups in the fall were in the 90s) and say something like, "Wow, we predicted 3.2 grams of product, and there IS 3.2 grams of product."

So this spring, I basically did all of this the same. However, I had some extra time and a group of students keen to discuss a different way of looking at stoichiometry. Attached below you'll find the PowerPoint I used along with the notes handout I created. Some of the material for the BCA notes and PowerPoint came from the previous blog posts mentioned above. I deliberately used the same sample problem as the original "traditional" set of notes so the students could focus on the process rather than the outcome. I also went beyond what I've seen with the typical BCA process by adding a bit of emphasis to the conversion to and from grams.

After my notes, I polled my students (anonymously with a Google Form) using the following questions:

Question 1

On a scale from 1-5, how confident do you feel in your ability to use the “algorithm” method we’ve already learned. (5 is most confident, 1 is not confident at all)

Question 2

Justify your answer to question 1. In other words, explain your rating. What do you like about the “algorithm” method? What confuses you about the method?

Question 3

On a scale from 1-5, how confident do you feel in your ability to use the BCA Table method we’ve just learned. (5 is most confident, 1 is not confident at all)

Question 4

Justify your answer to question 3. In other words, explain your rating. Did the BCA Table method help you understand the concepts behind stoichiometry calculations? What do you like about the BCA method? What confuses you about the method? Do you plan to use it again or stick with the “algorithm” method?

While I recognize that student opinion data isn't as valuable as the results from assessments, I do like to ask my students for feedback to see how they are processing things and where I can help. Below are some of the results and discussion of the survey results.

Student Rating Comparison

The average rating for the algorithm method was 4.05, while the average rating for the BCA Table method was 3.76. With only one class of 21 students, I certainly didn't run a statistical analysis to see if there is significance to this difference. Overall students had general confidence in both methods - and this was certainly evident based on one formative assessment the day after the notes on BCA tables. Interestingly enough, only two of the 21 students actually used BCA tables on the quiz. Most students simply used the algorithm method.

When comparing the ratings of the two methods, the numbers bear out (qualitatively) that more than half the students had either the same or increased confidence in using the BCA Table method. However, 9/21 is a relatively large number of students that didn't have as much confidence in using the BCA Table method.

The main reason for this decreased confidence may not have been the method itself, but the lack of practice. To be fair to the discussion, it is true that I didn't have as much time to dedicate to this as I would have liked for a full comparison. For many of the students that had decreased confidence in the BCA Table method, a lack of practice was mentioned as a reason. The following comment is indicative of these comments, "I understand the concept and ideas behind it, but I think I need more practice using the concept to get it right all the time and feel confident about it."

For students with confidence in both methods, two comments stood out:

"This is also confident for me because this is using concept of stoichiometry. I like that in BCA, we can find every mole of reactant and product as reaction occurs. I think I can both use BCA and algorithm method but I think BCA show me better understanding of stoichiometry while algorithm method is used for faster conversion."

"I understand the BCA method in full, but compared to the other way of solving, it is a lot more time consuming which could be helpful if one was looking at a more in depth analysis as why the problems are solved as they are."

While admitting that my use - and teaching - of the BCA Tables method is limited to this lesson, I do think the algorithm method is more efficient. However, the BCA Table method can lead to a better conceptual understanding of how stoichiometry works.

For the students that had increased confidence in the method, they generally thought BCA Tables provided a better understanding. One student said, "It is a simple method that allows us to see the process and it helps to give a clearer understanding of Stoichiometry."

This is certainly not intended to be a complete discussion of which method is "better." My sample size is 21 students, with one small attempt at using BCA tables. I also think that as teachers we should be open to some methods being more usable by some students, while other methods will work better for other students.

Having seen them in action now, I can certainly see the benefit on a few levels. While I don't think the algorithm method prevents conceptual understanding, it can often be done by students that don't really "get" the underlying reason the method works. The BCA Table method does seem to provide a more solid understanding of the mole relationships of the reaction in question. I can also see real benefit to BCA Tables for limiting reagents and equilibrium problems.

Update: After I'd written my first draft of this blog post, I was reviewing the making of buffers to my IB Chemistry students and found another scenario where BCA tables might prove helpful: Explaining why titrating 1/2 the amount of a weak acid with a strong base can give a suitable acidic buffer. The picture below shows my whiteboard from class - where a "de facto" BCA table is being used.

De Facto BCA Table for Buffer Creation

I will likely try BCA Tables again. I have not yet decided if I want to go all in with this method or not. But adding more tools to my teaching toolkit is always something I enjoy.

For those of you that use BCA Tables, I do have a question: How do you tie in the grams to moles conversion? Do you use "typical" dimensional analysis/conversions? And how do you teach students to use the mole ratios to complete the "C" row of the table? I found that students struggled with the ratio a bit (and maybe this means I struggled to teach it correctly) in trying to decide 2/3 or 3/2 for a ratio, for example. How do you address this?

What am I doing to help kids achieve?

How do I know when they are there?

What is the evidence?

I have to admit. I am really lucky. I teach with some of the best teachers on the planet. Mike Geyer is one of them. Here is how this story may help you....

Have you ever seen something that was really cool and you thought, "I could use this in the classroom, but I just don't have time to work on this now?" This might be one of those ideas. This is the kind of thing that could get you started as a fun summer project.

Mike introduced me to something called a "Raspberry Pi". A raspberry pi is the one of the world's cheapest computers. It is a $35 computer that runs off of an eight gig SD card. Anyone can program it with Python (a relatively simple coding lanquage) and it can do small things. Here is what Mike did. First he purchased a Raspberry Pi. He then went to the tech guy in our building and hunted down old computer peripherals such as a keyboard, mouse, screen and some cables. Next, he wrote a simple program that has the computer (Raspberry Pi) count continously and display the count on the computer screen. Mike displayed the time and the date of when it started and he made a nice little display explaining a bit about "Raspberry Pi" and how "Qwerty" the computer was counting to a mole. Students have to figure out when it will hit one mole. There are all kinds of ways it can be done. One student started taking pictures of it with her cell phone over time. She had a number on a screen and a time and date of the picture. She was able to start graphing and develop a reasonably good estimate on when the computer would hit one mole. Mike published this in J. Chem. Ed (J. Chem. Educ., 2014, 91 (11), pp 2005–2006).  

Here is what I like about this project.  First, kids are shown a neat and simple way to get into programing.  The "black box" effect is dispelled which is the purpose of the raspberry pi.  Next, it has them trying to predict what a mole is and when it will reach a mole.  Finally, there are many ways to solve the answer so they are allowed to think "outside" the box.  It did take some initial work on Mike's part but once it is up, it is simple to run and to forget about it.  I have used it as an extra credit project with great success.  Most of the teachers in our building have taken a liking to "Qwerty".

Sometimes our jobs can seem overwhelming.  The summer time is a great time to re-energize and to consider new projects or activities that would be fun for us and for the kids.  This might be one you may want to add to your tool kit.  Mike did a great job with the article and who knows how else a raspberry pi might be used in the chemistry classroom?  A great challenge for the kids...might want to give this project a try....

Senioritis. Seniors come down with it. Underclassmen claim it's contagious. Teachers do not appear to be exempt. Let's face it: Summer is coming and we all know it. We all want it. We are all ready for it! However, the calendar declares it is not here yet. What do you do when your students are unmotivated? What do you do when the administration tells you your class will be cut short for another assembly? What do you do when much of your class is missing for a field trip, athletics, or another end of year activity? What do you do when you have three important chapters left while all these events are taking place? How do you stay enthusiastic when all you want to do is throw in the towel, too? Let's use this space to encourage one another to finish the year strong. Go!

In one of my last blog posts I wrote of how I sometimes enjoy ending a unit with a series of demonstrations and using them to elicit a dialog between the students and myself to check for understanding. It is always a fascinating experience to hear the misconceptions that many students have the day before the test.

My kinetics unit has a very good review lecture where I take the kids through the six factors that affect reaction rate (concentration, temperature, agitation, catalysts, surface area, and nature of the reactants) by showing them demonstrations that illustrate each of the factors.  Most importantly I am trying to get my students to engage in a detailed discussion of what is happening at the molecular level. My whole approach is based on the idea of “Effective Molecular Collisions”. What is it that we can do to produce more collisions and speed up a chemical reaction? I have a phrase that I keep using over and over again: “the species must hit, they must hit hard enough, and hit in the right place”. How can these six factors help with that?

I use light sticks to illustrate the effect temperature has and ask them why light sticks in cold water glow dimmer and ones in hot water glow brighter to try to elicit a discussion of the speed at which the molecules are moving. My main point is the faster molecules at the higher temperature hit harder, more often, and because of the increased number of collisions there is an increase chance of it being in the right orientation to produce a reaction.

To discuss agitation I have a demo that is sometimes called “The Blue Bottle” or the “Traffic Light Reaction". When it is shaken in a closed flask oxygen gas in the flask collides with a mixture of dextrose and base to cause color changes in different acid base indicators. I have even seen some instructions that allow you to try and tailor it to your schools colors in one demo book but I can’t seem to recall which one. I like to point out to the students that most young kids learn these first two factors while watching people cook. Raise the temperature and mix the pot to help cook things faster. My main point here is that if we were not shaking the bottle the molecules would not be hitting as often. Especially since the oxygen is in the gas phase and not exposed to the entire volume of the solution.

Chemistry Comes Alive - Blue Bottle

To discuss surface area I have often shown a candle making contact with a block of wood, then a paper towel, and then Lycopodium powder. This is a little dangerous and maybe one I should rethink. I am not sure I want to show the old classic empty paint can version of this demo and teach students how to make an impromptu bomb. Showing a metal like magnesium in ribbon versus turnings reacting with HCl can illustrate the point nicely. My main point here is that the greater surface area provides more likely points of collision.

Chemistry Comes Alive - Lycopodium powder

For catalysts I have a great demo that shows the decomposition of tartrate ion in hydrogen peroxide. The catalyst is cobalt chloride and it changes to a deep green color while reacting (illustrating a visible activated complex) and then returns to its red color. Ironically the AP Biology teacher in the next room watched me do this one day and is now incorporating it into his lessons on enzymes. My main point here is that the cobalt provides a catalytic intermediate and surface for the reaction to happen on.

Chemistry Comes Alive - Decomposition of tartrate ion in hydrogen peroxide

For the kind of eclectic idea of the nature of reactants I show several things that do react and several that don’t. I like to put copper in water (no reaction), show a precipitation reaction, and even the classic freeze a beaker to a board reaction. In a very large lecture hall I have even done this by explodingballoons of (1) helium, (2) hydrogen, and (3) hydrogen and oxygen in a 2:1 mix. This is not for the inexperienced teacher. My main point being that some things will not react fast (or at all) no matter what you do.

Chemistry Comes Alive - Precipitation

To illustrate concentrations effect on a reaction I like to use the iodine clock reaction. I have a student help me run several mixtures in a row where I start with 50 mL of the first solution (A) and 50 mL of the second solution (B) and time the appearance of blue color. We then do the same thing with only 40 mL of solution A that has been diluted with 10 mL of distilled water to maintain a constant volume. We continue with 30, 20, and 10 mL samples each diluted with enough distilled water to make a constant volume of 50 mL. Each one takes longer than the previous one and we plot the data (mL A versus time) to get a nice curve. I ask my students to predict how many mLs of A would be needed for a reaction of a specific time that I choose and have them predict from their graphs. I like to tell them that it is a “group quiz” and if they are correct they will get 50 points. For every second they are off I will reduce their score by ten points. It can get some heated debate started in the room. My main point here being that increasing the number of molecules present increases the chance of a reaction.

Chemistry Comes Alive - Iodine Clock

Depending on the time I have available for this lecture I have several other kinetics demos I like. I show “Hooberman Balls” to illustrate activation energy, the Old Nassau Clock reaction to show intermediates, and even an Oscillator to show the same thing. For rate determining steps I use a stack of funnels on a ring stand and pour water through them into a bucket. I learned that one from a Flinn workshop.

I have been doing these demonstrations for many years. My personality has allowed me to come up with many anecdotes about them that are high school appropriate. I am sure that all of you have your own ways of tailoring these demos to your stories. 

I am very grateful to Don Showalter of the University of Wisconsin at Stevens Point for showing me how to bring these specific reactions together. Also Larry Quimby and Lanny Larsen of my school (Francisco Bravo Medical Magnet High School) spent many years dong them with me and helping me develop the presentations. In fact we often brought all three of our classes together in one room to team-teach these days. It has proved to be a very effective method of reviewing having three teachers all together in the same room at the same time. I have a great love now for team teaching.

I have attached a file to this blog post that I use to set the lecture up. It has the recipes and equipment lists for the way I do the lecture. It also contains a copy of the graph we pass out for the iodine clock experiment. As always with any experiment or lab be sure to follow the necessary safety practices and remember the five P’s. Prior practice prevents poor presentations. I have provided a basic outline of the activities and materials below. You can find more detailed directions online, in a demo book or in some of the videos linked here. 

As always you can follow all my NErDy adventures on Twitter @morganchem

I have a confession to make: I don’t really understand how to correctly predict the electronic configurations (EC) of every element and ion. 

Until recently, I thought I had everything figured out.  In fact, for quite a long time I have believed that I knew – at least in principle – how to predict the EC of any element or ion. Having read several articles in the Journal of Chemical Education on this subject,^1,2 I understood that the ground state EC of an atom or ion is a function of the lowest energy arrangement of ALL electrons in that atom or ion. I understood that confusion arises in predicting EC when the focus is placed solely on the energy of the very last electron (or last few electrons) being “placed in” orbitals during the “building up” of the atom or ion.

But I also held a misconception regarding EC: I thought that the 4s orbital was ALWAYS higher in energy than the 3d orbital. My justification for this erroneous belief came from information in two papers published in JCE^1,2. I won’t bog you down with all the details, but keep in mind that the title of one of these papers is “4s is Always Above 3d!”¹ However, I learned that this idea is incorrect during a conversation on Twitter with Dave Doherty. Dave alerted me to another publication in JCE which claims that in K and Ca, the 4s orbital is lower in energy than the 3d!³ After reflecting on that conversation with Dave, the arguments in these various JCE publications, and the data contained within, I concluded that my understanding was wrong. 4s is NOT always above 3d!

What is my point in sharing all of this with you? Well, consider that I have been teaching science and chemistry for over 20 years, I have a Ph.D. in physical chemistry, and I still don’t completely understand what is going on with EC. It is my strong suspicion that other teachers of science and chemistry also struggle with completely understanding this material and explaining it to students. There appears to be no way to present the concept of EC to students in a manner that is easy to grasp and faithful to physical reality at the same time.  Consider that one author has argued that five concepts are needed to understand “The Full Story of the Electronic Configurations of the Transition Elements”³: d-orbital collapse, d vs. s electron repulsions, s-Rydberg destabilization, configurations and states in free and bound atoms, and relativistic spin-orbit coupling. I don’t know about you, but these concepts are going to be tough for my students to understand – and present challenges to my intellectual horsepower, too! Eric Scerri has written extensively on the electronic configurations of the elements, and he has noted that “contrary to what some educators may wish for, there is no simple qualitative rule of thumb that can cope with this complicated situation”.⁴

I shared with Dave my frustrations on this subject.  I communicated with him my desire to present the prediction of the EC of elements and ions to my students in a manner that:

1. Is understandable to both teachers and students alike.

2. Chemical educators can generally agree upon.

3. Is faithful to physical reality as much as possible.

4. Can be justified with theory that both teachers and students can understand.

Dave met my challenge head on, and put together some ideas on how to present a coherent and accurate picture of electron configuration. I encourage you to read his article, Clarifying Electron Configurations! 

Dave Doherty develops physically accurate, 3D particle models of atomic and molecular structure, chemical reactions, and other chemical concepts and the Atomsmith software⁵ that allows students and teachers to interact with and “perform experiments” on these models. He can be reached at ddoherty@bitwixt.com. He uses @Atomsmith1 on Twitter to share ideas about chemistry education and he uses @AtomNMolecules to challenge college chemistry professors on topics such as electron configuration.

References:

1. Pilar, F. L. 4s is Always Above 3d!, J. Chem. Educ., 1978, 55, 2 – 6.

2. Reed, J. L. The Genius of Slater’s Rules, J. Chem. Educ., 1999, 76, 802 – 804.

3.  Schwarz, W. H. E. The Full Story of the Electron Configurations of the Transition Elements, J. Chem. Educ., 2010, 87, 444 – 448.

4. Scerri, E. The trouble with the aufbau principle, 2013, http://www.rsc.org/eic/2013/11/aufbau-electron-configuration.

5. Atomsmith Classroom, Bitwixt Software Systems, www.bitwixt.com. Available for Mac and Windows computers, and as an online HTML5 app for browsers on all platforms.

The Aufbau Principle: the (n + l) Rule

We’ve all seen and use the so-called Aufbau Diagram (Figure 1). It is a mnemonic used to remember the order of “filling” of atomic orbitals during the construction of the ground state electron configurations of the elements. The presentation of this diagram is largely disconnected from any physical meaning. Here’s what we tell our students: “Memorize the diagram, learn to use it, and you’re guaranteed to get the right answer.”

Aufbau Diagram

Figure 1. The Aufbau Diagram: Atomic orbitals are filled starting at 1s and continuing, from the upper left, in the order indicated by the arrows.

Is there a way to connect this diagram to its physical meaning? Yes! That is the goal of this article.

How was this diagram constructed in the first place? It turns out that it is a representation of a method of predicting the “order of filling” called the Madelung rule, which is also called the (n + l) rule. The “n” and “l” in the (n + l) rule are the quantum numbers used to specify the state of a given electron orbital in an atom.  n is the principal quantum number and is related to the size of the orbital. l is the angular momentum quantum number and is related to the shape of the orbital.

Here’s how the (n + l) rule works. The (relative) energies of the orbitals can be predicted by the sum of n + l for each orbital, according to the following rules:

a. Orbitals are filled in order of increasing (n + l), which represents the relative energy.
b. If two orbitals have the same value of (n + l), they are filled in order of increasing n.

The diagram in Figure 1 is the result of these rules.

Figure 2 is a version of the diagram that displays the dependence on (n + l) for each orbital, where E represents the relative energy of the orbitals. The orbitals are filled according to the values of E for each orbital: E=1 for 1s, E=2 for 2s, E=3 for 2p and 3s, and so on. According to rule (b) above, when two orbitals have the same E, such as E=3 for 2p and 3s, the orbital with lower n (2p) is filled first.

Detailed Aufbau Diagram

Figure 2. An Aufbau diagram that illustrates the (n+l) rule.

The (n + l) rule is a remarkably clever and useful tool. It correctly predicts the order of orbital energies through element 20 (calcium). It also correctly predicts many electron configurations beyond that. And here we arrive at a very important point: predicting the relative energies of each orbital is not the same thing as predicting correct electron configurations. More on this point later.

Why does the (n + l) rule work? It’s not magic and now we’ll discuss the connection between the rule and its physical meaning. To understand the connection, we need to start with how the quantum numbers n and l are related to the energy of an orbital. We’ll use 3D models (actually 2D images of the 3D models) of atomic orbitals to demonstrate. [Sorry to disappoint those looking for a deep dive into quantum mechanical calculations. These models are visual representations of the results of those calculations.]

In Figure 3, we see a representation of the orbitals occupied by the electrons in the ground state of the element krypton (for clarity, the orbitals have been separated from one another). Notice that as the quantum number n increases (from 1 to 4 in krypton), so does the overall size of the orbital.

Electronic Configuration of Krypton

Figure 3. the electron configuration of krypton. (Generated using the Electron Configuration Lab of Atomsmith Classroom¹)

How is the size of the orbital related to its energy? Recall that the potential energy of attraction between protons and electrons, which have opposite charges, depends on the distance between them: the closer an electron gets to the protons in the nucleus, the lower its energy will be. Compare the sizes of the 1s (n = 1) and 4s (n = 4) orbitals (Figure 3). Because the 1s orbital is smaller, the average distance of an electron to the nucleus will be smaller than that of the electrons in the 4s orbital. That’s the connection – the higher n is, the higher the energy of the orbital. 

What about the l in the (n + l) rule? As mentioned above, l, the angular momentum quantum number, determines the shape of an orbital. In all orbitals for which n > 1, there are areas, called nodes, in which it is extremely unlikely to find an electron. There are two types of nodes: radial and planar (or angular).  Figure 4 illustrates the radial node in a 2s orbital (l = 0) and a planar node in a 2p orbital (l = 1). Note that radial node (Figure 4, center) does not cross the nucleus, whereas planar nodes (Figure 4, right) do. s orbitals (which all have l = 0) contain only radial nodes. All other orbitals (p, d, f, etc., for which l > 0) contain both radial and planar nodes.

2s and 2p orbitals

Figure 4. Left: 2s and 2p orbitals, overlapped. Center: radial node (l = 0) in a 2s orbital (green circle). Right: a planar node (l = 1) in a 2p orbital (green line). The 2s and 2p orbitals (center and right) have been “sliced” in Atomsmith’s Orbital Lab.

The total number of nodes (radial + planar) in an orbital is equal to (n – 1). Of these, l nodes are planar.

How does the number of planar nodes affect the energy of an orbital? Look again at the radial and planar nodes in Figure 4: the planar node crosses the nucleus – where the positively charged protons are. Radial nodes do not cross the nucleus. 

If a node is an area where an electron is not likely to be found, then electrons in orbitals with planar nodes are likely to be found farther from the nucleus (on average). As we discussed earlier, large distances from the nucleus means higher energy. Thus, the higher the value of l, the more planar nodes an orbital has, and the higher the orbital energy.

So the (n + l) rule is a way to account for the two main factors that affect the relative energies of atomic orbitals: the size of the orbital (depends on n) and the number of planar nodes (= l). In cases where (n + l) is the same for two orbitals (e.g., 2p and 3s), the (n + l) rule says that the orbital with lower n has lower energy. In other words, the size of the orbital has a larger effect on orbital energy than the number of planar nodes.

Like all Models, Push Aufbau (n + l) Far Enough and it Fails.

The (n + l) rule is a model. And, as we tell our students, all models have limits. The (n + l) rule works quite well up to Z = 20, calcium (Z is the atomic number).  What does “works well” mean? It successfully predicts two things: 

1. the relative energies of the orbitals
2. the order in which the orbitals are occupied

It may not be obvious that these two things are different. But they are, and the differences start to matter at Z = 21, scandium – the beginning of the transition metals.

For Z = 20, calcium, the (n + l) rule says:

• the 4s orbital is lower energy than the 3d orbital
• the 4s orbital is occupied and the 3d orbitals are not (1s² 2s² 2p⁶ 3s² 3p⁶ 4s²).

These are both correct!

For Z = 21, scandium, the (n + l) rule says:

• the 4s orbital is lower energy than the 3d orbital
• the 4s orbital is occupied and one 3d orbitals is occupied (1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹).

Here’s where the (n + l) rule first fails. #2 (the occupation) is correct, but #1 is incorrect. For transitions metals, 3d is lower in energy than 4s! Figure 5 shows the relationship between orbital energy and atomic number (Z). Notice that the curves of the 4s and 3d orbital energies cross at Z = 21. 

4s vs. 3d Energies

Figure 5. The relationship between orbital electronic energy and atomic number (Z).

However, and this is an important point: even though the (n + l) rule gets the orbital energies wrong, it still gets the electron configuration (orbital occupancies) right!

How is it possible that, for transition metals, the 3d orbitals are lower in energy, but they are not preferentially occupied?

The short answer is that the orbital energies are not the only important factor in determining how the orbitals are occupied. The long answer? It’s complicated – very complicated.

Prof. Dr. W.H. Eugen Schwarz, a theoretical chemist at the University of Siegen in Germany has published a number of papers on this very subject. His results are clearly beyond the scope of any introductory chemistry course, but we hope to give you a flavor of how other factors besides orbital energies may influence the occupancy of atomic orbitals in an electron configuration.

Schwarz elucidates five factors that influence the electron configuration of transition metals:²

1. d-orbital collapse
2. d versus s electron repulsions
3. s Rydberg destabilization
4. configurations and states in free and bound atoms
5. relativistic spin-orbit coupling

We’ll look exclusively at the second factor: d versus s electron repulsions.

Let’s consider titanium (Z = 22). Its electron configuration is 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d², which the (n + l) rule correctly predicts. If the electron configuration depended solely on the orbital energies, we would expect: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁴ – with no electrons in the 4s orbital.

Why don’t the last four electrons preferentially occupy the 3d orbitals, which are lower in energy than the 4s orbitals?

Consider Figure 6a, where we see models of the 4s and 3d orbitals, separated in Atomsmith Classroom’s Electron Configuration Lab. If the electron configuration of titanium was      1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁴, four of the five 3d orbitals would contain one electron each. And the 4s orbital would be unoccupied.

In Figure 6b, the 3d and 4s orbitals have been superimposed on one another around the nucleus. Just as protons and electrons attract each other due to their opposite charges, electrons repel each other because they have the same charge. This repulsion results in a higher energy – things are simply getting crowded.

What can the electrons do to minimize these repulsions? Notice that the 4s orbital is larger than the 3d orbitals. If two of these electrons find their way into the 4s orbital instead of 3d orbitals, they have more space to spread out to minimize the repulsion. This is the basis of Schwarz’s “d versus s electron repulsions.”

4s and 3d Orbitals, Expanded

Figure 6a. The 4s and 3d orbitals separated.

fig06b_collapsed_white.png

Figure 6b. 4s and 3d orbitals overlapped.

How Can You Use This?

Let’s summarize what we’ve discussed so far:

Up to Z = 20 (calcium), the (n + l) rule (and the Aufbau diagram) correctly predicts:

• Orbital energy levels

The order of occupancy of the orbitals

• The physical meaning of the (n + l) rule (and its ability to make these predictions) is related to the size (n) and shape (l) of a given orbital.

For Z > 20 (starting at the transition metals):

• The (n + l) rule is not able to correctly predict orbital energy levels.
• Even when we know the orbital energies, this knowledge is not sufficient to predict the order of filling. Other factors, such as “d vs. s electron repulsions” (crowding) must be considered. (Schwarz discusses four more of these).
• Although its physical meaning is no long sufficient, the (n + l) rule still correctly predicts the order of filling. Except where it doesn’t and we invoke “exceptions.” 

The first point to be taken from this is that the (n + l) rule is a model and that it works… until it doesn’t. If you choose to teach it as a model and connect it to some of the physical meaning discussed above, it’s a great example of how models can be both useful and also fail.

The “story” outlined above has the potential to be much more fulfilling for your students than “Memorize the diagram, learn to use it, and you’re guaranteed to get the right answer.” But it’s a tough story to tell by just waving your hands. You need a model to tell it, and the model needs the following features:

1. Students must understand the basics of coulomb interactions: opposite charges are lower in energy when they are close together; the repulsions of like charges result in increased energy.
2. Representations of the atomic orbitals that are physically accurate in both size (n) and shape (l)
3. 3D is better than 2D
4. The orbitals should be separable and superimposable
5. Interactivity is desirable

You can’t tell this story without the ability to show your students the relative sizes and shapes (i.e., the nodes) of the orbitals (#2). Pictures of the orbitals in a textbook can work for many students; but all students will benefit from the ability to interactively observe (#5)  the sizes and shapes in 3D (#3), and to separate and superimpose them (#4; Figures 6a and 6b) so that they can gain an appreciation for how crowded an atom really is.

The Aufbau principal, first envisioned by Niels Bohr in 1920, and its implementation as the (n + l) rule is a very useful abstraction. Connected to its physical meaning, it can become part of powerful mental model that students can draw on to build (and explain) their understanding of the structure of the atom. This kind of connection demonstrates the real promise of 3D particulate representations of atomic and molecular structure and phenomena. Many more of these kinds of stories will be told. 

Author's note: The idea for this article arose from a discussion between the author and Tom Kuntzleman. Tom describes this interaction in a blog post, Conversations, Confessions, Confusions (and hopefully some Clarity) on Electronic Configurations.

References:

1. Atomsmith Classroom, Bitwixt Software Systems, www.bitwixt.com(link is external). Available for Mac and Windows computers, and as an online HTML5 app for browsers on all platforms.
2. Schwarz, W. H. E. The Full Story of the Electron Configurations of the Transition Elements, J. Chem. Educ., 2010, 87, 444 – 448.

            With the end of the school year approaching, educators are not only developing their semester exams, they are preparing for the upcoming school year as well. Although each individual educator has their own approach to improving their curriculum, many will be spending their summer aligning their curriculum to the Next Generation Science Standards. Currently eighteen states have adopted the Next Generation Science Standards, with additional states developing their own modified version. The idea of revising curriculum for each and every course can be daunting as educators try to identify a common theme that can be applied throughout the entire department. So where do we start? How do we thread a common theme for the professional development provided in our subject area?

            As a chemistry educator I have spent the last year working with colleagues to align our chemistry standards with NGSS. After reviewing the standards it was apparent that our chemistry department would require a shift in our instructional approach but where do we begin?

            A resource that I have found essential to developing our revised instructional approach is Brian J. Reiser’s “What Professional Development Strategies Are Needed for Successful Implementation of the Next Generation Science Standards?”. Reiser provides a thorough explanation of the interconnected goals of core ideas, practices, and coherence, and the resulting shifts in instructional strategy that should occur. In addition, he provides the reader with several challenges that educators may face when implementing the NGSS Framework such as; how to structure a lesson, how to develop explanatory models, developing an argument and reaching a consensus, and building classroom culture. This document provides justification for each instructional shift and provides ideas on how to develop curriculum in which students are continuously building on their explanatory model.

            Although this document can be an excellent resource for any individual educator I suggest sharing it with your department. According to Reiser, “the application of these ideas need to be connected to particular subject matter contexts, such as helping teachers investigate how to help students develop explanatory accounts using the particle model of matter, or evidence based arguments” (p.13). If all professional learning communities, or departments, are on the same page then professional development can then be adapted from the ideas and standards for each curriculum. I strongly recommend implementing the ideas from this resource into your curriculum, whether or not your state has adopted the Next Generation Standards or not.  

Work Cited:

I have been on a mission lately to make scientists out of my students. I am long past my fears that they are not capable of discovering the world for themselves or that they won’t learn the content if we spend too much time on science practices. What I have to work on now is orchestrating the experience. The pedagogy underlying Modeling Instruction has become the backbone for much of my instruction lately. This method of instruction not only gives my students an engaging, authentic scientific experience but has resulted in deeper content knowledge.

One thing I have struggled with is how to capture that scientific experience. How do I give my students constructive, quantifiable feedback on their scientific habits? How do I incorporate a more formal writing process that addresses the Common Core ELA standards? Some of those questions have been answered through the recently-published Argument-Driven Inquiry in Chemistry, published by NTSA. Filled with thirty “Investigation” and “Application” labs that span the curriculum, this resource helps to formalize the scientific experience. The authors’ suggested instructional methods parallel Modeling Instruction, including self-designed lab procedures and white boarded circle discussions. The labs and accompanying resources then take those Claim-Evidence-Reasoning whiteboard conversations and formalize them into an investigation report.

The lab scenarios are simple, yet engaging. When we did a version of Lab 17 (Limiting reactants), students were seeking out the ideal ratio of vinegar to baking soda. One student remarked “This is the coolest lab ever!” as we mixed these two reactants together much as he probably did in preschool. Students were actively seeking data that would provide them a means to answer the question. They were independently critiquing each other’s data and conclusions. They were asking if they could perform new and different trials to try and gather more conclusive data. When they completed their reports, they really wanted me to be convinced that their claim was correct. They were genuinely proud of their work.

I have particularly enjoyed the investigation report rubric. While there are certainly formatting considerations in lab reports, to me the most important aspect is the content. While the rubric does seek to address format and writing style, the emphasis is on the student’s articulation of their thinking. Do my students’ reports show that they can connect all the dots from a question posed to a procedure that will provide valid data? Can my students analyze their data in a way that allows them to draw a conclusion? More broadly, can my students consider a question that’s a little beyond the forefront of their knowledge and use their skills as scientists to move their understanding forward? I can use the rubric to answer those questions.

This publication is also available as an e-book. NSTA members may be able to purchase this book at a discounted price through the NSTA Science Store.

Organic chemistry was when I fell in love with chemistry. Also known as Chem 210 at the University of Michigan, it was the first time I actually started to connect what was going on at the nanoscopic level to the macroscopic world. Since then, I’ve been hooked.

As a teacher, I have grown a lot over seven years (including student teaching). How do I get my students to make connections to the nanoscopic without, you know, curved arrow mechanisms? I have been fortunate throughout the years to piece by piece gain new perspectives that I thought I’d summarize and share with you - in fact, I’ve effectively rearranged my whole first year chemistry course to be built upon a progressively more in depth view of the nanoscopic. In this post, I will focus on only my reactions unit. This is in the second semester after students learn about the mole and empirical formula.

In our reactions unit in the second semester, we go all in.

1. First, students learn how to balance, classify reaction types, and predict products.

This part is where my former mentor, Julie Andrew of CU Boulder, really helped take this unit to the next level. She shared a lab with me (that I’ve edited from her original document) where not only do students practice predicting products, but they get practice in experiencing when predictions fail. Some of the reactions are designed so that they get results that DON’T match their predictions. This allows for students to (1) make a real conclusion if their predictions matched their observations and (2) ask new questions that lead me to showing splint tests. Here is an example. When students predict the products for the reaction of hydrochloric acid and sodium bicarbonate, they predict sodium chloride and carbonic acid, since they see it as a double replacement. This is 100% valid based on what they have learned. However, this is what actually happens: HCl + NaHCO3→ NaCl + H2O + CO2.

In our whole class debrief, we go through EVERY reaction after they share their data on the board. I tell them they are smart for making that prediction once they find it fails- it’s what scientists do ALL the time- make predictions based on prior knowledge to sometimes find that they were wrong. It is helpful to give them the "maybe" option. See a sample class data set below- this gave us great ways to start class conversations.

B. Here’s where we dive into the nanoscopic - the animation project. I introduce the project, and then do a few lessons first.

1. 1. Lesson: Dissolving - ionic vs covalent compounds (edited from a posted PHET - credit is within the document). My students really struggled with the vocabulary of dissolve vs. dissociate, even though they could find errors in particulate diagrams. So the next day, this was their warm up. It was quite telling (PS- yes the phosphate has dissolved in water and has dissociated from the sodium. However, the phosphate ion itself doesn’t dissociate further).

2. 2. Notes: Redox vs. Precipitation Reactions

   3. Now, groups of 2-4 students choose two reactions from their previous lab to make animations to connect the nano and macroscopic representations. They must choose a precipitation reaction and a redox reaction. Here is a link to the videos they made.

One Bigger Con:

1. Time. It takes time to teach this and time to make the animation. I have posted a pseudo-POGIL-style PHET activity below that prepared students for animations. I also posted the project guidelines (thank you Julie Andrew at CU Boulder for the nuts and bolts of this!!!!!). Word of the wise: google slides is also awesome, and in some cases, more awesome than ChemSense referenced in the instructions.

A Few Pros:

1. Time “lost” in this project reaps benefits in dividends in stoichiometry. This is the first year (!!!) I’ve taught stoichiometry with BCA tables and the animation projects before were an unintentionally amazing scaffold. I have NEVER seen my students so successful with stoichiometry as I’m seeing now, and my hunch is that the animation projects helped make the BCA tables more accessible.

2. I sent the animations to MEL Science. They haven’t gotten back to us yet, but my students were pretty stoked to send off the animations to a company for feedback.

I hope this is helpful and sparks ideas for the future, and I’d love to hear more of what you do too!

What am I doing to help kids achieve?

How do I know when they are there?

What is the evidence?

  I am facing what many teachers are facing. It is AP week, I am trying to continue "as usual" with doing labs and learning but this time of year is anything but "as usual". There is a rates lab we do this time of year which is a good lab, rather involved with a significant amount of set up and work. I got an idea for a slightly different rates lab from Bob Worley. I found a similar large scale version from Flinn Scientific. Thanks to Bob, I decided to do a microscale version.

  Here is the lab. Students take a solution of sodium thiosulfate and add a few drops of hydrochloric acid. The reaction ends with sulfur forming and making the solution cloudy. There is a suggested series of reactions that occur before the sulfur forms. There are some concerns with doing this on a large scale. First, there is a good chance of getting sulfur dioxide which smells bad and can form sulfuric acid. Second, it is hard to know exactly when the reaction ends.

  Here is what the students did on a microscale. We counted drops of sodium thiosulfate. In our system, 20 drops of 0.1 M sodium thiosulfate and 2 drops of 1 M HCl had a reaction that lasted about 2 minutes. Students did other reactions with fewer drops of sodium thiosulfate and more drops of water (18 drops of sodium thiosulfate, 2 drops of water....16 drops of sodium thiosulfate and 4 drops of water etc...). The also tested substances, like copper, to see if it would act as a catalyst. Although we did not get to it, I also examined temperature on my own. Instead of using hot water baths, I placed the chemicals in a microwave for about 15 seconds. The solutions heated up just enough to show a difference in the reaction rate (caution..do not heat the sodium thiosulfate above 60 degrees). This was simple, quick, easy and gave great results. Students also counted the drops of solution that would make 1 mL so they could use this as a conversion to calculate the concentrations of the diluted solutions. Here is the part about this lab that I really liked. I had students use their cell phones and a watch to capture the reaction on video so they could go back and compare the reactions to determine the time for them to finish. Here is the video.

  If they had trouble seeing how each reaction ended and the times, they could go back as many times as they needed and watch the video. Overall, the data was not bad. I would probably try this again next year with a few tweeks. I wanted it to be more inquiry but we just did not have the time to fully investigate or have students develop their own experiments, but you could do it and on a microscale it is O.K. if mistakes are made. Also, because of the small scale, the sulfur dioxide was not a problem. Give it a try...let me know what you think....

What am I doing to help kids achieve?

How do I know when they are there?

What is the evidence?

I have to admit. I am really lucky. I teach with some of the best teachers on the planet. Mike Geyer is one of them. Here is how this story may help you....

Have you ever seen something that was really cool and you thought, "I could use this in the classroom, but I just don't have time to work on this now?" This might be one of those ideas. This is the kind of thing that could get you started as a fun summer project.

Mike introduced me to something called a "Raspberry Pi". A raspberry pi is the one of the world's cheapest computers. It is a $35 computer that runs off of an eight gig SD card. Anyone can program it with Python (a relatively simple coding lanquage) and it can do small things. Here is what Mike did. First he purchased a Raspberry Pi. He then went to the tech guy in our building and hunted down old computer peripherals such as a keyboard, mouse, screen and some cables. Next, he wrote a simple program that has the computer (Raspberry Pi) count continously and display the count on the computer screen. Mike displayed the time and the date of when it started and he made a nice little display explaining a bit about "Raspberry Pi" and how "Qwerty" the computer was counting to a mole. Students have to figure out when it will hit one mole. There are all kinds of ways it can be done. One student started taking pictures of it with her cell phone over time. She had a number on a screen and a time and date of the picture. She was able to start graphing and develop a reasonably good estimate on when the computer would hit one mole. Mike published this in J. Chem. Ed (J. Chem. Educ., 2014, 91 (11), pp 2005–2006).  

Here is what I like about this project.  First, kids are shown a neat and simple way to get into programing.  The "black box" effect is dispelled which is the purpose of the raspberry pi.  Next, it has them trying to predict what a mole is and when it will reach a mole.  Finally, there are many ways to solve the answer so they are allowed to think "outside" the box.  It did take some initial work on Mike's part but once it is up, it is simple to run and to forget about it.  I have used it as an extra credit project with great success.  Most of the teachers in our building have taken a liking to "Qwerty".

Sometimes our jobs can seem overwhelming.  The summer time is a great time to re-energize and to consider new projects or activities that would be fun for us and for the kids.  This might be one you may want to add to your tool kit.  Mike did a great job with the article and who knows how else a raspberry pi might be used in the chemistry classroom?  A great challenge for the kids...might want to give this project a try....

Senioritis. Seniors come down with it. Underclassmen claim it's contagious. Teachers do not appear to be exempt. Let's face it: Summer is coming and we all know it. We all want it. We are all ready for it! However, the calendar declares it is not here yet. What do you do when your students are unmotivated? What do you do when the administration tells you your class will be cut short for another assembly? What do you do when much of your class is missing for a field trip, athletics, or another end of year activity? What do you do when you have three important chapters left while all these events are taking place? How do you stay enthusiastic when all you want to do is throw in the towel, too? Let's use this space to encourage one another to finish the year strong. Go!