Monthly Archives: December 2015

Chemistry, more like cheMYSTERY to me! – Particles with internal structure


In this model building series, we last left off in Unit 4, a small but mighty chunk of curriculum. This unit introduced us to Dalton and his tiny particles called atoms. Here is what we learned:

  • All matter is made indestructible particles called atoms.
  • Different types of atoms are called elements.
  • All atoms of the same element are identical. Different elements have different properties.
  • Atoms combine chemically in simple, whole number ratios to make compounds.

I you remember, Dalton wasn’t sure how atoms combined together. It will be up to J.J. Thomson (and your students!) to answer that question!

NOTE: I do not follow the original modeling order where the mole comes next. I tried that my first year and it just seemed disjointed. I do not address the mole until right before stoichiometry.

Unit 5 is all about attractions. Bonds are not a stick or a hook that holds atoms together, they are electrostatic attractions. Bonds are also not something atoms “want”, because atoms are not people. I have wrestled with this unit for the last 3 years and I think I finally have something I like.

I kick off Unit 5 with an old favorite: the sticky tape lab.

Sticky Tape Lab

I know some people use this lab as a demo because it can be time consuming and sometimes the data are questionable but I think it is worth taking the time for. In the Sticky Tape lab, students observe the interactions between 2 charged pieces of tape and other materials including another set of charged tape, foil and paper.


The tricky part of this lab is getting the tape charged correctly. I give each group a roll of tape and tell the students to give it to the best direction follower in the group. They are usually pretty self-aware. I then make the class go through the process of laying the base tape, bottom tape and top tape down on the desk, peeling up the bottom and top tapes together, stroking the 2 pieces of tape and then quickly ripping them apart as a class. Then, I go around and check to make sure every group’s tape is properly charged by discretely holding a piece of foil to each piece of tape before I allow students to collect data. Once students collect their required data, I encourage them to experiment with items around their desks. I also encourage them to rub those items on someone’s head and then see how they are attracted to the tapes, foil and paper.

The discussion of this paradigm lab really helps put the model in students’ minds. It is all about explaining microscopic phenomena using macroscopic observations. Students can quickly guess that the tape somehow becomes charged, but the important part is what that means for our model. Simply moving atoms would not make a piece of tape charged. There must be a particle within the atom that has a charge! The Thomson Plum Pudding Model is born! Charged particles were transferred from one tape to another, making one positively charged and one negatively charged. This is also a good place to talk about Benjamin Franklin and his designation of positive and negative charges.

Plum Pudding Model (1)

Students understand that the top tape and bottom tape are attracted to each other because opposites attract, but they have a hard time explaining how both tapes are attracted to the neutral paper and foil. I like the Balloons and Static Electricity PhET for this. Students can easily see that charged objects can displace the mobile negative charge in atoms to produce a partial charge. We then talk about how the electrons can move more easily in a metal because the positive core does not hold onto the electrons as tightly compared to a non-metal (soupy pudding vs sticky pudding). This explains why the tape was strongly attracted to the foil and only weakly attracted to the paper. This model also explains why metals conduct electricity and non-metals do not which can be easily demonstrated with a 9-volt battery/light bulb circuit. The discussion of electricity is a perfect lead in to conductivity testing.

Conductivity Testing

In the past, I have done conductivity testing of various atomic, molecular and ionic substances as a demonstration with a large, 110 V conductivity tester, but this year I decided to get crafty. I sacrificed a string of LED Christmas lights to make these mini conductivity testers.


These are just a simple circuit with speaker wire as the leads, 9-volt batteries as the power source and Popsicle sticks as the base. The only tricky part was getting a good connection between the battery and the wire. Aluminum foil and a lot of tape proved very useful for this. This tutorial was also very helpful. I set up 6 stations for students to rotate through to test various solids and solutions for conductivity. The LED lights worked great at showing different levels of conductivity by lighting dimly or brightly.

Students were able to classify their data into 4 categories: elemental solids that conduct electricity, elemental solids that do not conduct electricity, solutions that conduct electricity and solutions that do not conduct electricity. After students noticed that the solutions that conducted electricity contained a metal, we named these ionic compounds and the other type of compound, molecular. Since ionic compounds conduct electricity, they must be composed of charged particles. The obvious next questions is, which is negative and which is positive?

Micro-Electrolysis of Copper (II) Chloride

This micro-electrolysis is an activity I have added to this unit to introduce anions and cations. I instruct students on how to set up a very simple micro-electrolysis with aquarium tubing, a 9-volt battery and mechanical pencil lead (.7 mm or thicker works best).


After letting the electrolysis run for a few minutes, students see bubbles forming on the positive electrode and with some careful wafting, they can identify it as chlorine gas. When students look carefully, they see the negative electrode is turning a reddish-brown color. Students immediately call this rust. I always ask, “what is rust?” and the students reply “iron oxide?” I then ask, “is there any iron in the solution?” Students then realize that the reddish-brown substance cannot be rust and must be copper. Since the copper (metal) is attracted to the negative electrode, it must be positively charged. That means the chlorine (non-metal) must be negatively charged. This is when I introduce the terms “cation” and “anion.”

Patterns of Charge 

The Modeling materials has a worksheet called “Predicting Formulas” which I have students complete after talking about anions and cations. This worksheet gives students a variety of ionic compounds and helps them find the patterns in which the ions combine. I always intro this worksheet with, “we know from the last unit that we can find the formulas for compounds using mass ratios” so students understand where these formulas come from. After completing and discussing this worksheet, students can identify the basic patterns of charge for the main group elements.

We have been zeroing in on ionic compounds for a little while, but it is time to zoom back out and look at molecular and atomic substances as well.

Structure with MolView

In the past, I have had access to 7 laptops that I could install the Mercury Software on to look at the structures of various ionic, molecular and atomic substances. I am at a new school this year so I had to find a ChromeBook alternative. Enter MolView. MolView is an awesome Mercury Software alternative. It does not have all the compounds that Mercury does and you have to do an advanced search in the Crystallography Open Database to get unit cell structures, but it gets the job done.


I had each student manipulating the structures on a ChromeBook and I also put the structures up on the SMARTboard so we could discuss them as a class.

I made a big deal this year about ionic compounds being bonded throughout because ions are charged spheres, meaning they attract particles of the opposite charge in every direction.

From looking at the structures, students constructed rules for classifying ionic, molecular and atomic substances. I always like to show students the structures of graphite and diamond to get the discussion started on “why structure matters?” Maybe that’s just the geologist in me!

After this activity, I have students complete the “Why Structure Matters” worksheet from the Modeling materials to relate structure to melting and boiling points.

This year, I added something new before getting to nomenclature. After talking about melting and boiling points, the next obvious place to go seemed to be intermolecular forces. In the past, I talked about how there are attractions between molecules that are not as strong as ionic and covalent bonds and in Unit 3, energy had to be put into a system to overcome these attractions to change phase, but I never gave these attractions a name.

Intermolecular Attractions (Forces)

I don’t like the term intermolecular forces so I call it intermolecular attractions (IMAs), because it is more descriptive of what is actually going on. For IMAs, I borrowed a lab from my colleague across the hall and “model”fied (that’s a thing, right?) it. Students timed the evaporation of 6 molecular substances: pentane, hexane (switching out for butane next year), ethanol, methanol, ethyl acetate and acetone.

img_0340 (1)I thought students might be bored by this lab because it is kind of like watching paint dry but they were actually very enthused about how quickly some of the substances evaporated and how they “disappeared” before their eyes. I heard some great hypotheses as the students talked about which ones would evaporate fastest: “it must have something to do with mass” and “these ones have oxygen in them and these don’t.”

Students saw that the evaporation times broke the substances into 3 groups: molecules without oxygen, molecules with oxygen but no OH group, and molecules with an OH group. I named the attractions in the first group “induced dipole-dipole attractions” and the second and third group “permanent dipole-dipole attractions” (because of the electronegative oxygen). I explained that the molecules with the OH group have a special kind of permanent dipole-dipole attraction called hydrogen bonding.  I of course had to introduce the term “dipole” and we talked about why permanent dipole-dipole attractions seemed to be stronger than induced dipole-dipole attractions. I did not use the terms dispersion forces or Van der Waals forces because they are not descriptive of what is actually happening.

Thoughts before moving on to nomenclature

I think this unit is the toughest modeling unit to teach because you have to teach bonding without the Bohr model. The great thing about that is you are not breeding the “atoms want 8 electrons” misconception. Atoms don’t “want” anything, they are atoms. Bonding is all about electrostatic attractions, not a set of rules. My advice is hit this hard!

Ions are formed by gaining or losing electrons. When an ion forms it is charged all over so it attracts particles of opposite charge in all directions. This is why ionic compounds do not exist in discrete, formula units. This is also why 1 sodium atom bonds with 6 chlorine atoms but only has a +1 charge.

By contrast, molecular compounds bond within molecules because the electrons of each atom are attracted to the positive core of the other atom. This is why molecular compounds do not form lattices but instead are held together by weaker intermolecular attractions.

I highly recommend reading Beyond Appearances: Students’ misconceptions about basic chemical ideas (Kind, 2004). The ideas in this paper really helped me get the big picture of this unit.

Ionic and Molecular Nomenclature

Moving on, the last thing to hit in Unit 5 is ionic and molecular nomenclature. I also took a new approach to this topic this year and had students work more independently than usual. I created a “Chemistry Ninja Warrior” system where students had to “level up” to different types of nomenclature. Different levels earned different cool stickers.

5813524696662016 (1)

Each level had a test (worksheet) that students had to demonstrate mastery on before they moved to the next level. The goal was for every student to reach level “ninja turtle” and more advanced students could move beyond that. Students worked independently or with other students at their level on POGIL activities to learn the nomenclature rules.

I liked that my students got to work at a pace that worked for them and I got to spend more time with the students who need  1:1 attention. I used the POGIL activities this year as is but next year I think I will edit them after seeing some of the snags my students ran into.

After spending some significant time on naming, the only thing left is a practicum!

Unit 5 Practicum

This practicum is more like a “demonstration of knowledge” than a lab challenge. Each group is given a set of tables containing names of elements and a pair of dice. How the students roll the dice determines the compound they will build. They must write the formula and name of the compound they roll. For molecular compounds, students roll the dice again to get the number of each type of atom in the molecule. I also have students construct a few rules for naming and differentiating ionic and molecular compounds. This is not my favorite practicum but it does a nice job of wrapping up the unit.

Whew! I think that is a hard unit to wrap your head around! Let’s sum up the model so far…

  • All atoms contain mobile, negatively charged particles called electrons whose charge is balanced by the positive (pudding) core of the atom. (Thomson’s Plum Pudding Model)
  • In metals, the positive core has a weaker attraction to the electrons so electrons can move more freely than in non-metals, allowing metals to conduct electricity.
  • Metals tend to lose electrons and become positively charged cations and non-metals tend to gain electrons and form negatively charged anions.
  • Ions are charged all over and attract ions of opposite charge from all directions. When ions of opposite charges are attracted to each other, they form ionic bonds. Ionic substances are bonded throughout and have high melting/boiling points.
  • When the electrons of two non-metal atoms are attracted to the other’s positive core, a covalent bond is formed. Molecular compounds are bonded within molecules but the molecules are only attracted to each other through intermolecular attractions. Molecular substances have lower melting/boiling points compared to ionic substances.
  • Molecules can be attracted to each other through induced dipole-dipole attractions and permanent dipole-dipole attractions.
  • Ionic compounds are named by writing the metal first and then dropping the ending of the non-metal and adding the suffix -ide.
  • Molecular compounds are named by using the prefixes -mono, -di, -tri, -tetra, etc. to denote how many atoms of each element are present in the compound. The first element only gets a prefix if there is more than 1. For the second element, you must drop the ending and add the suffix -ide.

Thanks for sticking with me through that one! Stay tuned for Unit 6: Chemical Reactions!