Note: A number of the links in this document provide a hint or additional information if you hover over them.
The purpose of this dry lab is to help you practise:
Applying simple Lewis electron-pair bond theory.
Using valence shell electron-pair repulsion theory (VSEPR) to predict the geometry at atomic centres in molecules or ions.
Recognizing possible canonical (resonance) structures.
Selecting the appropriate type of hydbridization of atomic orbitals to correspond to the geometry.
The program, (link below; read this first), selects semi-randomly from a data base of simple compounds and ions, and allows you to construct Lewis structures, where a bond between two atoms results from the sharing of a pair of valence shell electrons. Such bonding can be described mathematically using the wave-mechanical treatment known as "valence bond theory" which was initially developed for this purpose.
While there are other kinds of bonds, for example, multi-centre bonds which are perhaps better described by molecular orbital theory, and a variety of other interesting exceptions, the first step in describing the bonding in most molecules or molecule-ions (complex ions) involves selecting possible structures based on two-centre two-electron bonds.
Start the program by clicking this - it will appear in a pop-up window. (You may need to enable pop-ups if your browser is blocking them.)
Click "pick problem" to have the applet select a molecule, ion, or ionic compound for you to work on. The choice is made randomly from a database of around 60 species. Do 10 for inclusion in your report, and more if you want to. It is also possible to enter your own for more practice - see the notes below for more information.
The formula of a molecule, ion or ionic compound will be displayed in the box next to the "pick problem" button. If the problem is an ionic compound, you must recognize this, and just do the complex (polyatomic) ion. The monatomic counter-ion(s) must not be used.
There are two ways to proceed:
Ignore the the procedure you were given in class and construct the diagram your own way. If you have memorized the structure, or have another systematic way of getting to it which works, this might be fine. Then again, you might get a lot of error messages! If things go well, you will end up at step 7 below.
If you want to do it your own way without any intermediate verification, click the last checkbox ("All of the above!"), which will select them all.
or you can
Go through the steps one or more at a time using the other check boxes. Clicking a checkbox will select it and all the others above it. When you click "Check", the program will not check those aspects of the structure beyond the selected step. That way the number of error messages (which will appear in the large window at the bottom) will be reduced, and hopefully be less confusing.(The steps correspond to the procedure for drawing a Lewis structure which you were probably shown in class.)
Enter the total number of valence electrons. If this box is checked, the little window next to it should allow you to enter the number. When you click "check", an error message will appear if you got the number of electrons wrong. Do not forget the overall charge, if any!
Select the central atom. Type the symbol in the larger centre square box of the structure. If you have selected this step, click "check". (The "enter" key has no effect.) The program will give an error message if you have made a bad choice.
Draw the structure - single bonds only. Type the symbols for the other peripheral atoms using any of the six other larger square boxes. Click with the cursor once roughly between the centre atom and each peripheral one, and a single bond should appear. If the box for this step is selected and you click "Check", the program will (re)check step 2, and that you have correctly selected the peripheral atoms.
Add lone pairs to all atoms. Click the little circles by each symbol to add electrons. If this step is selected and you click "check", the program (re)checks the input from steps 2 and 3, and then that the peripheral atoms have an octet (except H), and that the remaining non-bonding electrons have been placed on the centre atom. (Only pairs of electrons are accepted without complaint by this version of the program - paramagnetic molecules are not handled.)
Make multiple bonds as appropriate. Click bonds again to convert a single bond to a double or triple bond, and adjust the number of lone pairs accordingly (by clicking on them again to remove them.) If this step is selected, clicking "Check" will re-check the formula and electron counts (steps 2, 3 and 4).
Specify the formal charges. Type the formal charges, if any, in the little square boxes by the atom symbols. If this step is selected, and "Check" is clicked, steps 2, 3, 4, and 5 are (re)checked, together with the formal charges you have entered.
If the structure is acceptable, clicking "OK" will bring up two questions in the large window at the bottom. The answers to these questions should be included in your report. In some cases, you may need to search (the internet) for an answer. Please write the answers you find in your own words - do not plagiarize.
After noting everything you need for your report, click "pick problem" to get the next one to work on.
For each of the 10 structures, your report should contain the following items:
A clear hand-drawn copy of the Lewis structure as it was displayed on the computer screen when it was successfully completed.In addition, do not forget to mention any counter-ions which are part of the compound which you were working on.
Specify and draw the prototype "electron pair" geometry which goes with the "steric number" (number of "electron groupings") you have arrived at. Specify the corresponding "hybridization", and the observed "molecular geometry". There may be helpful diagrams in the two questions posed by the applet after you have constructed the Lewis diagam. Marks will be awarded for artistic merit!
Give the answers to the two questions.
Ionic Compound. If a neutral compound is shown, be aware that it might nonetheless be ionic: one of the ions will be mononuclear and the other polynuclear. The monatomic counter-ion should not appear on your Lewis structure. Take care not to construct covalent bonds where there should be none! As a trivial example, KOH is really K+ and OH−; you would work on the OH− ion
Valence Electrons. You can make sure you have correctly counted the number of electrons that must be assigned to bonds or lone-pairs by entering the number in a little window that should appear if you check the box by this step. Click in the window and enter a number. Then click the "Check" button.
Symbol. Get into the habit of correctly writing atom symbols! Upper case and lower case matter e.g. Xenon is "Xe", not "xe" or "XE".
Lone Pairs. Each atom is associated with a larger square window for the symbol, three pairs of circular check-boxes for non-bonding electron pairs, and a smaller square window for the formal charge, if any. Click in the appropriate square to enter a symbol or a charge, and click in the circles to place electrons. The circles will toggle between states (electron/non electron) if you click repeatedly. For convenience, double clicking one of a pair should set/unset the other one, too.
Single and Multiple Bonds. As you click between the centre atom window and any of the peripheral atom windows, a bond or bonds will be drawn, whether or not the windows are in use. The program cycles through single, double, triple bonds and then no bond again.
Formal Charges. The program only accepts formal charges written according to the convention that has the number, if any, preceding the sign, thus "2+" rather than "+2". Charges exceeding "3+" or less than "3−" are not recognized.
Canonical Structures. Many Lewis stuctures, where all atoms have their octet look OK on paper and represent valid canonical structures, but these structures contribute less to the observed averaged structure than others with violations of the octet rule.
Thus, as the centre atom in covalent compounds, it is not unusual for the early 2nd period elements Li, Be and B to have less than an octet, but not the later ones, C, N, O or F. On the other hand, non-metals of the 3rd and lower periods often exceed their octet by using "vacant" d-orbitals in sp3d or sp3d2 hybridizations for steric numbers greater than 4 or for forming π-bonds.