Chemical Bonding patterns
Unit: IIB; periodic table.
Position within Unit: After introduction of atomic structure, before introducing properties of elements (metals, non-metals, etc.). This activity is an introduction to the periodic table and its structure, and also to ionic bonds. Extension can include covalent bonds, and possibly begin discussion of properties of elements (Noble gases are touched upon, but you might be able to add others).
Title: Elementary classification
Objectives: / To develop an understanding of the periodic table and its structure. In particular, to understand that the periodic table is structured as a way to organize the elements according to chemical properties that vary in a periodic fashion, and that this organization is based on the natural patterns that effect chemical bonding. Also serves as an introduction to the concept of ionic bonds.CA Standards: / Main activity: Grades 9-12: Chemistry, 1;a, d: Chemistry 2; a (in part)
Extension Adds: Chemistry 2; a (complete), b
Materials: / You need to cut up a bunch of little colored paper squares -- 3 cmx3cm works well-- you will need at least 6 squares of green per group, 7 yellow squares/group, 3 green squares/group and 2 red squares/group It's best if you do this before class, or the students tend to get bogged down cutting up the paper and miss out on the conceptual stuff. If time is short, you can also fill them out beforehand for the students to save in-class time.
If available, samples of some of the elements so students can see them are nice to have (graphite for carbon and sulfur work fairly well) – you could also demonstrate with some of the more flashy elements if available: dropping a small (very small!) amount of sodium onto the surface of some water is a classic, flashy demo. Neither of these is necessary.
Important vocabulary: Periodic, Atom, Electron, Proton, Neutron, Ion, bond, symmetry
Activities: / The Anticipatory set(5 minutes):
Use some examples of ways that organization can make their lives easier. For example, you could have some materials scattered randomly around the classroom and have students collect them, then have things organized and make the students collect those to compare the difference in ease. Make the connection between this physical organization and organizing information. Ask students if any organization works, or if there might be some ways of organizing information that make more sense. Again, an example might help: How is a dictionary organized? Could you organize by subject instead of alphabetically? (yes) – but which is generally more useful?
The main activity is students working their way through the attached worksheet, which you should provide students. The worksheet has the specific details for the activity; additional information is provided below.
Activity A: (5 min)
The first activity has the students organize the first 18 elements based solely on physical properties, provided in the table. Students will generally lump them into a few groups (either solids liquids and gases, or by color). It doesn't matter how this comes out and don't spend much time on this: Use this as an opportunity to remind them about atoms and elements.
Activity B: (30 min)
Next, students will organize the 18 elements based on their tendency to gain and lose electrons, following the details in the handout. They should, after some trial and error, end up with a structure that looks like the top 3 rows of the periodic table. The biggest problem students tend to have with this is the idea of allowing gaps, so you may need to help them get this. (Note: only the first 18 elements are used to keep the numbers manageable and also because beginning in the 4th row, some new patterns emerge that would be very difficult to approach in this fashion).
Whatever you do, don't tell the student's that they'll be building the periodic table! The whole point is that the patterns are easy, and lead naturally to the table. They should be able to build it just by following the simple rules provided. You should try to avoid any mention of the periodic table and if there is one up in the room, take it down (preferably at least a day before the lesson so as not to draw attention to its absence). Students who realize they are building the Periodic table tend to try to remember what it looks like and reconstruct it from memory, rather construct it using the patterns -- and therefore get less out of the activity, especially if they don't already know the patterns.
There is an alternative pattern that works given the basic parameters of the lesson, except that it isn't symmetrical: The alternative pattern places the noble gases on the left side of the table rather than the right, and shifted down a row. Students who get this aren't really wrong! Depending on where you are going with the lesson, you may want to just leave it like this and readjust later (a nice way to show alternative solutions -- As in most science, one right answer just doesn't exist) or "fix" it right away (generally by trying to get the student to make it symmetrical, or sometimes just by moving helium to the right place, then have the students re-align everything). This tends to happen a lot because students seem to prefer asymmetry to leaving gaps. This is a good time to work on the meaning of symmetry.
The teacher has to resist the urge to show students the "right" solution as long as possible. Give students time to try a bunch of alternatives; just point out where the solutions don't match and offer simple suggestions and hints, but don't fix it for them. One thing that happens here is that as the quicker students get it, others will see it and the idea will spread through the room like an infection. This is OK (and speeds up the process), just try to get to each group and make sure they see the patterns as outlined below.
After they have the basic structure, ask students in their groups to show you where all the atoms that tend to gain one electron are (or gain 2, lose 1 etc.) until they see the pattern -- don't just show it to them, ask them questions that will make them discover it for themselves.
There is one big possible misconception -- and this is why students should not put + & - on the cards when they fill them out. Electrons have a negative charge, so an atom that gains an electron forms an ion with negative charge, while an atom that loses an electron becomes a positive ion. This "backwardness" is sometimes tough for students to deal with, and this activity doesn't help much on that front.
Another possible misconception: this activity can give the impression that the more central elements (N and C, for example) form ionic bonds readily. In fact, they do not (though some can be made to exist in the lab). As students move into the covalent bond section of the activity, be sure to stress the tendency of most of the elements in the "middle" of the 18-element table to form covalent bonds rather than ionic bonds. Of course there is also Hydrogen, which is rather unique in its tendency to form positive ions and be able to make covalent bonds readily.
Introduce Ionic Bonding (10 min)
The questions at the bottom of the page require the students to use the table they have built, plus the examples of stable, ionically bonded compounds given, to determine the pattern of simple ionic bonds (you can express this as either gaining or losing electrons in the outermost shell, or simply try to get them to notice that the number of electrons gained by one atom always equals the number of electrons lost by the other atoms: electrons gained-electrons lost=0).
Conclusion:
Show students a complete periodic table and explain to them that the patterns they saw as they completed the lab are why the periodic table has the strange shape it has: it is a way of classifying the elements that is extremely useful for predicting which atoms can form different kinds of chemical bonds. With this introduction the students are now primed to look for patterns in the periodic table and you can continue to introduce the many patterns within this amazing structure.
Extension:
Page 2 of the handout extends the basic structure and pattern searching behavior to covalent bonds: student find the basic pattern of covalent bond formation follows a slightly different, but related rule for "good" bond formation. This can serve as a basic intro to the Octet rule. You could extend the basic Idea of this to most of the periodic properties of elements so that, instead of showing students the patterns, they discover them for themselves.
Basic assessment is done as they work, the ultimate goal being to build the periodic table. The answers to the questions on the worksheet can also serve as assessment, especially question 5. You can also give them some novel combinations of atoms to determine whether they will form ionic/covalent bonds. Notice that they do not have a card for Potassium (K) – they need to figure out where in the table K belongs based on its atomic #.
Questions 1-5, Answers:
1) What do all of these compounds have in common? (Hint: look at the total number of electrons gained/lost). Note: This pattern holds true for ionic bonds, but not necessarily for covalent bonds. The number of electrons gained by one atom always equals the number of electrons lost by the other atom: electrons gained-electrons lost=0
2) Which classification scheme, A or B, provides you with an easier way of determining which elements might bond easily with each other? B, the classification based on electron gain/loss, works much better than A, the classification by physical properties.
3) Are there any elements that you think would be unlikely to form bonds with any other elements? Which classification scheme (A or B) did you use to identify these elements?The elements that tend to neither gain nor lose electrons (i.e. the noble gases); Classification B makes these obvious and groups them together.
4) Which classification scheme is more useful for predicting chemical reactions? B again
5) Use one of your schemes to predict which of the following compounds should be stable:
MgS; HCl; ClS; CN; NH; MgK, KCl. (The atomic # of potassium, K, is 19)
Stable: MgS; HCl; KCl
Unstable:ClS; CN; NH; MgK
Extension Answers:
1) Could the molecule N2 be stable? YES If so, how many electrons must each atom gain (share)? 3
2) What about CN? NOIf so, how many electrons would each atom have to gain (share)? No answer needed, but: Carbon needs 4, Nitrogen needs 3. Hydrogen Cyanide, HCN, is relatively stable
3) Is it possible for the element Lithium (Li) to form stable covalent bonds? NO Why or why not? Lithium would need to share 7 electrons, but only has one available to share: an atom cannot share more electrons than it has available to share under most conditions.
4) Is it possible for the element Phosphorus to form stable covalent bonds? Yes Why or why not? Needs 3 electrons and has five to share. Now is not the time to get into why phosphorus can sometimes form 5 covalent (but not very stable) bonds – save that for organic chemistry and after they know a lot more about electron configuration!
The last question would be a good opportunity to have your students produce a Venn diagram of the similarities and differences of Ionic and covalent bonds.
Reflection / Each time you run this lesson, you will discover better ways to explain what is going on and connect the info to your students. Keep track of these and change the lesson to suit your individual situation
Chemical Bonding patterns
Element / Symbol / Atomic Number / # Electrons normally gained (+) or lost (-) / Physical Properties (under normal conditions)Aluminum / Al / 13 / -3 / Silver Solid
Argon / Ar / 18 / 0 / Colorless Gas
Beryllium / Be / 4 / -2 / White Solid
Boron / B / 5 / -3 / Brown Solid
Carbon / C / 6 / + or – 4 / Clear crystal
Chlorine / Cl / 17 / +1 / Greenish-yellow Gas
Fluorine / F / 9 / +1 / Greenish-yellow Gas
Helium / He / 2 / 0 / Colorless Gas
Hydrogen / H / 1 / -1 / Colorless Gas
Lithium / Li / 3 / -1 / Silver/white solid
Magnesium / Mg / 12 / -2 / Silver-white solid
Neon / Ne / 10 / 0 / Colorless gas
Nitrogen / N / 7 / +3 / Colorless gas
Oxygen / O / 8 / +2 / Colorless gas
Phosphorus / P / 15 / +3 / White, waxy solid
Silicon / Si / 14 / + or – 4 / Grey solid
Sodium / Na / 11 / -1 / Silvery, waxy Solid
Sulfur / S / 16 / +2 / Yellow Crystal
A) In your notebook, sort the elements in the table above according to their physical properties.
B) Now we will sort them differently. Use the colored cards provided to help you.
You will use the four different colors to represent what happens to the electrons of each element.
i) Use Green for all elements that gain electrons:
ii) Use Yellow for all elements that normally lose electrons:
iii) Use Red for all elements that can either gain or lose electrons (i.e. those labeled “+ or –“):
iv) Use Blue for all elements that neither gain nor lose electrons.
Label each card with the symbol for an element in the center, the atomic number for that element in the upper left-hand corner, and the number of electrons gained or lost by that element in the upper right-hand corner. Do NOT use a + or – sign; remember to use the color-coding system above as you label cards. (see sample below).
Once your labeling is complete, arrange your cards so that:
Their atomic numbers are in sequence (i.e. you can count 1-18 easily);
AND the elements are grouped together by their colors (tendency to gain or lose electrons);
AND the cards are arranged in a symmetrical pattern.
You may use multiple rows and gaps are allowable between cards (but keep to a minimum).
Check with your lab instructor before copying your arrangement into your notebook.
Some easily formed, stable compounds are NaCl, BeO, HF and LiCl.
1) What do all of these compounds have in common? (Hint: look at the total number of electrons gained/lost). Note: This pattern holds true for ionic bonds, but not necessarily for covalent bonds.
2) Which classification scheme, A or B, provides you with an easier way of determining which elements might bond easily with each other?
3) Are there any elements that you think would be unlikely to form bonds with any other elements? Which classification scheme (A or B) did you use to identify these elements?
4) Which classification scheme is more useful for predicting chemical reactions?
5) Use one of your schemes to predict which of the following compounds should be stable:
MgS; HCl; ClS; CN; NH; MgK, KCl. (The atomic # of potassium, K, is 19)
Extension:Covalent Bonds:
Some covalently bonded molecules follow the pattern you determined previously. For example: H2O, CH4 and CO2 are all stable, covalently bonded molecules that follow the same pattern (If your pattern doesn’t work in this way, go back and revise your thinking until you see a pattern that does).
Other covalently bonded molecules don’t follow this simple pattern (For example: hydrogen peroxide, H2O2; ethane C2H6; Glucose, C6H12O6). There is another pattern that these molecules follow that is also quite simple. Use the clues below and the scheme B you developed to help you figure out this pattern.
Clue #1: remember that the elements He, Ne and Ar are chemically un-reactive – what property do they share?
Clue #2: In Ionic bonds, you looked at the number of electrons gained or lost. Number of electrons is still important in Covalent bonds, but atoms involved in covalent bonds don’t gain or lose electrons. Instead, they share electrons – each atom must share one of its own electrons in order to get an electron from the other. These shared electrons can be “double counted”; that is, each atom involved in a covalent bond gets to count the pair of electrons shared (one by each atom) as if it belonged exclusively to itself.
Clue #3: If we look at the ability to covalently bond to hydrogen we see the following pattern:
Ne cannot bond any hydrogen atoms.
One F atom can bond stably to one hydrogen atom. (HF)
One O atom can bond stably to two hydrogen atoms. (H2O)
One N atom can bond stably to three hydrogen atoms. (NH3)
One C atom can bond stably to four hydrogen atoms. (CH4)
Count the number of electrons each of the above atoms has (remember, they get to count one extra electron for each H atom they bond with) – what do the atoms in each of these covalently bonded compounds have in common? Which atom do all of these resemble in their number of electrons (their own plus shared) when they are stably bonded?