Exploring chemical bonding
Grade level(s):High School (9-12), Grade 9, Grade 10, Grade 11, Grade 12
Atoms bond with each other to achieve a more energetically stable form. Atoms are more stable when their outer electron shell is complete. Atoms bond with each other by either sharing electrons (covalent bond) or transferring electrons (ionic bond).
Atom, molecule, ion, compound, electrons, neutrons, protons, electron shell/level, valence electrons, chemical bonding, ionic bond, covalent bond, electron transfer, metals, non-metals, noble gases, Octet rule
What you need:
You can borrow the following items from the SEP resource center:
- Kit 302: Exploring Chemical Bonds Kit (if you do not have access to the SEP resource center, refer to the "Getting ready" section of this lesson to put your own kit together)
- Periodic tables of the elements (one for each pair of students)
Students work in pairs or groups of three
Students will need enough desk space to lay out the content of the envelopes, so lab benches or tables will work better than individual student desks.
Students will engage in an exploration demonstrating the Octet rule and chemical bonding using paper models of elements forming covalent and ionic compounds.
Students should have explored and grasped the following concepts:
- Matter is everything that has mass and takes up space (See lesson "What is matter?" on SEPlessons.org)
- Matter is made up of tiny particles, called atoms. Atoms contain electrons, neutrons and protons (subatomic particles).
- The periodic table is a way of organizing the elements according to their characteristics and chemical behavior.
- The periodic table can be used to determine the number of protons, electrons and neutrons of atoms of different element.
- The number of protons defines an element.
- The number of valence electrons (electrons in the outer shell) determines the chemical properties of the element.
Students will be able to....
- predict whether two atoms will form a covalent or an ionic bond based on their valence electrons and their position in the periodic table
- model electron transfer between atoms to form ionic bonds and electron sharing between atoms to form covalent bonds
A molecule or compound is made when two or more atoms form a chemical bond, linking them together. The two types of bonds, addressed in this activity, are ionic bonds and covalent bonds. Atoms tend to bond in such a way that they each have a full valence (outer) shell. Molecules or ions tend to be most stable when the outermost electron shells of their constituent atoms contain the maximum number of electrons (for most electron shells that number is 8 = Octet rule). Ionic bonds are made between ions of a metal and a non-metal atom, whereas covalent bonds are made between non-metal atoms.
In an ionic bond, the atoms first transfer electrons between each other, change into ions that then are bound together by the attraction between the oppositely-charged ions. For example, sodium and chloride form an ionic bond, to make NaCl, or table salt. Chlorine (Cl) has seven valence electrons in its outer orbit, but to be in a more stable condition, it needs eight electrons in its outer orbit. On the other hand, Sodium has one valence electron and it would need eight electrons to fill up its outer electron level. A more energetically efficient way to achieve a full outer electron shell for Sodium is to "shed" the single electron in its outer shell instead. Sodium "donates" its single valence electron to Chlorine so that both have 8 electrons in their outer shell. The attraction between the resulting ions, Na+ and Cl-, forms the ionic bond.
In a covalent bond, the atoms are bound by shared electrons. A good example of a covalent bond is that which occurs between two hydrogen atoms. Atoms of hydrogen (H) have one valence electron in their outer (and only) electron shell. Since the capacity of this shell is two electrons, each hydrogen atom will "want" to pick up a second electron. In an effort to pick up a second electron, hydrogen atoms will react with nearby hydrogen (H) atoms to form the compound H2. Both atoms now share their 2 common electrons and achieve the stability of a full valence shell.
If the electron is shared equally between the atoms forming a covalent bond, like in the case of H2 , then the bond is said to be non-polar. Electrons are not always shared equally between two bonding atoms: one atom might exert more of a force on the electron than the other. This "pull" is termed electronegativity and measures the attraction for electrons a particular atom has. Atoms with high electronegativities — such as fluorine, oxygen, and nitrogen — exert a greater pull on electrons than atoms with lower electronegativities. In a bonding situation this can lead to unequal sharing of electrons between atoms, as electrons will spend more time closer to the atom with the higher electronegativity. When an electron is more attracted to one atom than to another, forming a polar covalent bond. A great example for a polar covalent bond is water:
Ionic and covalent compounds
Due to the strong attractive forces between the ions, ionic compounds are solids with a high melting and boiling point. When dissolved in water, the ions separate, resulting in a solution that conducts electricity. Covalent compounds have a much lower melting and boiling point than ionic compounds and can be solids, liquids or gases. Some covalent compounds are water soluble, some are not. But non conduct electricity when dissolved in water (unlike ionic compounds).
Over the years the model of an atom has changed. For an interesting review check out this link: http://www.clickandlearn.org/Gr9_Sci/atoms/modelsoftheatom.html
For simplicity we will be using the Bohr model throughout this lesson. It is not the latest model, but sufficiently explains the concepts of atomic structure as well as molecular bonding addressed in this lesson.
Bohr's model of the atom describes the electrons as orbiting in discrete, precisely defined circular orbits, similar to planets orbiting the sun. Electrons can only occupy certain allowed orbitals. For an electron to occupy an allowed orbit, a certain amount of energy must be available. Only a specified maximum number of electrons can occupy an orbital. Under normal circumstances, electrons occupy the lowest energy level orbitals closest to the nucleus. By absorbing additional energy, electrons can be promoted to higher orbitals, and release that energy when they return back to lower energy levels. The first energy level can hold a maximum of 2 electrons, the second and all subsequent energy levels can hold a maximum of 8 electrons (Note: This is highly simplified and partially inaccurate, however this simplification allows the students to gain a basic understanding of atomic structure and bonding before they will learn details in their high school chemistry class once they master these basic ideas.)
If you don't have access to the SEP resource center to check out the "Chemical bonding kit" follow the instructions below to put your own kit together.
- Print out the attached "Chemical_Bonds_Template" on card-stock. One set per pair of students. Laminate the printouts for repeated use.
- Cut out the "atoms" and punch a hole in the middle of each of the small circles on all of the rings, using a single hole punch. For the inner holes for the larger atoms you will need a hole punch with a long reach (at least 2"). Craft or office supplies stores carry them.
- Put brass fasteners in the holes of each atom according to the number of electrons it possesses. Elements of Group I will have one electron (one brass fastener) in the outer ring, elements of Group II will have two electrons (fasteners) in the outer rings and so on. Inner rings should be completly filled with fasteners.
NOTE: In order to safe time and fasteners, you can also NOT hole punch the inner rings and only use fasteners to show the electrons on the outermost electron level. However, make sure to tell your students that the inner electron levels are all full even though they do not hold fasteners.
- Label and assemble 3 envelopes for each pair:
- Envelope 1:
Argon (Ar), Neon (Ne), Helium (He)
- Envelope 2:
Magnesium (Mg) and Sulfur (S)
Sodium (Na) and Chlorine (Cl)
Lithium (li) and Fluorine (F)
Beryllium (Be) and Oxygen (O)
- Envelope 3:
2 Chlorine (Cl) atoms
2 Hydrogen (H) atoms
2 Fluorine (F) atoms
Chlorine (Cl) and Hydrogen (H)
- Envelope 1:
- Add a card in each evelope, showing which atoms combine to make compounds (see attached "Compounds_Info_Envelopes.pdf")
- Copy the attached "Task card" for your students
- Prepare "charge" labels by using round color coding labels (colored round sticky dots) and label half a sheet of labels with "-" and the other half with "+". You will need one sheet for each pair of students.
Lesson Implementation / Outline
- Tell your students that so far you have explored the question "What is matter made of?" and have worked with the periodic table and gained some familiarity with the elements. Some elements are found in nature in their elemental form but most elements usually don't occur in the elementary form on earth, but combine naturally with each other to create more energy-stable materials that we experience all around us like air, water, rocks and all living things. In this next activity they will explore HOW and also WHY elements combine with each other.
- Explain to students that they will work in pairs and receive three labeled envelopes, containing a group of atoms and an information card showing which atoms combine to form compounds. Their task it is to look at each group and see if they see any similarities or patterns. Hold up an example of one of the atoms in the envelope. Explain that the element symbol is in the middle, the circles represent the electron levels and the brass fasteners show the electrons.
- Instruct students to open one envelope at a time (1-3), working together to find commonalities among the atoms or compounds inside and to write those down before they move on to the next envelope. Suggest that placing atoms that form compounds side by side is helpful. They can use the periodic table to look up the elements that they have in front of you. After about 10-15 minutes they will regroup with the rest of the class and see what they have discovered.
- After students received verbal instructions, hand out the "Taskcard", the 3 envelopes and a periodic table to each group of students.
- While students work, roam the room and listen to students discussions. Answer questions as needed, but don't give away what they are supposed to discover.
- After 10-15 minutes (or when students had a chance to look at all 3 envelopes), ask the class what they have discovered. Use equal share- out strategies to make sure several pairs will have a chance to report out one discovery. Collect responses on overhead or board.
Envelope 1: He, Ne, Ar
All have a full outer electron level; are not combined with other atoms to form compounds; are all elements of group 8 = noble gases; all are non-metals.
Envelope 2: NaCl, LiF, MgO, BeS (ionic bonds)
All elements from group 1,2,7,8. Valence electrons of the two atoms in each compound add up to 8. Each compound contains one metal and one non-metal.
Envelope 3: Cl2, F2, H2, HCl (covalent bonds)
Always two non-metals combining; often 2 of the same elements bonding together.
- Ask students why they think neither of the noble gas elements bonded with another element. Tell students that atoms "strive" to be in their most stable form possible and that it turns out that they are most stable when their outer electron level is full. Since the nobel gases already have a complete outer electron level, they are already in their most stable form and therefor non-reactive. Other atoms bond together in order to achieve a more energetically stable form/ a full outer energy level. The first energy level is the smallest one and can only hold a maximum of two electrons. The second and third and the following electron levels can each hold a maximum of 8 electrons.
- Have students look at their Sodium Chloride from envelope #1. Ask: How many electrons does Chlorine have in its outer energy level? 
How many more would it need to be energetically stable? 
Sodium has one electron in its outer energy level. How many more electrons would it need to fill up that energy level? 
What could these two atoms do in order to both have a full outer energy level? [Take some student ideas]
Explain that instead of trying to get 7 additional electrons, it is energetically better for Sodium just to get rid of this single electron on the outside. If Sodium and Chlorine come together, Sodium transfers one electron to Chlorine. Now Chlorine has a complete outer energy level. For Sodium, now that it gave away its outer electron, that outer energy level is no longer existing and what is now the outermost energy level is complete and therefore Sodium also achieved a more stable form.
Have students simulate the electron transfer by taking the single electron (fastener) from Sodium and adding it to the Chlorine atom.
Ask students what charge the Chlorine and Sodium are now. Chlorine gained one negatively charged electron, so it is now negatively charged and Sodium lost one negatively charged electron. It now has more positive charges than negative charges and is overall positively charged. Remind students that we call atoms that are charged, ions. This electron transfer resulted in a Chlorine ion and a Sodium ion. Just like the opposite poles of a magnet, positive and negative charges attract. This attraction is what then bonds the two atoms together.
Pass out the color "+" and "-" stickers to students and have them stick the correct one on their newly created Chlorine and Florine ion.
Explain that this kind of bond is called an "ionic bond" because it is between two ions that are the result of an electron transfer between atoms.
Have students go through the process with their partner, using a different element pair (for example, Mg - O). Have students report out how these two atoms form the ionic bond and what charges the created ions have.
Reinforce that ionic bonds are formed between a metal and a non-metal, that elements from Group 1 (with one valence electron) bond with elements from Group 7 (with 7 valence electrons) in a one to one ratio and elements from Group 2 with elements from Group 6 in a one to one ratio. Elements of group 2 can combine with elements of group 7 in a one to two ratio (for example MgF2).
- Have students look at a pair from envelope #3, for example the 2 Fluorine atoms. Now that they know that atoms "want" to achieve a more stable form by filling up their outer electron level, ask how these 2 Fluorine atoms could possibly achieve that.
Each of the Fluorine atoms needs one additional electron to complete their outer energy level. If one takes one electron from the other it has 8 electrons in its outer shell and is full, but the second Fluorine atom now only has 6 electrons. That's not an option that works for both atoms. But if they each share one electron with the other one, they BOTH have 8 electrons in the outer level and achieved their more stable form. Demonstrate the sharing of electrons with the models. Overlap the two atom models so that two holes of each model line up and put the brass fasteners through, connecting the two atoms. Explain that the shared electrons spend part of the time circling one nucleus and part of the time circling the other one.
Explain that this kind of bond between two atoms is called a "covalent bond", the prefix "co-" meaning mutually/together/jointly as the atoms are sharing their electron pairs. Have students simulate the formation of covalent bonds between the other atoms in envelope three.
Ask students to pull off the charge label dots, rearrange the brass fasteners so that each atom has the correct amounts of "electrons" in its outer energy level and make sure to put the atoms in the correct envelopes together with the compound information card. Use the inventory sheets to help students put the atoms and envelopes together correctly.
After students explored the compounds given to them in the envelope, you can challenge them to predict what other compounds they can make with the available atoms and what bond they would form. Some students might discover that they can also make covalent molecules with more than 2 atoms, for example H2O. If nobody makes that suggestion, ask students to use their two Hydrogen atoms and one Oxygen atom to make water. Ask what kind of bond those atoms would form and have them model it. Here you also can get into the concept of polar and non-polar covalent bonds.
Have students reflect on what they understood about chemical bonding and what questions they still have. This can be done in writing as a journal reflection, as a Think-Pair-Share or as a whole class share out on the board.
Extensions and Reflections
This lesson could be followed by a lesson on nominclature of ionic compounds and chemical formulas using the same materials provided in the kit.
|ChemBonds taskcard.pdf||43.79 KB|