Brian Smith

CLU/AMGEN

Physical Science

July ‘03

Chemical Bonding in Solids

 

In a Nutshell …

Valence electrons are responsible for bonding between atoms and determine whether atoms will form ionic, covalent or metallic bonds.  The properties of the combination of atoms are the result of the bonds between them.

Covalent bonds form between atoms of similar electronegativity by sharing electrons between the atoms.  The sharing of electrons enables each atom to fill its outer electron shell (octet rule).  This results in a strong attractive force between atoms; the covalent bond.

Atoms having a significant difference in electronegativity transfer electrons to form ionic bonds.  The more electronegative atom (electron acceptor) accepts one or more valence electrons from the less electronegative atom (electron donor).  The acceptor fills its outer electron shell by adding to it and the donor ends up with a full outer electron shell (previously the filled shell beneath the partially-filled outer shell) by donating the electrons in its unfilled outer shell.  The electron donor forms a positively charged particle (cation) and the acceptor becomes negatively charged (an anion).  The electrostatic force between the cation & anion is the ionic bond.

In metals, valence electrons are delocalized.  That is, they are very loosely connected to the nucleus to which they “belong” and may be some distance from that nucleus.  As a result, they are able to move freely and metals are therefore good conductors of heat & electricity as well as being malleable & ductile.

Place in the Curriculum:

Chemical bonding will follow a thorough study of atomic structure, electron configuration and the Periodic Table.

Central Concepts:

1.    Valence electrons determine an atom’s chemical properties.

2.   Atoms combine to reduce their overall potential energy.

3.   The nature of the bond is determined by the number of valence electrons each atom needs to gain or lose.

a.    Metals form metallic bonds in which valence electrons are delocalized.

b.   A metal and a non-metal will donate and accept, respectively, one or more electrons to form an ionic bond.  As a result, each will have a filled outer electron shell; a noble gas configuration.

c.    Non-metals will combine by sharing electrons, equally or unequally (polar or non-polar), to form a covalent bond.  Each will usually have a filled outer electron shell containing its own electrons and those it shares with other atoms.

Objectives:

1.    Determine the number of valence electrons in an atom from its position in the periodic table.

2.   Draw Lewis-dot structures of the representative elements.

3.   State and apply the octet (duet) rule.

4.   Describe the role of noble gas electron configurations in ion formation.

5.   Describe the formation of cations and anions using orbital diagrams, Lewis-dot formulas, and standard ion formulas.

6.   State the characteristics of an ionic bond and recognize compounds having ionic bonds.

7.   Relate the properties of ionic substances to ionic bonding.

8.   Explain the electrical conductivity of melted and of aqueous solutions of ionic compounds.

9.   Describe the metallic bond in terms of vacant valence orbitals and loosely-bound valence electrons.  Relate some properties of metals to the type of bonding present.

10.  Describe the formation of a covalent bond between two nonmetallic elements.

Activities:

1.    Demonstrations

a.    Electrical Conductivity of Solutions … see page 1

b.   Magnetic Analogy for Bonding Forces … see page 2

c.    Force at a Distance … see page 3

2.   Labs

a.    Physical Properties and Chemical Bonding in Solids … see page 4

b.   Metals and Ionic Crystals … see page 5

c.    Combining Cations & Anions Using Model Ions … see page 6

3.   Key Questions … see page 7

Assessment:

1.    Rubric for Teacher Observation & Evaluation … see page 8 … under development

2.   Rubrics for Lab Write-ups … see page 9 … under development

3.   Written Quizzes & Tests … see page 10 … under development

 


page 1

Electrical Conductivity of Solutions

Purpose:

This demonstration provides experimental evidence on the nature of ionic and molecular substances in solution.

Materials:

·         Electrical conductivity apparatus (commercial or home-made)

·         6-8 beakers, 250 mL

·         Solid sodium chloride, NaCl (table salt) and solid sucrose, C12H22O11 (table sugar)

Safety:

If the conductivity tester used is powered by a 120v source, then use caution to prevent electric shock.

DO NOT ALLOW STUDENTS TO USE THE APPARATUS.

If you want students to test conductivity, obtain one or more low voltage (battery powered) conductivity testers

The substances used are innocuous and do not require special handling.  Solutions may be safely disposed of by flushing down the drain with water.  The solid NaCl and sugar may be used for all demonstrations and then disposed of in the trash.

Procedure:

Half-fill a 250mL beaker with solid NaCl and a second 250mL beaker with solid sucrose.  Test the electrical conductivity of the solid NaCl with the tester (light bulb remains dark). Clean the electrodes and then test solid sucrose (light bulb remains dark).  Ask students to verbalize and then record what's observed in each case.

Half-fill three 250mL beakers with distilled water.  Test the electrical conductivity of the distilled water in each of the beakers (light bulb remains dark*).  Slowly add a substantial quantity of sugar to one of the beakers with the conductivity tester in place.  Stir with a GLASS stirring rod until it’s clear that at least some of the sugar has dissolved (light bulb remains dark*).  Remove the conductivity tester from the first beaker and rinse its contacts (wires) with distilled water.  Place it in a second beaker and slowly add a substantial quantity of salt.  Stir with a GLASS stirring rod until it’s clear that at least some of the salt has dissolved (light bulb will gradually begin to glow as the salt dissolves … finally becoming quite bright as the amount of dissolved salt reaches a threshold level at which the full voltage is conducted through the solution*).  Remove the conductivity tester from the second beaker and rinse its contacts (wires) with distilled water.  Place it in a third beaker containing only distilled water.  Ask students to verbalize and then record what's observed in each case.

Analysis:

Referring to the atomic structure of the sugar molecule and the salt formula unit, draw out an inference from the students that, as the sodium and chloride ions (good time to review cations & anions and what holds them together in salt’s solid crystal lattice) accumulate, the electric current is carried by the charged particles.  Sugar, a molecular substance, produces no charged particles and doesn’t conduct electricity in solution.

* If your community water supply treats its tap water with chloroamines, any distilled water that you prepare on site will probably still conduct electricity.  Use commercially prepared distilled or deionized water.


 

page 2

Magnetic Analogy for Bonding Forces

Purpose:

This demonstration illustrates the nature of the electrostatic attraction and repulsion between like and unlike charges using a magnetic analog.

Materials:

·         4 pairs of ceramic ring or disk magnets.  In each pair, the parts need to be different in size and/or appearance.  That is, they should be significantly different sizes … or significantly different appearances (one solid & one with a hole in the center) … or, better yet, both ! (2 pair is the minimum needed)

·         overhead projector and screen

Safety:

There are no special safety concerns.

Procedure:

On the overhead projector, show that like poles repel and unlike poles attract as follows:

ü       Set up four pairs of magnets so that the different size and/or shape parts of the pair are attracted to each other.  These pairs represent the attraction between the nucleus and an electron.  Move one pair into the center of view.  Then approach with a second pair so that its "nucleus" approaches the other "nucleus."  No attraction is observed.  Move this pair so that the "electrons" are between the two "nuclei." This leads to a stable arrangement.  If enough magnet pairs are available, continue adding “nucleus-electron” pairs to show how the positive and negative parts of the “atoms” attract/repel each other.

ü       With the stable arrangement in the center, show that if either "electrons" or "nuclei" are forced closer together, they repel.  Thus an "equilibrium distance" between particles with like charges is created.

ü       Have students verbalize and then record what's observed.

Practice with the magnets prior to doing the demonstration.  It requires patience to move the magnets carefully to show formation of a "bond."  Either ring or disk magnets may be used, but the magnets should be face-magnetized.  That is, the faces should be the poles.  It's possible to use ring magnets to represent the nuclei and disk magnets to represent the electrons.  In either case, use large diameter magnets for nuclei and small diameter magnets for electrons.

Analysis:

Referring to atomic structure, .point out that this is analogous to the electrostatic forces that cause an ionic bond to form.  Both magnetic and electrostatic forces behave the same way; like charges or poles repel and unlike ones attract.  Point out that a "nucleus-electron" pair is stable because there's only one attractive force and no repulsive forces. When two pairs approach, new attractive forces arise between the "nucleus" of one pair and the "electron" of the other pair, and vice versa.  New repulsions also are present between the two "nuclei" and the two "electrons."  The result is four attractions but only two repulsions, hence the two pairs form a stable arrangement.


page 3

Force at a Distance

This simple demonstration shows that forces can exist even though objects are not in direct contact.

A paper clip is attached to a fine thread or long hair.  The other end of the thread or hair is taped to the base of the stand.  Adjust the length of the thread or hair so that there is a small gap between the magnet and the paper clip.

A piece of paper can be passed between the magnet and the paper clip to show that there is no direct, physical connection between the two.

 

 

 

 

 

 

Figure 4.  Force between a magnet and a paper clip.

 


page 4

Physical Properties and Chemical Bonding in Solids

Note:  This is appropriate for use as a student activity or a demonstration, either student or teacher.

Purpose:

This activity enables students to “see” how the atoms in solids bond together to form chemical bonds.

Materials:

Styrofoam balls (preferably different colors and sizes … blocks could be used as well), toothpicks and contact cement.

Procedure:

Establish the “ground rules” first …

ü       toothpicks represent electrons available to be shared

ü       contact cement represents the electrostatic attraction between charged particles formed by giving up (cations) or accepting (anions) outer shell electrons.

Using styrofoam balls of different colors and/or sizes (use different shapes if balls are not available, or if you want to add variety), have students prepare models of atoms that form cations & anions (such as the alkali metals, alkaline earth metals and the halogens) by coating the top, bottom, front, back, right & left sides with contact cement.  Then have students prepare models of atoms that typically share electrons by placing toothpicks at points on the styrofoam where chemical bonds will form as a result of electron sharing.

Have the students assemble the ionic crystal lattice.  If you haven’t covered the types of crystal lattices, confine your work to the face-centered cubic style illustrated by NaCl.  As they assemble the model atoms, the contact cement will form a tight bond between the ions.  As they assemble the shared electron compounds (water & methane are good examples) they should stack the molecules in a pile to represent a sample of the material in its solid state.

Have students separate the molecules of the compound(s) they’ve created to represent the amount of effort (energy) needed to melt the material.

Analysis:

Since the contact cement is analogous to the electrostatic forces in ionic crystals and the individual molecules piled together are indicative of the bonding between molecules, the contrasting difficulty of separating the ions as opposed to the molecules represents the energy required to change the material from its solid to its liquid state.


page 5

Metals and Ionic Crystals

Note:  This is appropriate for use as a student activity or a demonstration, either student or teacher.

Purpose:

This activity enables students to “see” how the electron distribution in ionic crystals and metals affects their physical properties.

Materials:

Styrofoam balls (preferably different colors and sizes … blocks could be used as well), gooey stuff (peanut butter works well) and contact cement.

Procedure:

Establish the “ground rules” first …

ü       the gooey stuff represents the delocalized arrangement of electrons found in metals.

ü       contact cement represents the electrostatic attraction between charged particles formed by giving up (cations) or accepting (anions) outer shell electrons.

Using styrofoam balls of different colors and/or sizes (use different shapes if balls are not available, or if you want to add variety), have students prepare models of atoms that form cations & anions (such as the alkali metals, alkaline earth metals and the halogens) by coating the top, bottom, front, back, right & left sides with contact cement.  Then have students prepare models of metallic atoms by coating the styrofoam balls with gooey stuff (use a lot and the physical properties of malleability and ductility will come across clearly).  If you use only one kind of styrofoam ball, your model will be that of a metallic element whereas more than one kind will represent a metal alloy.

Have the students assemble the ionic crystal lattice.  If you haven’t covered the types of crystal lattices, confine your work to the face-centered cubic style illustrated by NaCl.  As they assemble the model atoms, the contact cement will form a tight bond between the ions.  Then have students assemble metallic atoms by sticking them together with gooey stuff (use a lot and the physical properties of malleability and ductility will come across clearly) … just stick ‘em together any ol’ way.

Malleability:  Have students test the models for malleability by gently applying pressure to the assembled atoms.  The ionic crystals will resist changing shape until the pressure reaches a point at which the electrostatic bonds are broken and the crystal shatters.  The metallic atoms will flow past one another ‘til they form a single layer.

Ductility:  Have students push the mass of assembled atoms through an opening that’s smaller than the overall size of the mass of assembled atoms.  The metallic atoms will reorganize and squeeze through whereas the ionic crystal will resist going through intact.

Analysis:

Since the contact cement is analogous to the electrostatic forces in ionic crystals and the gooey stuff is indicative of the delocalized electrons in a metal, the ease of rearranging metallic atoms as contrasted with the resistance to rearrangement of the ions in a crystal illustrates their contrasting malleability and ductility.

 


page 6

Combining Cations & Anions Using Model Ions

Note:  This is appropriate for use as a student activity or a demonstration, either student or teacher.

Purpose:

This activity enables students to “see” how positively charged ions (electron donors or cations) and negatively charged ions (electron acceptors or anions) combine in certain ratios to form compounds.

Materials:

Blank paper & small sticky notes (Post-its).

Procedure:

Establish the “ground rules” first …

ü       the paper rectangle with the ion’s (monatomic or polyatomic) symbol represents all of the atom (or combination of atoms in the case of polyatomic ions) except for the outer shell or valence electrons. 

ü       sticky notes represent the outer shell or valence electrons.

ü       a paper rectangle with no sticky notes (because they’ve been donated to another atom) has a filled outer shell (the shell that was previously just below the unfilled outer shell).

ü       the number of sticky notes moved represents the amount of energy required to make the atoms combine.

ü       each electron accepted increases an atom’s negative charge by one.

ü       each electron donated increases an atom’s positive charge by one.

Have students tear a sheet of paper into 16 pieces.  On each of these, put the symbol of a representative element.  (Teacher Note: Use your judgment regarding which elements and/or polyatomic ions you include)

Attach sticky notes to each “atom” according to the number of electrons in its outer shell.

Place two model atoms side-by-side (Na & Cl are good to start with).  Transfer electrons from one atom to the other with the least expenditure of energy (see ground rules above).  If one of the species of atom cannot donate enough electrons or accept enough electrons so that all the atoms end up with a filled outer shell, add more atoms of the appropriate species ‘til, by donating and accepting electrons, all atoms have a filled outer shell.  The paper rectangle represent the formula unit of the ionic compound formed.

Analysis:

Metals attain noble gas electron configuration by donating outer shell electrons ‘til the outer shell is empty and the shell immediately below, which is full, becomes the outer shell.  Non-metals accept electrons ‘til their outer shell is filled and they achieve noble gas electron configuration.  Noble gas electron configuration is stable.

 


page 7

Key Questions

1.        What are valence electrons ?

Electrons in the highest occupied energy level of an atom are known as valence electrons. These electrons determine what kind of chemical bonds, if any, the atom can form.

2.       How can the total valence electrons for an element be determined ?

Electron configurations may be used to determine the number of valence electrons for an element. For example, hydrogen has only one electron-its valence electron. The configuration is 1s1. For representative elements, the number of valence electrons equals the total electron population at the highest principal energy level (n), as indicated by electron configurations.

3.       How many valence electrons does a sodium, silicon, beryllium, and oxygen atom have ?

sodium = 1; silicon = 4; beryllium = 2; oxygen = 6

4.       What is the relation between the number of valence electrons in atoms of an element and the element's placement in the periodic table ?

The number of valence electrons determines the group placement of an element. For example, hydrogen has one valence electron; it's in the alkali metal family. All other elements in this family, Li, Na, K, Rb, and Cs, also have only one valence electron. On the other hand, fluorine has seven valence electrons, as shown by its configuration 1s2 2s2 2p5. This places it in the halogen family.

5.       Draw orbital diagrams for atoms of sodium and fluorine. Use the diagrams to write Lewis-dot formulas for these elements.

 

 

 

 

 

6.       How does the periodic table help to determine the number of valence electrons for an element ?

Traditional group numbers (at least as commonly used in the U.S. prior to recent international moves to re-number the groups from 1 to 18) for representative or main group elements are the key.  A group IA element has one valence electron, IIA has two, IIIA has 3, and so forth.  This is not surprising, since valence electrons determine chemical properties which, in turn, determine element placements in the table.  For example, fluorine is in group VIIA and has seven valence electrons. All other elements in group VIIA … Cl, Br, and I … likewise have seven valence electrons.

7.       Why are molecules more stable than separated atoms, particularly among representative elements ?

The origin of chemical bond stability depends upon the attractive and repulsive electrostatic forces present.  Electron-nucleus interaction always furnishes attractive forces while nucleus-nucleus and electron-electron interactions furnish repulsive forces.  A simple way to show the origin of this stability is in terms of possible attraction for one or more electrons simultaneously by two nuclei close to each other.  Whether two given atoms can form a bond depends upon the filling of orbitals in the separate atoms.  Figure 6 illustrates this concept.  The hydrogen molecule is used to show that when two hydrogen atoms are close together, there's a possibility of more attractive forces than repulsive forces.  On the other hand, when two helium atoms approach, a net attractive force does not occur, although there are four new attractive forces.  This can be explained … at least in terms of this simple analysis … by the equal number of new repulsive forces also formed, as shown in Figure 6.

 

Figure 6. Bonding of the hydrogen molecule.

 

8.       How many bonds can be formed by atoms of representative elements ?

The number of bonds that an atom of a representative element can form depends upon the orbital occupancy of valence electrons in the atom. Beginning with the halogen family, the valence electron shell has seven electrons, three pairs and one unpaired electron in s and p orbitals … for example, 3s2 3px2 3py2 3pz1 in chlorine.  This enables any halogen element to form a single covalent bond by forming one electron pair with an unpaired valence electron from another atom.  This is shown by the formation of HF, illustrated in Figure 7.

 

 

 

 

 

 


In the case of oxygen family elements, each element's atom has six valence electrons.  Two of these are unpaired … for example, 2s2 2px2 2py1 2pz1 in oxygen.  Two single covalent bonds, as for oxygen in water, may be formed under these circumstances, as illustrated in Figure 8.

Figure 8: The Bonding in H2O

 

9.       What is the octet rule ?

When atoms react, they often change electron populations to acquire the stable electron configuration of a noble gas.  That is, eight electrons in the outer energy level.  For hydrogen, of course, this would be a "duet" rule with two hydrogen atoms sharing two electrons.  This is not a hard & fast rule since there are many exceptions.  It is useful in predicting the bonding expected when many atoms form compounds.  If the term "octet rule" is objectionable, an alternative is to point out that atoms tend to seek noble gas electron configurations either by electron sharing (covalent bonding) or transferring (ionic bonding) when forming compounds.

10.    In chemical reactions, do metals and nonmetals behave the same or differently with respect to sharing or transferring electrons ?

Metals generally have lower electronegativities than do nonmetals.  Thus, metal atoms attract electrons less strongly and tend to lose electrons to acquire an octet (noble gas electron configuration).  This gives the metal atom a net positive charge, resulting in a cation.  Nonmetals, on the other hand, behave in the opposite manner, having higher electronegativities than metals.  Nonmetal atoms tend to gain electrons to acquire a noble gas electron configuration, giving them a net negative electric charge.  Nonmetals tend to gain electrons and form negatively-charged ions (anions).  These generalities hold reasonably well for many reactions involving representative elements, particularly if higher members of the carbon family are excluded.

11.  Draw orbital diagrams for the sodium ion, Na+, and the chloride ion, Cl-, showing the outermost energy level only.  Write the Lewis-dot formulas and electron configurations for these species.

Teacher's Note:      Any simple monatomic ion will work with this question,  For example, Li+, F-, Mg2+, S2-, etc.  Lewis-dot formulas of single ions are usually enclosed in brackets and the ionic charge indicated.

 

 

 


12.    Draw Lewis-dot structures for NH3, H2O, Cl2O, C2H4, and SiO2.

Teacher's Note:            Many compounds may be used for this question.  Avoid "problem" molecules such as those with an odd number of valence electrons that cannot follow the octet rule (such as NO), at least initially.  When teaching students how to draw Lewis-dot structures, a useful technique is to develop a helpful set of rules, such as these:

a.    Count the total number of valence electrons in the structure.

b.   For ions, add one electron for each negative charge and subtract one electron for each positive charge.

c.    Draw the skeleton using dashes to represent electron pairs joining two atoms together until the skeleton is complete.

d.   Add dot-pairs until all valence electrons are accounted for and each atom has an octet of electrons (duet for hydrogen).

e.    If Step d is impossible when N, C, O, or S are involved, try double or triple bonds (two pairs or three pairs of dots) to form octets.

13.    Compare the covalent bonds formed between elements of similar electronegativity such as carbon and hydrogen and covalent bonds formed between elements with significantly  different electronegativity, such as hydrogen and chlorine.

Student answers will probably vary considerably.  At minimum, some reference to equal or unequal electron sharing should be made.  Students should note that unequal sharing produces a charge separation.  In any case, students should point out that polar covalent bonds, depending on molecular geometry, often give a molecule properties that affect its behavior.

14.    Describe the bonding trend expected when fluorine bonds with each element in the second row of the periodic table, including itself … F2, OF2, NF3, CF4, BF3, BeF2, and LiF.

Students should recognize that a variation in bond type from homonuclear covalent to essentially ionic takes place.  In their answers, students may note the arbitrary nature of deciding at which point polar covalent bonds are better regarded as ionic.

15. What happens to the system's total potential energy when two isolated atoms capable of bonding come into close proximity ?

As atoms approach, the atomic nuclei are attracted to the valence electrons.  If both atoms have half-filled or empty valence orbitals, then bonding may occur, lowering the potential energy of the system due to attractive forces reducing the separation between atoms.  See Figure 11.

 

 

 

 

 

 

 

16. Metals have characteristic properties that can be explained in terms of bonding.  Briefly describe the bonding in metals that explains such properties.

The basic ideas are illustrated in the lab involving chemical bonding in solids found on page 4.

17. Some substances are molecular with low melting points, soft structure, and low solubility in water.  What kind of bonding could account for these properties ?

It's probable that students will not give complete answers without help from you.  Three types of intermolecular interactions (not including the strong covalent bonding within the molecules themselves) may be responsible.  These interactions are collectively referred to as van der Wals forces.  The weakest intermolecular forces are London forces, as seen in noble gases, carbon dioxide, and other non-polar, low-melting substances.  These forces can be viewed as due to motion of electrons and formation of temporary dipoles.  The forces become stronger as the total number of electrons within molecules and the surface area of the molecules increase.  For example, carbon dioxide melts at a considerably higher temperature than does molecular hydrogen.

Another type of weak intermolecular bonding force is a dipole-dipole force.  Two kinds of these forces, distinguished by the energy required to break them, are possible.  One is the attractive force between opposite charges of polar molecules.  It is stronger than London forces but much weaker than ionic forces in ionic solids.  Molecular solids and liquids with this type on bonding generally have higher melting points.

Some compounds appear to have abnormally high melting points when compared to compounds of similar size, shape, and total electrons.  In many of these cases, such compounds exhibit hydrogen bonding, involving highly electronegative nitrogen, oxygen, or fluorine atoms.  The small size and high electronegativity of such atoms cause highly unequal sharing of the electron pair forming the covalent bond to hydrogen; there is substantial separation of charge.  The intermolecular attractions (between the hydrogen atom and a lone pair of electrons on a N, O, or F atom from an adjacent molecule) arising in this fashion are about an order of magnitude stronger than ordinary dipole-dipole bonding.  Water is an example of a hydrogen-bonded substance.  If its intermolecular forces were simple dipole-dipole forces, it would melt and boil at lower temperatures than hydrogen sulfide, H2S, which is a gas at room temperature.

 


page 8

Rubric for Teacher Observation & Evaluation

Under Development

 


page 9

Rubrics for Lab Write-ups

Under Development

 


page 10

Written Quizzes & Tests

Under Development