ELECTRON ARRANGEMENT& BONDING

Teacher's Guide

Electron Arrangement and Bonding Teacher's GuideThis teacher's guide is designed for use with theElectron Arrangement and Bonding series ofprograms produced by TVOntario and available onvideotape to educational institutions and nonprofitorganizations. The programs are broadcast byTVOntario, the television service of The OntarioEducational Communications Authority. For broad-cast dates, consult the appropriate TVOntarioschedule. Ordering information for videotapes andthis publication appears on page 24.

Canadian Cataloguing in Publication DataStratton, John.

Electron arrangement and bonding. Teacher'sguide

To be used with the television program, Electronarrangement and bonding._Bibliography: p.I SBN 0-88944-081-6

1. Electron arrangement and bonding (Televisionprogram) 2. Electrons - Study and teaching(Secondary) 3. Chemical bonds - Study andteaching (Secondary) I. TVOntario. II. Title.

QD461.S771985 54122 C85-093032-4

©Copyright 1985 by The Ontario EducationalCommunications Authority.

All rights reserved.

The SeriesProject leader: David ChamberlainNarrator: James MoriartyContent Development: William KonradConsultants: Horst Beck, Phil GedenComputer animation: Animations Drouin Inc.Audio: Mars StudiosDesigned and produced by Northey Productionsfor TVOntario,1984

The GuideWriter: John StrattonConsultant: John EixEditor. Deborah BurrettDesigner: Michael E. Bowness

Note to TeachersThe Electron Arrangement and Bonding series forsenior chemistry students explains the importanceof the electron to chemical compounds. The six10-minute programs follow the development ofthe atomic model and its contribution to theunderstanding of the function of the electron.Aspects such as energy levels, orbitals, charge,relative mass, bonding, and electron configurationare all clearly demonstrated with computeranimation. Students will learn about the propertiesof an atom and how these may be predicted bystudying electron arrangement and bonding.

This guide contains a detailed description ofeach program. It also provides learning objectives,activities for before and after viewing, and furtherreading suggestions.

ContentsProgram 1: Introducing the Players ............. 1Program 2: The Rutherford-BohrAtom .......... 7Program 3: Electron Arrangement ............. 10Program 4: How Atoms Bond ................ 13Program 5: Molecular Substances and

Covalent Crystals . 17Program 6: Metals and Ionic Solids . . . . . . . . . . . 21

Printed in Canada 1706/85

Introducing— the PlayersObjectivesStudents should be able to:

1. Understand the concept of an atom as the smallest completeparticle of a substance.

2. Illustrate, using asimple diagram, Rutherford's model ofthe atom.

3. Describe the electron, proton, and neutron with respect to theparticle's location, charge, and relative mass.

4. Understand the importance of the proton in identifying anelement.

5. Understand the importance of the neutron in contributing onlymass to an atom of an element.

6. Explain why electrons in Rutherford's model of the atom"should" spiral into the nucleus.

Program DescriptionProgram 1 introduces the atom, the smallest particle of matter. Theprogram also introduces the three main subatomic particles: theelectron, the proton, and the neutron. The location, charge, andrelative mass of each of these particles is demonstrated; thestability of the nucleus is discussed; and the fact that everyelement can be identified by the number of protons in the nucleusof the atom is explained. In the history of the discovery of atomicstructure, the work of Ernest Rutherford and Thomas Chadwick isgiven prominence. In particular, Rutherford's model of the atom isreviewed in terms of its advances and shortcomings. Finally, theprogram raises two questions about atomic behavior: "Why dosome atoms form bonds with other atoms while some do not formbonds at all?" and "Since charged accelerating particles loseenergy through radiation, and since electrons orbiting a nucleusare charged accelerating particles, why don't the electrons spirali nto the nucleus?" The answers to both these questions require arefined model of electron arrangement. This is the subject of theprograms that follow.

Structure of the Atom series or provide a time line modelled onthe one that follows.1897 Electrons were discovered.1902 Lord Kelvin (William Thomson) suggested that the atom

was a sphere of positive charge with electrons embeddedi n the positive fluid.

1903 J.J. Thomson suggested that electrons were arranged incertain fixed positions within the positive sphere.

1903 Philipp Lenard suggested that the atom was mostly emptyspace with electrons paired with positive particles of thesame size. The pairs, which he named "dynamids," werespread throughout the atom, leaving empty space betweenthem.

1904 H. Nagaoka suggested that electrons revolve at a fixedspeed in a ring around a large positive centre.

Electrons

Lord Kelvin's Model J.J. Thomson's Model

Dynamid

Positive Sphere

Space

Large Positive Centre

Before ViewingBriefly describe early models of the atom to give the students ani dea of what Rutherford had to work with as he developed hismodel. You could show the first few programs of TVOntario's

P Lenard's Model H. Nagaoka's Model

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After discussing the evolution of the atom up to Rutherford's time,have students draw diagrams of early models of the atom andcompare them with Rutherford's ideas. This should prepare thestudents for introduction to the idea presented in the program ofelementary particles or "building blocks" of atoms.

After ViewingHave students complete activities 1 and 2.

Activity 1 involves a demonstration by the teacher of four _ _experiments performed by Hittorf, Crookes, Perrin, and J.J.Thomson when they investigated the structure of the atom.Students should note their observations and answer the questionsthat follow each part. Concluding Activity 1 is an exercise inwhich students are asked to simulate Millikan's thinking when hedetermined the charge on the electron. Activity 2 simulatesRutherford's scattering experiments involving gold foil and alphaparticles. This activity should help students understand whyRutherford thought the atom was mostly empty space.

ActivitiesActivity 1: Gathering Evidence

I n this activity, you will demonstrate forstudents as many of the experiments done byHittorf, Crookes, Perrin, and Thomson aspossible with the equipment available in yourschool. You will also use an analogy to helpstudents simulate Millikan's reasoning whenhe determined the charge on the electron.

Materialscathode ray tubeshigh-voltage sourcevacuum pumpmagnetsealed film cans containing different numbers

of the same objectbalanceconnecting wires

Part A Production of Cathode RaysAn apparatus consisting of a glass tubecontaining two metal plates is evacuatedslowly. The plates are connected to ahigh-voltage source. The plate connected tothe negative terminal is called the cathode.The plate connected to the positive terminali s called the anode.

As the pressure is reduced, a discharge ofelectricity occurs. At about 6.7 kPa, thinzig-zag sparks or streamers pass from thecathode to the anode.

At about 1.3 kPa, a bright column of lightfills the tube. The color of light depends on thegas in the tube. Sodium gas produces a yellowcolor and neon gas produces a red color.

At about 0.07 kPa, the column of lightrecedes towards the anode, and the glassopposite the cathode begins to glow.

Questions

a) Explain why the streamers of light in thetube were called "cathode rays."

b) What evidence suggests that atoms ofthe electrodes are not travelling from oneto the other?

c) What evidence suggests that energy ismoving between the metal plates?

d) What evidence suggests that the stream-ers are not particles of the gas?

e) List two plausible explanations for theproduction of the streamer.

f) Suggest a plausible explanation for thetube glowing at very low pressures.

Part B Hittorf's Experiment, 1869A solid object placed in front of the cathodecasts a shadow on the glass.g) Explain how this observation could be

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Part C

Cathode rays are passed through a slit andalong a zinc sulfide screen. A greenish beami s observed along the screen.

When a magnet is brought close to the tubethe greenish beam is deflected.

Part D Perrin's ExperimentPerrin found that when cathode rays struck anobject, the leaves of a gold-leaf electroscopeplaced beside the object spread apart. He alsofound that when a magnet was used to deflectthe cathode rays from the object, theelectroscope remained unchanged.j) If the greenish beam on the screen was

caused by a ray of light passing throughthe tube, would the magnet deflect it?

k) I f magnets deflect the streamers we callcathode rays, then how must thestreamers be produced?

used to support the idea that "rays" were,coming from the cathode.

h) Does this observation rule out thepossibility that particles were passingbetween the electrodes? Explain.

i) If particles are passing between theelectrodes, in what direction are theymoving?

Crookes' Experiment, 1879

I) What additional information about thestreamers is provided by Perrin's electro-scope experiments?

m) What are "cathode rays"?

Part E Thomson's Experiment, 1897Thomson deflected a beam of cathode rays byapplying a known voltage across the electricalplates. He then varied the strength of the

electromagnet until the beam came back to its.original position. From the voltage across theplates and the strength of the magnetic fieldproduced by the electromagnet, he was able__to determine the ratio ofthemass to charge(m/e) of the particle passing from the cathodeto the anode. Thomson found that the m/eratio didn't change even though he usedvarious voltages, various metals for theelectrodes and various gases in the tube.Thomson concluded that the nature of thecathode rays was independent of how theywere produced.n) Explain why Thomson's observations

dealt the death blow to the idea thatcathode rays were rays of light passingdown the tube.

o) What evidence supports Thomson'sconclusion "that atoms of all substancescontain the same kind of negativeparticles"?

p) I n 1861, six years before Thomsonmeasured the mass-to-charge ratio,Stoney suggested the name electronfor the particle produced in cathode raytubes. However, Thomson is credited withthe discovery of the electron. Suggestplausible reasons why Thomson is cred-i ted with the discovery of the electronrather than Crookes, Perrin, or Stoney.

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Part F Millikan's Experiment, 1909Millikan blew a cloud of fine oil droplets withatomizer(A) into the large dust-free chamber(C). Some of the droplets fell through thepinhole (P) into the space between two metalplates (M) and (N). Millikan was able to observethe behavior of the droplets between theseplates, he was able to make the droplets riseor fall. From measurements of the voltage andthe time required for the droplet to rise or fall agiven distance, he was able to determine thetotal charge on a droplet. From this informa-tion, he determined the charge on an electron.

Since the experiment is difficult to perform,the following analogy may help you simulatethe reasoning Millikan used to assign a chargeto the electron. You are given several sealedfil m cans (representing oil droplets), whose

total mass (representing charge) you candetermine with a balance. You can assumethat all the film cans have the same masswhen empty and that each can holds adifferent number of the same unknown object.Your problem is to determine the mass of oneof the contained objects without opening afilm can.q) Describe how you plan to determine the

mass of one of the objects in the sealedfil m cans.

r) Record the data and the necessarycalculations to determine the mass ofone of the objects.

s) The following table includes some of thedata gathered by Millikan. Use the samereasoning as you used in the analogyabove to determine the charge on an

Activity 2: A Simulation of theGold Foil ExperimentThe Nuclear ModelI n 1909 Marsden and Geiger, two members ofthe physics department working underRutherford at Manchester University, directeda particles from a radioactive source at atarget of very thin gold foil. Their observationsare simulated in this activity.

Here the gold foil will be represented byplasticene with an object embedded in it; or bytwo large pieces of cardboard, one pasted over

each side of a smal I piece of metal sheeting;or by a target concealed under a shield. Theparticles will be represented by cdissectingneedles or marbles. The dissecting needlescan be used to probe the plasticene andcardboard. The concealed targett can beprobed by rolling marbles at the target.

Materialsplasticene with objects embedded in itmetal sheeting between piecesoof cardboarddissecting needlesconcealed target, ramp, and marbles

Procedure1. Probe the "gold foil" (plasticene, card-

board or concealed target) with "alphaparticles" (dissecting needles or marbles).

2. Draw a picture of and descriibe yourobservations.

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ExperimentNumber

Total Chargeon Oil Drop(Coulombs)

electron. (Accepted value:

1 112.32 128.63 144.74 161.05 128.56 192.67 144.78 161.0

Discussion

1. Marsden and Geiger used a zinc sulfidescreen to detect where the a particlescame out of the gold foil. This screenproduces a flash of light (a "scintillation")when an a particle strikes it. Observingthe scintillations with a microscrope, thev

a) When did the probe "bounce back" inyour experiment?

b) Suggest why some particles bouncedback, and many were deflectedslightly. (HINT: Remember that theparticle is positively charged.)

c) Most of your probes went throughwithout hitting anything. What doesthis imply about the size of the objectconcealed in the plasticene, in thecardboard, or under the shield?

d) Most of the a particles passeddirectly through the gold foil. Whatdoes this imply about the amount ofspace occupied by the stuff of atoms?

3. Draw a picture of a model for an atom thatwould explain the gold foil experiment.

4. Within two years of the a particlescattering experiments, Rutherford putforward the following description ofatoms.

An atom has a small central core inwhich all the positive charge and mostof the mass of the atom are concen-trated. The small positively charged coreforms the nucleus of the atom. Most

5.

of the atom is empty space in whichsufficient electrons to balance thecharge on the nucleus revolve.

Rutherford's atomic model has often beenlikened to a miniature solar system, theelectrons being "planets" revolving arounda nucleus "sun"a) Draw Rutherford models for the

following atoms: He(2 electrons),O(8 electrons), and CI(17 electrons).

Geiger and Marsden found they coulddetermine the number of positive chargesi n the nucleus from the angle of deflection .of the ce particles. Mosely, in 1913,confirmed their measurements using x-raydata. Mosely called the number of positivecharges in the nucleus the atomic number.(The symbol for atomic number is Z.)Further research has shown that thenucleus of an atom consists of particlescalled protons and neutrons. ( The symbolsfor protons and neutrons are p and n,respectively.) The proton has a relativecharge of 1 + and an approximate relativemass of 1. The neutron, discovered byChadwick in 1932, is neutral and has anapproximate relative mass of 1.

Radioactive Source

Zinc Sulfide Screen

Particle Location

e orbiting(electron) nucleusp i n

(proton) nucleusn i n

(neutron) nucleus

Since neutrons and protons are both found in,an atom's nucleus, they are collectively callednucleons. The total number of nucleons (thenumber of protons + the number of neutrons)is called the mass number. (The symbol formass number is A.) The mass number of themost abundant isotope of an element is equalto the whole number closest to the relativemass of the element.

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straight through the gold foil or weredeflected only slightly. A few particles,however, were turned through a largeangle, and occasionally an a particlebounced back. Rutherford said thisphenomenon was "almost as incredible asi f you had fired a 15-inch shell at a piece oftissue paper and it came back and hityou:'

Use Thomson's model of the atom todraw a picture of a very thin foil. Explainwhy, if you thought atoms looked likeThomson's model, the bouncing back of aparticles would be remarkable.

2. Compare your observations in thesimulation of the gold foil experiment withthose of Geiger and Marsden. How arethey the same? How are they different?

found that most of the particles went

For Further ReadingBrady, J.E. and J.R. Holum. Fundamentals of

Chemistry. New York: John Wiley and SonsInc., 1984.

Gray, H.B. and G.R Haight, Jr. Basic Principlesof Chemistry. New York: W.A. Benjamin Inc.,1967.

Magie, W.F., ed. A Source Book in Physics.Cambridge, Massachusetts: HarvardUniversity Press, 1969.

Paul, D. et al. The New Physics. Toronto: Holt,Rinehart and Winston of Canada Ltd., 1977.

Project Physics Incorporated. Models of theAtom. New York: Holt, Rinehart and WinstonI nc., 1968-69.

Toon, E.R. et al. Foundations of Chemistry.New York: Holt, Rinehart and Winston Inc.,1968.

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The Rutherford-Bohr Atom

Students should be able to:

1. Define an energy level.2. Explain the energy changes that occur when an electron

moves from a higher-to-lower or lower-to-higher energy levelwithin an atom.

3. State Bohr's findings regarding energy levels and the manneri n which electrons occupy these levels.

4. Illustrate with examples the fact that elements with similarelectron arrangements have similar chemical and physicalproperties.

Program DescriptionThe problems inherent in Rutherford's model of the atom wereresolved by Niels Bohr. He developed the hypothesis thatelectrons cannot possess just any amount of energy but ratheralways occupy specific energy levels, or "orbits". This programi ntroduces Bohr's hypothesis and goes on to explore the nature ofthe transfer of electrons from one specific level of energy toanother. Using animation, the program also demonstrates Bohr'shypothesis about the relationship between the potential andkinetic energy of electrons in circular and elliptical orbits. Alsodiscussed is the relationship between the arrangement ofelectrons in an atom and the chemical properties of that atom.This, in turn, is used to explain the similarities in the properties ofelements as grouped in the periodic table.

Before ViewingIt would be helpful to review with the students some of theunanswered questions Bohr was trying to explain with his modelof the atom.1. Why are atoms able to combine with other atoms?2. Why do some elements have very similar chemical properties

to those of other elements?3. Why is there an obvious order to the elements such that the

periodic table can be set up?4. How can atoms remain stable with electrons rotating around

the nucleus?5. How are line spectra explained?

Since most students have likely not observed line spectra before,Activity 1 should be completed before viewing the program. UseActivity 1 to establish a phenomenon which Rutherford's modelcould not explain. Point out that this was not a new problembut rather one that was not explained by any previous atomicmodel. The teacher may prefer to leave the calculations of theenergy of each of the hydrogen lines until after viewing. Thesecalculations will be helpful later on in Activity 2.

After ViewingStudents should carry out activities 2 to 5 after viewing theprogram. In Activity 2 the students think about Bohr's postulates.While the calculations in Part 3 of this activity are very limited,they are useful for the quantitative picture they provide for thehydrogen atom.

The concept of quantization is not always an easy one to grasp:Activity 3 provides a simple demonstration to clarify the term. Thefourth activity introduces the idea of a similarity in the atomicstructures of some atoms and ties this in with the similar chemicalproperties of the same elements. Bohr's model was successful inthat it provided a correlation between atomic structure, chemicalproperties, and the periodic table.

ActivitiesActivity 1: Line SpectraApparatusdiffraction grating spectroscope (see diagram

below)gas discharge tubes (helium, hydrogen,

mercury, etc.)high-voltage D.C. source to operate the above

tubeswhite light source (fluorescent ceiling lights)

Method1. Operation: Look through the diffraction

grating at the narrow end of thespectroscope pointing the slit towards theli ght source to be analyzed. The spectrum

Objectives

will appear on the right side of the slitbelow the scale where the wavelengths ofthe absorption or emission lines can beread. The visibility of the spectrum can beimproved by holding the hand around thenarrow end of the spectroscope using thethumb and the index finger to keep straylight from around the eye.

2. Calibration: The numbers on the scale ofthe spectroscope represent 100 nm units.Looking through the spectroscope at afluorescent lamp, all colors of thespectrum will be visible with two brighterlines-one violet line at 436 nrn and onegreen line at 546 nm. If the lines are notexactly in place, the difference must beadded or subtracted respectively in alldeterminations.

3. Observe the spectra of the gases in eachdischarge tube. Record the wavelengths ofthe lines for each gas.

Discussion1. Compare the wavelengths of the emitted

lines for each gas. Do any of the gaseshave the same set of lines?

2. Calculate the energy of each line inhydrogen's spectrum.

Diffraction Grating

Activity 2: Bohr's Postulates andQuantization1. Bohr's first postulate regarding the

structure of the atom is that an atompossesses stable states in which electrons will not lose energy by radiation evenif they are in motion about the nucleus ofthe atom. According to classical physics,an electron in orbit about a nucleus is acharged accelerating particle and thereforeshould lose energy by emitting radiation.

What was Bohr implying about elec-trons in atoms when he formulated thispostulate?

2. Bohr's second postulate regarding atomicstructure is that an electron can movefrom one stable state to another stablestate and when it does it will absorb oremit an amount of energy that is equal tothe difference in energy between the twostable states. In 1923, Herschel proposedthat elements could be identified by theline spectra through a prism.

How do you think Bohr correlated hispostulate with the existence of linespectra?

3. Particles moving in circular orbits ofradius r, with a speed v, and a mass m,have an angular momentum mvr. Bohr'sthird postulate was that atoms are onlyallowed to have orbits for which thevalues of mvr are definite. Using the threepostulates, Bohr was able to calculate thevalues of certain properties of the stablestates in hydrogen.

To get an idea of the values associatedwith certain properties of the stable orbitsi n hydrogen, use the following formulaederived by Bohr:a) Calculate the radius of the first three

orbits (n =1. 2. and 3) usina

b) Calculate the speed of the electron inthe first three orbits (n =1, 2, and 3)using

c) Calculate the energy of the electron inthe first three orbits (n=1, 2, and 3)using

d) When you observed hydrogen'sspectrum in Activity 1, the linesshould have had wavelengths of about656 nm, 486 nm, 434 nm, and 410 nm.These lines are the result of energy,-.being released when electronsdropped from the third to the second,fourth to the second, fifth to thesecond, and sixth to the secondenergy level respectively in thehydrogen atom. Use the formula inPart c to calculate the energy of theli ght that Bohr predicted would bereleased in each of the above energylevel changes. Compare this with ,,the value you calculated using theobserved lines in Activity 1.

Activity 3: A Demonstration of

Quantization

Apparatussmooth cardboard or bristol boardwood or books to provide support as shownball bearing

Method

( Height is variable.)

1. Cut and fold the cardboard as shown inthe diagram. (All dimensions are sugges-tions only.)

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. Roll the ball bearing along the bottom (flat) ACtivity 4: The Properties of thepart of the cardboard with enough energy Elementsthat it freely rolls about one quarter of theway up the incline before coming backdown.

3. Repeat step 2 several times, pushing theball bearing harder each time until afterfour or five pushes, the bearing is able toreach and stay at the top.

Bohr predicted the maximum number ofelectrons that could be found in differentenergy levels of atoms. He proposed that thephysical and chemical properties of anelement were related to the arrangement ofthe electrons in these energy levels. Inparticular the number of electrons in the lastpr outer level of the atom is related to itschemical properties. Why is it reasonable toassume this?

ObservationsThe ball bearing represents an electron in anatom.

Think of the bottom of the ramp as thel owest energy level or ground state of theatom and the flat top as the second or nextenergy level above the previous one. What isthe effect on the electron of giving it differentquantities of energy?

Discussion1. Describe in terms of energy what must be

done in order to move the marble from thebottom of the ramp to the box at the top.Also describe what happens if insufficientenergy is given to the marble.

2. I f we assume that the time to transferfrom bottom to top is so small it can't bemeasured, then the marble has only twomeasurable energies. What is the termused to describe such a situation?

3. I n what ways is the marble moving up (ordown) the ramp analogous to an electronin an atom?

Electron Arrangements in Alkali Metals

1. The following is a list of elements. Theatomic number of the elements is givenalong with the symbol. Using the abovetable as a guide, place electrons in theenergy levels and decide which elementscan be grouped together into families withsimilar chemical properties. Check theperiodic table to see if you are correct.

2. Describe the arrangement of electrons inthe energy levels of the followingelements: O (atomic number 8), S(16),Se(34), and Te(52). Explain why you wouldexpect these elements to have similarchemical properties.

3. Make a drawing of each of the noblegases showing the proper number ofelectrons in successive energy levels. Thenoble gases are He (atomic number 2),Ne(10), Ar(18), Kr(36) and Ke(54). Whatsimilarities do the inert gases show thatwould account for their similar chemicalproperties? What inert gas is theexception to the rule?

Further ReadingBrady, J.E. and J.R. Holum. Fundamentals of

Chemistry. New York: John Wiley and SonsInc., 1984.

Gray, H.B. and G.P Haight Jr. Basic Principlesof Chemistry. New York: WA. BenjaminI nc., 1967

Pauling, L. General Chemistry. San Francisco:W.H. Freeman and Company, 1970.

Project Physics Incorporated. Models of theAtom. New York: Holt, Rinehart and WinstonI nc., 1968-69.

Toon, E.R. et al. Foundations of Chemistry.New York: Holt, Rinehart and WinstonI nc., 1968.

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Element Energy levelwith number of electrons

1 2 3 4 5 6 7Li 2 1Na 2 8 1K 2 8 8 1Rb 2 8 18 8 1Cs 2 8 18 18 8 1Fr 2 8 18 32 18 8 1

1 0

Electron Arrangement"ObjectivesStudents should be able to:1. State the deficiencies in Bohr's theory.2. Describe the refinements in the atomic model proposed by

De Broglie, Schrodinger, and Heisenberg.3. Illustrate, using simple diagrams, the shapes of the s, p, and

d orbitals.4. Explain the meaning of the term "probability distribution".5. Show how the number of different orbitals, total number of

orbitals, and number of electrons in an orbital is related to thenumber (n) of the energy level.

6. Compare the arrangement of the orbitals in hydrogen with thati n multi-electron atoms.

7. Recognize the section of the periodic table to which anelement belongs, based on the knowledge of the last orbitalsto receive electrons in the element.

Program DescriptionThis program traces the work of De Broglie, Schrodinger, andHeisenberg in refining the Rutherford-Bohr model of the atom.Bohr's model could not be used to predict the probability of anelectron moving from one energy level to another, nor could itexplain the line spectra of multi-electron atoms. De Broglie, takinga different tack, demonstrated that electrons behaved as if theyhad wave properties. Schrodinger and Heisenberg combined DeBroglie's ideas with Rutherford's concept of a positively chargednucleus and Bohr's concept of energy levels. They arrived at atheory which proposed a) that electrons are not in a definite orbitaround the nucleus, but only have a probability of being found incertain locations about the nucleus; and b) that Bohr's energyl evels consist of a number of sub-levels, or orbitals. This programexplores the nature of these orbitals and the arrangement ofelectrons within them. It further demonstrates the importance ofthe number of electrons in the outer orbitals in determining thechemical properties of the atom.

Then use a chart, transparency or slide of several line spectra tocompare hydrogen's spectra with those of multi-electron atoms.Use this comparison to demonstrate that although Bohr's modelaccounted for the fine spectra of hydrogen, it could not accountforvariations in the line spectra of atoms with more than oneelectron.

After ViewingBoth Activities 1 and 2 should be done after viewing Program 3since they are based on the work of Schrodinger and Heisenberg.Activity 1 uses a simulation to establish the meaning andi nterpretation of a probability distribution curve. Activity 2 providespractice with the terminology and concepts of this program-orbitals, electron configuration, and their relationship to thechemical properties of atoms.

ActivitiesActivity 1: Electron LocationI n this activity you will simulate an electrondistribution by attempting to throw penniesi nto a film can. The distribution of the penniesonce they have all been thrown is analogousto a "snapshot" of an electron. The film canrepresents the nucleus.

Materialspennies (or bingo chips, corn kernels, etc.)a film can1 m2 sheet of paperstringfelt-tip markertape

Procedure

Before ViewingReview the ways in which Bohr's model of atomic structure wassuccessful. Also review the terminology used in his explanationsof atomic behavior: orbit, energy level, quantization, line spectra.

1. Set the film can in the centre of the paperand trace a circle around the can. Set thecan aside temporarily.

2. Put one end of the string at the centre ofthe circle.

a) Compare your penny probabilitydistribution with the 1S- and2S-electron probability distributions.How are they the same? How arethey different?

b) Describe what happens to theprobability of finding an electron asdistance from the centre of thenucleus increases.

The above electron probability distribu-tions come from the solutions toSchrodinger's Equation. His equationand Bohr's method produced the samevalues for hydrogen. Why wasSchrodinger's model considered better?By arbitrarily selecting a certainprobability- e.g., 95%-we can deter-mine a maximum distance from thenucleus within which there is a 95%probability of finding the electron. Whenwe do this for a 1S-electron atom we geta sphere within which there is a 95%chance of finding the electrons. Describethe variations of the probability of findingthe electron within that sphere.Similarly, we could draw a sphere of 95%probability for 2S electrons. Such asphere is called the 2S orbital. Describethe variations in the probability of findingthe electron within that sphere.How are the various s orbitals the same?Different?The sphere is the typical shape of the sorbital. Draw the shapes of the three porbitals. Don't forget to show theirorientation.

9. What is meant by the term orbital?10. You may have been lucky enough to

actually toss a penny into the can. Thiscorresponds to an electron being foundi n the nucleus. What evidence suggeststhat the probability of finding an electronin the nucleus is not zero?

Activity 2: Electron ArrangementThe following questions will help you reviewthe terminology and concepts of electronarrangement.1. The movement of electric current through

a wire produces a magnetic field aroundthe wire. Similarly, spinning electrons havebeen found to behave as tiny magnetswith theirown magnetic field. Thesemagnetic fields have been found to pointi n one of two directions.a) Explain what forces would make it

difficult for two electrons with thesame spin to occupy the same orbital.

b) Why is it possible for two electronswith opposite spins to occupy thesame orbital?

c) Explain why it would be unlikely thatthree or more electrons might occupythe same orbital.

2. a) If a particular energy level is denotedby n (where n = 1, 2, 3...), what is thenumber of types of orbitals, totalnumber of orbitals and maximumnumber of electrons that may befound in that energy level (in termsof n)?

b) Complete the following table for thefirst four energy levels.

Energy Types of Total NumberLevel Orbitals Number of of

(n) Present of Orbitals Electrons1234

3. Hold the felt marker 5 cm along the stri ngand drawa circle.

4. Repeat step 3 holding the marker at 10, 15,20...50 cm to draw concentric circles onthe paper.

5. Tape the film can to the centre of thepaper (where the circle was originallydrawn around it).

6. Divide the pennies amongst four studentsso that each has approximately 25 4.pennies.

7. One student should stand at each side ofthe paper and each should try to toss thepennies into the can.

8. When all pennies have been tossed, countthe number of pennies in the 0-5 cm circle 5.(include any in the film can), the 5-10 cmcircle, 10-15 cm circle, etc.

9. Plot the number of pennies in each circleon a graph. Plot the distance, the 'i ndependent variable, along the horizontalaxis. Plot the number of pennies, thedependent variable, along the vertical axis.Join the points to form a smooth curve.The graph you have drawn is called aprobability distribution graph. 6.

Discussion

1. Describe the shape of the curve you havedrawn.

2. Describe what happens to the probability 7.(chance) of finding a penny as distancefrom the centre increases. 8.

3. The following graph is the probabilitydistribution graph for a 1S-electron and a2S-electron atom.

12

3. a) I n simple terms, explain why theorbitals of a particular energy level arefound at different energies in allatoms with more than one electronwhile the same orbitals all have thesame energy in hydrogen atoms.

b) How do the energies of the s, p, and dorbitals compare in any one level?

4. a) What is meant by the "electronconfiguration" of an atom?

b) Using a diagram similar to thefollowing, draw the electron config-uration for each of the followingelements; Al, Si, P, S, CI, Ar, K, Caand Sc.

5. Because the outer regions of atoms arethe first parts to come in contact whenchemical reactions take place, we mightexpect that all elements with similar outershells would have similar chemicalproperties.a) Draw an outline of the periodic table

as shown below.

b)

c)

d)

e)

c) Even though the orbitals are not atthe same energy as predicted by Bohr,the electron configurations of some ofthe elements in 2b above would befound to be the same as Bohr wouldhave found. Which elements wouldhave different electron configurationsthan Bohr's predictions because of adifferent order in filling the orbitalswith electrons?

Count the number of columns in eachof the marked-off areas (A), (B), (C),and (D) of the periodic table.What are the types of orbitals (s, p, d,or f) that are the last to receiveelectrons in each of the areas (A), (B),( C), and (D)?What is the maximum number ofelectrons that can occupy each of thes, p, d, and f sets of orbitals? Howdoes your answer compare with thecount made in Part a above? Explainthe connection between these twoanswers.What is the same (in most cases)about the electron configurations ofthe elements found in each of thecolumns shown on the above diagramof the periodic table? What would youexpect to find if you checked thechemical properties of these sameelements? Explain your answer.

Further ReadingBrady, J.E. and J.R. Holum. Fundamentals of

Chemistry. New York: John Wiley and SonsInc., 1984.

Gray, H.B. and G.P. Haight, Jr. Basic Principlesof Chemistry. New York: W.A. BenjaminInc., 1967.

Pauling, L. General Chemistry. San Francisco:W.H. Freeman and Company, 1970.

Pauling, L. The Nature of the Chemical Bond.New York: Cornell University Press, 1967.

Project Physics Incorporated. Models of theAtom. New York: Holt, Rinehart andWinston Inc., 1968-69.

Sharpe, S.W. Atoms and Matter. Toronto:John Wiley and Sons Canada Ltd., 1978.

Toon, E.R. et al. Foundations of Chemistry.New York: Holt, Rinehart and WinstonI nc., 1968.

How Atoms BondObjectivesStudents should be able to:

1. Show by means of simple diagrams that forces of attractionand repulsion exist between two closely spaced atoms.

2. Describe how the distance between two atoms influences thestrength of the attractive and repulsive forces between them.

3. Explain how overlapping orbitals allow two atoms to share thesame electrons.

4. Define a covalent bond.5. Illustrate, using a simple orbital energy level diagram, why the

elements of the same period of the periodic table have adecreasing volurne.

6. Explain how it is possible for one atom to pull on the electronof another atom with sufficient force to remove the electron.

7. Define an ionic bond.

better to leave discussion questions 2 to 4 of this activity untilafter the program.

After ViewingDiscussion questions 2 to 4 of Activity 1 should be completedafter viewing the program. Students will know of the various forcesthat affect atoms in close proximity and should be able to makethe connection with the behavior of the magnets used in the firstpart of this activity.

Activity 2 should be done after viewing the program. Thisactivity involves trends in the sizes of atoms and ions in the rowsof the periodic table as nuclear charges increase. The fact thatsome atoms "lose" electrons to other atoms is discussed in thisactivity. This follows up on the program's presentation of theformation of ionic bonds.

Program DescriptionThis program shows how the atomic structure elucidated bySchrodinger and Heisenberg finally explains the phenomenon ofatomic bonding. The program discusses the forces of attractionand repulsion within an atom and then demonstrates the ways inwhich these forces interact when atoms are in close proximity toone another. The two basic types of atomic bond are explained: 1)covalent bonds, which are formed when the orbitals of two atomsoverlap allowing them to share electrons; and 2) ionic bonds,which are formed when an electron is transferred from one atomto another creating two ions of opposite charges which are thenattracted and pulled toward each other. The program also explainswhy some types of atoms will not form bonds. It further points outthat atoms may also join in combinations of covalent and ionicbonds-a subject that will be explored further in Program 5.

Before ViewingHave the students do Activity 1 up to the end of the firstdiscussion question before viewing the program. This activityrequires the students to measure and then graph the effect ofdistance on attractive and repulsive forces. The activity isstraightforward and will give the students some experience withthe terms introduced at the beginning of the program. It would be

Activities

Activity 1: The Effect of Distanceon Attractive and RepulsiveForces

Part A: Distance and Repulsive Forces

Apparatus2 ceramic magnets approximately 2.5 cm in

diameter and with a 1 cm in diameter hole inthe middle

wooden dowelling approximately 45 cm inl ength

2 retort rods with stands2 "finger" clampstube of stopcock grease or other suitable

greasemasking tape

Method1. Set up the apparatus as shown in the

following diagram. Tape one of the ceramic

1 3

1 4

magnets to the wooden dowel, approxi-mately 10 cm from one end of the dowel.Mark off the dowel at 0.5 cm intervals,starting at the edge of the taped magnet.Spread grease along the dowel from theedge of the taped magnet to the pointwhere the clamp is holding the dowel. Thesecond magnet is free to move along thedowel. This second magnet must beplaced so that there is repulsion betweenit and the taped magnet. If they attract,remove the free magnet and reverse it onthe dowel.

2. Carefully move the free magnet towardthe fixed magnet until there is enoughrepulsive force to cause the free magnetto be just pushed away when it isreleased. Using the markings on thedowel, measure the distance between thetwo magnets.

Once this point has been determined,the object of the experiment is to movethe free magnet closer to the fixedmagnet, note the distance between them,then release the free magnet and note howfar it is pushed along the dowel by therepulsive force between the two magnets.Repeat this four times at each setting. Thedata is recorded in a table similar to the

following. (A set of experimental data arei ncluded in brackets to show the type ofresults that students may be expected toobtain.)

DistancebetweenMagnets

at Start (cm)

Part B: Distance and Attractive ForcesApparatus2 ceramic magnets as in Part Amasking tape1 spring balance that reads as small a weight

as possible1 ruler

Distance Travelled byFree Magnet (cm)

Method1. Tape one of the ceramic magnets to a

counter top.2. Tape the second magnet to the hook on

the spring balance after ensuring that thetwo magnets will attract each otherwhenthe one taped to the balance is loweredonto the one taped to the countertop.See the following diagram.

3. After taping the magnet to the springbalance, remove it from the vicinity of theother magnet and adjust the scale (if thisi s possible) so that it reads 0 g with themagnet attached. If this is not possible,note the weight reading with the magnetand tape in place.

4. The following measurements are tricky sobe patient and proceed slowly. Lower themagnet on the balance down toward themagnet taped to the counter top. As theforce of attraction increases the balance-magnet will be pulled downward. Soon theforce will be strong enough to quickly pull

Trial Trial Trial Trial Trial Average1 2 3 4 5

(3.0) (0)(2.5) •. (0.5)(2.0) (1.4)(1.5) (2.8)(1.0) (4.8)(0.5) (7.5)(0) (16.6)

the magnets together.Determine the pull (in grams) required to

separate the two magnets. Watch thescale carefully as the top magnet willrelease quickly. Repeat this four times andrecord the data in the following table.(Subtract the original weight of the magneteach time if necessary.)

5. Using the ruler, place the top magnet0.5 cm above the lower magnet. Keep themagnet in position by holding it betweenyour fingers. Pull up the spring balanceuntil it feels as if the magnet will stay atthis position. Release the magnet and readthe scale. Repeat four times and recordyour results in a table such as the onebelow. (Subtract the original weight of themagnet if necessary.)

6. Repeat step 5, placing the magnet anextra 0.5 cm above the lower magnet untilthere is no more force of attraction.Record your results in the table. (A set ofexperimental data are included in bracketsto show the type of results students maybe expected to obtain.)

a) Which pan`. of this curve representsthe attraction between the twohydrogen atoms? Explain.

b) Which part of this curve representsthe repulsion between the twohydrogen atoms? Explain.

c) What does the low point of the curverepresent? Explain.

d) What would be needed in terms ofenergy to increase the distancebetween the hydrogen atoms?

1. Using the data in your tables, plot twocurves on one graph set up as follows.

The distance travelled by the free magneti n Part A is not a force quantity, but it wascaused by a repulsive force. Therefore thePart A-curve above the horizontal axisshows graphically the effect of distanceon repulsive forces.Since repulsive and attractive forces areopposites, and since positive values areused to represent the repulsive forces,then the attractive forces will berepresented by negative values. Therefore,by plotting the "maximum pull" valuesfrom Part B as negative numbers andplacing the appropriate points below thehorizontal axis, this curve shows theeffects of distance on attractive forces.The energy changes that occur when twohydrogen atorns approach each other isshown in the following diagram:

Activity 2: Atomic and Ionic Sizes1. The size of an atom is related to its

chemical properties. The following datashow the atomic radii of the elements inthe third period of the periodic table.

1 5

4.

2.

3.

Distancebetween Maximum Pull Required toMagnets Keep the Magnets at the

at Start (cm) Indicated Distance (g)Trial Trial Trial Trial Trial Average

1 2 3 4 5(0) (6.75)(0.5) (3.25)(1.0) (1.5)(1.5) (0.75)(2.0) (0.35)(2.5) (0.25)(3.0) (0.15)(3.5) (0)

Na 1.86 11Mg 1.60 12AI 1.43 13Si 1.17 14P 1.10 15S 1.04 16CI 0.99 17Ar 0.94 18

Molecular Substances andCovalent CrystalsObjectivesThe students should be able to:

1. Show that the noble gases have similar electronarrangements.

2. Illustrate, using electron structures, how selected elementsform diatomic molecules.

3. Recognize that the electron structures of selected atoms thathave formed bonds are like the electron structures of thenoble gases.

4. Demonstrate, using electron structures, how single, double,and triple covalent bonds can form in selected diatomicmolecules.

5. Describe an experiment that illustrates the different strengthsof single, double, and triple covalent bonds.

6. Define a polar molecule.7. Show how the structure of a molecule can be used to explain

and predict someof its properties.8. Describe why carbon is capable of forming such a large

number of molecules compared to most other elements.9. Understand how a covalent crystal forms.

Program DescriptionThis program builds on the students' understanding of covalentand ionic bonding. The formation of diatomic molecules isexplained and it is shown how the electron arrangements of atomswithin these molecules resemble the electron arrangements of thenoble gases. The program goes on to discuss single, double, andtriple covalent bonds and to demonstrate the strength of each ofthese three types. It then introduces polar molecules-moleculeswhich combine both covalent and ionic bonding-and examinesthe strength of the bonds formed between such molecules.Finally, the special bonding properties of carbon are explored andthe structure of a covalent crystal-in this case, diamond-isexplained.

Before ViewingReview the electron configurations of the inert gases and pointout their similarities. Establish the fact that these are the onlyelements whose outer s and p orbitals are filled, and the only onesthat show virtually no tendency to combine with other elements.You could introduce the Octet Rule as well. The fact thatrepresentative elements tend to lose or gain electrons until theyhave eight in their outer shell will be helpful in understanding thebonding situations portrayed in the program. All the activities arebest done after viewing.

After ViewingActivity 1 gives the students practice at predicting how theelectron configuration of an element will change in order to satisfythe Octet Rule. Activity 2 is a teacher demonstration of the polarproperties of two different substances. In particular, this activityi ntroduces the fact that there are intermolecular forces set up as aresult of polarity. Finally, Activity 3 has the students constructdifferent carbon-containing molecules. This activity stresses thesomewhat unique ability of carbon to form up to four bonds withother elements, resulting in many different types of molecule.

ActivitiesActivity 1: Molecule formationThere are definite patterns that appear in theperiodic table with respect to the way atomsalter their electron configuration when they arebrought close to other atoms:a) The inert gases show virtually no

tendency to change their electronconfigurations from the ns2np6 arrangement in the outer levels. (He has 1 s2 as itsconfiguration and is equally stable.)

b) Elements in the first four families (orcolumns) on the left side of the periodictable and those in the five families

1 7

16

a)

b)

2. a)

b) What do you notice about theelements of the third period withrespect to the energy level that isbeing filled?

c) What is happening to the nuclearcharge on the elements of this periodas we move from left to right acrossthe period?

d) Explain the trend noted in the graphi n Question 1b by considering youranswers to Questions 2b and 2c.

3. The following table shows the sizes of thei ons of the same period-three elements.

Plot a graph of atomic number againstatomic radius. Put the atomic numberon the horizontal axis and atomicradius on the vertical axis.Describe the trend in atomic sizegoing from left to right across theperiod.Using diagrams similar to thefollowing, show the electron con-figurations of the elements of thethird period.

a) On the same graph used to plot theatomic number against atomic radiusi n Question 1a, plot another graphshowing atomic number against ionicradius.

b) Describe the trend in ionic size goingfrom left to right across the period.What important differences do younotice in the trend of atomic radiiwhen you compare this to the trend inionic radii?

c) Considering the electron configura-tions, atomic radii, and nuclearcharges of the elements as discussedin Question 2, explain why theelements near the end (or right side)of the period would be more likely toform negatively charged ions. Whywould the elements at the start (or leftside) of the period be more likely toform positively charged ions?

d) Explain the fact that i) positivelycharged ions are smaller than theneutral atoms from which they formedand ii) negatively charged ions arel arger than the neutral atoms fromwhich they are formed.

e) Imagine that magnesium atoms andchlorine atoms are approaching eachother in a container. What changesare likely to occur in their electronconfigurations as they react toproduce magnesium chloride?

Further ReadingBrady, J_E. and J.R. Holum. Fundamentals of

Chemistry. New York: John Wiley and SonsInc., 1984.

Gray, H.B. and G.P Haight, Jr. Basic Principlesof Chemistry. New York: W.A. BenjaminInc., 1967.

Madras, S. et al. Basic Modern Chemistry(3rd ed.). Toronto: McGraw-Hill RyersonLtd., 1978.

Pauling, L. General Chemistry. San Francisco:W.H. Freeman and Company,1970.

Pauling, L. The Nature of the Chemical Bond.New York: Cornell University Press, 1967.

Sharpe, S.W. Atoms and Matter. Toronto:John Wiley and Sons Canada Ltd., 1978.

Toon, E.R. et al. Foundations of Chemistry. NewYork: Holt, Rinehart and Winston Inc., 1968.

N a 0.95 + 1Mg 0.65 + 2AI 0.50 + 3Si 0.41 + 4P 2.12 -3S 1.84 -2CI 1.81 -1Ar does not apply none

1 8

preceding the inert gases on the right sideof the table tend to lose or gain electrons(to a maximum of four) so that they end upwith an electron configuration like that ofthe nearest inert gas.

c) The atoms of elements that attempt toremove electrons from other atoms willbond with each other by sharing enoughelectrons to resemble the nearest inertgas in their electron configuration.

1. Complete the following table showing theexpected loss or gain of electrons by eachelement during bond formation.

Two examples from the program are:

3. Consider the examples shown in Question2, NaCl and 02. One of them, NaCI,involves the formation of charges on theatoms because of the transfer of an

For the molecules you worked out inQuestion 2:

Activity 2: A Demonstration ofMolecular Polarity

Part A

Apparatus1 overhead pen1 overhead projector

Try the following combinations:

(i) magnesium with oxygen(ii) carbon with oxygen(iii) phosphorus with chlorine(iv) fluorine with chlorine(v) lithium with sulfur(vi) scandium with fluorine(vii) nitrogen with nitrogen

a) I ndicate appropriate charges on theatoms involved in bond formation byelectron-transfer.

b) Choose one molecule whose bondingmight be explained by either electron-sharing or electron-transfer. Explainyour choice.

4 circular ceramic magnets4 pieces of clear colorless acetate(3 cm x2 cm)

Method(Steps 1 and 2 must be prepared carefully bythe teacher before class.)1. Take two of the ceramic magnets and

place them side by side such that theyattract each other. If they repel, turn one ofthe magnets over. Tape a piece of acetateon each magnet so that half is on and halfhangs over the edge of the magnet.

Place a"+" sign on the protruding. part ofone of the acetates and a "-" sign on theother. The choice of magnets is arbitrary.

2. Repeat step 1 using the other two ceramicmagnets. Extra care must be taken toensure that the magnet that is marked"+" repels the one marked "+" from step1 above.

3. Place a " + " and "-" magnet on theoverhead projector. Slowly move one ofthe magnets toward the other until theypull together and bond. These representtwo bonded atoms forming a polarmolecule (with unequal electrondistribution).

4. Repeat step 3 with the remaining twomagnets on a different part of theprojector surface.

Element Electron Nearest Change inConfiguration Inert Gas Electron

ConfigurationNe l ose the 2-3s

electronsMg

BKSLiFCOCaN

2. Show the molecules that would beexpected to form if atoms of the indicatedelements were combined. Use the formatshown in the program.

A rectangle indicates overlappingorbitals involved in electronsharing.

An arrow shows electron transterfrom one orbital to another.

a)

b)

electron. It could be shown as

movement of electrons from one atom tothe other and so both atoms remainneutral.

The other molecule, did not involve

5.

6. With the "molecules" situated as shownbelow, slowly move one of the "molecul-es" sideways towards the other. Describethe effect of the "molecules" on eachother.

Discussion

1. I n what situation would polar moleculesrepel each other? In what situation wouldthey attract each other?

2. Whenever atoms bond with each other toform a molecule, the strength of the bondbetween the atoms is quite good whetherit is an ionic or covalent type of bond. Ineach of the following situations, describe

With the "molecules" situated as shownbelow, slowly move one of the "molecul-es" sideways towards the other. Describethe effect of the "molecules" on eachother.

the strength of the bond between the polarmolecules (i.e., the intermolecular force).

When A and B bonded,electron transfer occurred. Ahas a strong positive and Bhas a strong negative charge.

Part B

Apparatus

When C and D bonded,electron sharing occurred,but D was slightly moreelectronegative than C. Chas a slight (6) or partialpositive charge and D aslight negative charge.

When E and F bonded,electron sharing occurredand E and F had the sameelectronegativity.

carbon disulfide and bring the glass rodnear (but not touching) the stream of fluid.Observe any effect and close thestopcock.

3. Immediately open the stopcock of theburette with the water and bring the glassrod near the stream of water. Observe anyeffect and close the stopcock.

4. Repeat steps 2 and 3 after rubbing theebonite rod with cat's fur.

5. I f students are unfamiliar with staticelectricity, suspend one of the chargedrods using tape and string and bring theother rod close to it. Then bring a secondrod of the same substance and rubbedwith the same material near the first rod.Observe the effect of the charged rods oneach other.

Discussion1. Which liquid contains polar molecules and

which contains non-polar molecules?2. a) What can be said about the charges

on the glass and ebony rods after theywere rubbed with the indicatedmaterials?

b) Explain why the polar molecules wereattracted to both types of electriccharges.

3. When liquids are heated to their boilingpoint, the molecules separate from eachother and there is a change of state fromliquid to gas. The boiling point of a liquidi s an indication of the intermolecular bondstrengths. Which liquid, water or carbondisulfide, would you expect to have thehigher boiling point? Explain your answer.

1 retort rod2 burettes50 mL of carbon disulfide50 mL of distilled water1 glass rod1 piece of silk1 ebonite rod1 piece of cat's fur1 burette clamp (to hold 2 burettes)2 250 mL beakers

Method1. Fill one of the burettes with 50 mL of

water and the second with 50 mL ofcarbon disulfide. Place the burettes besideeach other in the burette clamp.

2. Rub the glass rod with the piece of silk toestablish a static charge on the rod. Openthe stopcock of the burette containing the

Activity 3: Bonding InvolvingCarbon AtomsApparatus1 student atomic model kit using wooden

spheres for atoms and springs for bonds

1 9

a)

b)

c)

20

Method

1. Choose six carbon atoms (usually theblack spheres) from the kit.

2. Using all six carbon atoms, but as manysprings (bonds) and hydrogen atoms(usually the white spheres) as needed,build molecular models, sketch them, andwrite the formula for each. You do not haveto use the same number (or any for thatmatter) of the hydrogen atoms each time amolecule is built, but all six carbon atomsmust be used.

3. Pool the results of all members of theclass.

Discussion

1. What is the electron configuration forcarbon?

2. Explain why carbon atoms are capable offorming four bonds when their electronconfiguration suggests that only twobonds are possible.

3. How many different molecules was theclass able to build using six carbonatoms? Some of the formulae may bethe same, but the molecules could bedifferent. Check the diagrams of themolecules carefully before deciding onthe number of different molecules.

4. How is carbon unique among theelements?

5. When building the molecules, it ispossible to start one using only carbonatoms. If there were enough carbon atoms,you could satisfy the bonding require-ments of the atoms. What is the namegiven to the substance you started to buildwith only carbon atoms?

Further ReadingBrady, J.E. and J.R. Holum. Fundamentals of

Chemistry. New York: John Wiley andSons Inc., 1984.

Gray, H.B. and G.P. Haight, Jr. Basic Principlesof Chemistry. New York: W.A. Benajmin Inc.,1967.

Madras. S. et al. Basic Modern Chemistry(3rd ed.). Toronto: McGraw-Hill Ryerson Ltd.,1978.

Pauling, L. General Chemistry. San Francisco:W.H. Freeman and Company, 1970.

Pauling, L. The Nature of the Chemical Bond.New York: Cornell University Press, 1967.

Sharpe, S.W. Atoms and Matter. Toronto: JohnWiley and Sons Canada Ltd., 1978.

Toon, E.R. et al. Foundations of Chemistry.New York: Holt, Rinehart and Winston Inc.,1968.

-Metals and Ionic SolidsObjectivesThe students should be able to:

1. Show where the metallic elements are located on a periodictable.

2. Describe the feature of the electron arrangement that iscommon to all metals.

3. Explain what is rneant by the fact that s orbitals lack a spatialorientation.

4. Describe the metallic bond.5. Illustrate, using simple diagrams, how metallic bonding can

explain electrical conductivity, thermal conductivity, malleabil-ity, and ductility.

6. Explain why ionic solids are a type of network solid.7. Understand why ionic solids are not conductors of electricity

or heat, have high melting points, are hard, and split if they arehit by another hard object.

8. Show how a knowledge of electron arrangements and bondingcan help to predict information about substances.

Program DescriptionThis program builds further on the students' understanding ofelectron arrangement and the resulting possible types of bondingby considering the different properties of metals and ionic solids.I t demonstrates how the properties which all metals have incommon-electrical conductivity, thermal conductivity, malleabil-i ty, and ductility-can be explained by the fact that all metallicelements have s orbitals as the outer orbital. The type ofnon-directional covalent bond formed by metal atoms iscontrasted to the type of bond in covalent crystals (discussed inProgram 5) and to the type of bond in ionic crystals. Ionic solidsdo not conduct either electricity or heat and they have a very highmelting point. The program shows how these common propertiesalso can be explained by the electron configurations of the atomsand the resulting forces of attraction and repulsion. In summary,the program points out how knowledge of electron arrangementand bonding allows us to predict a great deal of information aboutsubstances without having to test each one.

Before ViewingIn order to better acquaint the students with the properties ofmetals, Activity 1 should be done before viewing the program. It issuggested that the discussion part of this activity be left until afterthe program since it involves questions that refer to the model ofmetallic bonding described in the program. The remainingactivities are all best left until after the program.

After ViewingThe discussion part of Activity 1 should now be completed usingthe model of metallic bonding shown in the program. Activity 2 isan attempt to relate the physical properties of metals to theirelectron configurations. Activity 3 is a teacher demonstrationdesigned to show why ionic crystals shatter when hit, whilemetals bend or assume a new shape. Activity 4 has the studentspredict the relative bond strengths of two ionic compounds.

ActivitiesActivity 1: The Properties ofMetals

Apparatusstrips of the following metals, approximately

2.5 cm x12 cmcopperzinci ronaluminumnickel

crucible tongs1 piece of lead caning (1 m long) used in

making stained glass windows2 pairs of pliersfine emery cloth1.5V dry cell1 set of electrical leads1 ammeterBunsen burner

21

22

_ small piece of waxhammerany scrap piece of wood to be used as a

hammering surfacewhite light source (such as an overhead

projector)

Method1. Lustre or reflectivity: Clean the surface of

each of the five metal strips using the fineemery cloth. Hold the clean surface in abeam of white light and observe. Reflectthe light onto a piece of white paper. Howdoes the reflected light look?

2. Conductivity: Fasten the metal strip tothe positive terminal of the ammeter bytightening the threaded knob on theterminal. Clip the end of one of electricalleads to the other end of the metal strip.Fasten one end of the other lead to thenegative terminal of the ammeter. Hold theend of the lead connected to the negativeterminal of the ammeter on the negativeend of the dry cell. Momentarily touch theend of the other lead to the positive end ofthe dry cell. Do not leave the circuit closedas this could damage the ammeter. (Seethe following diagram.)

Observe the needle on the ammeter whenthe circuit is closed. Repeat the above foreach of the other metal strips.

3. Heat Conductivity: Place one of the piecesof wax on the end of one of the metalstrips. Carefully hold the strip with thecrucible tongs so that approximately 1 cmof the strip is in the flame of the Bunsenburner. Observe the wax as the metal stripi s heated. Repeat for each of the othermetal strips.

4. Ductility: Have a student hold one end ofthe 1 m long lead caning tightly with a pairof pliers. Have another student hold theother end tightly with the other pair ofpliers. Have the students pull the lead inopposite directions. Observe any changes.Measure the length of the lead after thepulling is completed.

5. Malleability: Place one end of the leadcaning used above on the scrap piece ofwood. Strike the piece of lead firmlyseveral times with the hammer. Observeany changes.

Discussion1. List the properties of metals based on the

observations you made when this activitywas performed.

2. The model of metallic bonding describedi n this program is:a) Metallic crystals consist of fixed

positive charges (the nuclei) sur-rounded by a "sea" of negativelycharged electrons.

b) The delocalized electrons move freelythroughout the metal but are heldwithin the metal by the attractions ofthe nuclei.

c) The nuclei are bonded together by theattraction for the electrons as theymove among them.

Using the model outlined above, explain eachof the five metallic properties listed inQuestion 1.

Activity 2: The Metallic BondMetals, as a general rule, have several orbitalsi n their outer shell but very few electrons

occupying them in the neutral atom. Theavailable empty orbitals can be occupied bythe freely-moving electrons within the metalliccrystal.1. Show the electron configurations of Li, Be,

Ca, and Mg. In each case give the numberof orbitals in the outer shell that containelectrons and the number that areavailable to be occupied by the electronsthat move in from other atoms.

2. It is practically impossible to stategeneralizations about the melting pointsof all the metals in the periodic table.However, it is not too difficult to compareone metal with another and explain theirdifferent melting points.a) Assuming that the outer shell

electrons all become part of the "sea"of electrons in a metallic crystal,explain why calcium has a highermelting point than potassium (838°Cfor Ca compared to 64°C for K). Showthe electron configuration of eachatom.

b) The melting points of the alkali metalsdecrease as you go down the family.Explain why. Show the electronconfigurations for Li, Na, and K.

3. a) The ability to conduct both electricityand heat increases from sodium tomagnesium to aluminum in the thirdperiod of the periodic table. Show theelectron configurations of each ofthese elements and explain these twotrends.

b) I t has been found that the ability of ametal to conduct electricity decreasesif the temperature is increased. Usingthe model of the metallic bondpresented in this program, explain thisobservation.

Activity 3: Ionic CrystalsApparatusapproximately 10 ceramic magnets (2.5 cm in

diameter)

clear acetate sheet1 small (15 cm) plastic rulertapeoverhead projectoroverhead pen

Method

1. Cut small pieces of acetate, large enoughto fit on the top of each of the ceramicmagnets. Tape one of the pieces ofacetate to one side of a ceramic magnetsuch that the hole is covered.

2. Place a "+ " sign on the acetate "window"on the magnet so that it is visible throughthe hole.

3.

4.

Place a second magnet beside the onefrom above such that the magnets attracteach other. (If they repel, turn the secondmagnet over.) After making sure themagnets attract, tape another acetatepiece onto the second magnet and place a"-" sign on the window.Repeat step 3 until there are five magnetsi n a row, all attracting, but with alternate"+" and "-" signs.

6. Set the small plastic ruler on its edge andpush it between the two rows of magnets.The ruler will be held in place by themagnets. Place the arrangement on theoverhead projector.

7. Hold one of the rows of magnets steadyand apply a gentle push, directed alongthe row, to the end magnet of the otherrow. Observe the effect on the "crystal" asthe "ions" move.

Hold steady.

Discussion

1. a) Place "+" and"-" signs in theappropriate places on the followingbefore and after diagrams of a crystalthat has had its ions displaced.

2. In an actual situation, when a crystal isstruck it usually shatters. Explain why thishappens.

3. a) The following before and afterdiagrams represent the atoms of ametal when applied force causes therow of atoms to shift. Place "+ "(nuclei) and "-" (electrons) in theappropriate places on both diagrams.

5. Place a second row of magnets beside thefirst row. Once again make sure that all themagnets attract. When completed, the tenmagnets should be arranged as follows.

Before After

b) Referring to the above diagrams,explain the observations noted in thedemonstration.

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b) Explain why metals end up simply -changing shape rather than shatter-ing like ionic crystals when a force isapplied to them.

Activity 4: Bond Strength in IonicCompoundsIt is generally recognized that the meltingpoint of a substance is an indication of thestrength of the bonds holding the atoms of anionic substance together. The strength of thebond between the atoms is determined by theradii and charges on the ions. The followingdiagrams show the relative size and placementof the ions in sodium fluoride and sodiumi odide.

The melting points of Nal and NaF areapproximately 660°C and 990°C respectively.Explain why the melting point of Nal is lower.(The Na ions are all +1, and F and I ions areall -1.)

Further ReadingBrady, J.E. and J.R. Holum. Fundamentals of

Chemistry. New York: John Wiley and SonsI nc., 1984.

Gray, H.B. and G.P Haight, Jr. Basic Principlesof Chemistry. New York: W.A. Benjamin Inc.,1967.

Madras, S. et al. Basic Modern Chemistry(3rd ed.). Toronto: McGraw-Hill Ryerson Ltd.,1978.

Pauling, L. General Chemistry. San Francisco:W.H. Freeman and Company, 1970.

Pauling, L. The Nature of the Chemical Bond.New York: Cornell University Press, 1967.

Sharpe, S. W. Atoms and Matter. Toronto:John Wiley and Sons Canada Ltd., 1978.

Toon, E.R. et al. Foundations of Chemistry.New York: Holt, Rinehart and Winston Inc.,1968.

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Videotapes BPNProgram 1: Introducing the Players 252501Program 2:TheRutherford-Bohr Atom 252502Program 3: Electron Arrangement 252503Program 4: How Atoms Bond 252504Program 5: Molecular Substances and

Coralent Crystals 252505Program 6: Metals and Ionic Solids 252506

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