Author Institution: Department of Chemistry, University of Minnesota, Minneapolis, Minnesot

Brasted, Robert C.

The Ohio Journal of Science

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No. 1 TEACHING GENERAL CHEMISTRY 15ATOMIC ARCHITECTURE FOR THE GENERALCOLLEGE CHEMISTRY STUDENTROBERT C. BRASTEDDepartment of Chemistry, University of Minnesota, Minneapolis, MinnesotaThe approach used in presenting atomic structure to the first year collegechemistry student is predicated upon several factors: 1. a very early presentationof the subject, after introduction of principles of units (fundamental and derived)and stressing of dimensions in the realm of the ultraminute; 2. a reasonablycomplete picture at this early stage, on the assumption that a usable working modelwill be developed; 3. an attempt to capture interest in the chemical area ratherthan devoting time at an early stage in the physics-chemistry borderline area ofkinetic molecular theory; and 4. introduction of the subject in a fashion whichallows a maximum of reasoning by the student and a minimum of memorizationin order to understand chemical and physical behavior of the elements and theircompounds. The descriptive chemistry of the elements discussed later in thecourse, the reactions of qualitative analysis, the principles of acids, bases,amphoterism, solution, conductivity, complex ion formation and many other topicsare then developed by considering the electron distribution within atoms. It isobvious that size and charge relationships also play an important part in thediscussion of these topics. These relationships are continuously stressed.The student body to which this presentation has been offered should be defined;1. Students both with and without high school chemistry background. Approxi-mately some 15 to 20 percent have had no high school chemistry. 2. Studentsin a state supported institution, with an enrollment in the order of 26,000. Theentrance requirements are such that any student in the upper 50 percent of hishigh school class may be admitted. 3. Students with major interests in virtuallyany field excepting chemistry, chemical engineering, and other fields of engineering.The majority are in the premedical, predental, prenursing, and agriculture cur-THE OHIO JOURNAL OF SCIENCE 59(1): 15, January, 1959.Hi CHEMISTRY SYMPOSIUM Vol. 59ricula. Actually, a similar approach in presenting chemistry is used with thechemistry majors. 4. Class size of approximately 250 with four or five suchsections being taught.The essential points, then, in the development of atomic architecture that arepresented very early in the course are briefly summarized: 1. experimental evidencesfor fundamental particles; 2. experimental evidences for relationships betweenthe nucleus and the extra nuclear portion of the atom; 3. nuclear particles andatomic number; 4. experimental evidence for the electron arrangement throughspectroscopic observation; and 5. the arrangement of electrons and their designa-tions by four quantum numbers in terms of energy levels is finally developed.These lectures on atomic structure are followed by a systematic presentationof the periodic table and the classification of the elements on the basis of atomicnumber. This is accompanied by a thorough consideration of types of chemicalbonds, both as to similarities and fundamental differences.The types of materials covered in the unit on atomic architecture will bementioned principally in topical form by referring to a number of figures. Thismaterial is presented by a combination of lectures, lecture demonstrations, theuse of lantern slides, and with "chalk-talks." The manner of the intensive treat-ment of these topics will vary somewhat depending upon the composition of thevarious groups and the needs of the lecturer.The main areas of the problems involved in studying the atoms of elementssystematically are presented in the following steps.I. The importance, nature, and behavior of the electron is developed throughdiscussions based upon—1. The Millikan oil drop experiment (fig. 1).2. Information derived from the vacuum tube in magnetic fields. Theconcept of canal rays comes also from vacuum tube experiments (fig. 2).S. The x-ray tube and the concept of bombardment of a target by electronsresulting in x-ray emmission. This observation fits nicely into later discussionsrelating energy of light to electron excitation and return of electrons to normalstates (fig. 3).II. The general observations of particles and energy from radioactivematerials (fig. 4).III. Particles derived from atomic nuclei. These particles are more or lessdogmatically described as coming from the nucleus at this point, with evidenceto be presented later (fig. 5).IV. Moseley's Law and the relation between x-ray wave length and massof the nucleus, in turn related to proton content of the nucleus (fig. 6). Thepoint is stressed that the chemical behavior of an element is a function of theproton number since the number of electrons and their final arrangement will berelated to this value.V. The neutron and its derivation by a nuclear reaction:The treatment of this topic includes a brief discussion of the history of the discoveryand the unusual properties of this nuclear particle (fig. 7).VI. The positron and its derivation by a nuclear reaction:[—>1eC+/3+ (positron).There is a discussion of "antiparticles" at this point. Thus, the positron is theantiparticle of the negative electron (fig. 8).No. 1 TEACHING GENERAL CHEMISTRY 17VII. The Rutherford gold leaf experiment: The great importance of thisexperiment, surprising in its fundamental simplicity, is discussed as showing forthe first time the approximate sizes of atoms and nuclei. There is a qualitativediscussion of the experiment to show that gold atoms have nuclei about 4 x 10~12cm in size, and that most of an atom must be empty space (fig. 9).VIII. The concept of the spectrum of an element obtained with spark andarc-activated elements and a spectrograph. This discussion is accompanied withelectron charge;408Q5 ESUHEATER FILAMENTFIGURE 1. Millikan oil drop experiment.FIGURE 2. Studies with cathode ray tubes.FIGURE 3. The x-ray tube (Coolidge type).demonstrations of flame tests (fig. 10). Chalk-talk techniques are extensivelyused at this time to bring out the following concepts:1. General concepts of the wave nature of light and the relation of energyto wave length.2. The concept of activated electrons within atoms and the release ofenergy as electrons drop from outer to inner orbits.3. The concept of the principal quantum number (fig. 11).IX. The need for additional energy levels to account for the multiple lines18 CHEMISTRY SYMPOSIUM Vol. 59of the spectrum (fig. 12). At this point there is consideration of the secondquantum number and the designations of the subshells.X. The third quantum number. The element sodium is used to bring outthe meaning of the third quantum number (fig. 13).XL The fourth or spin quantum number. The discussion brings out thefact that all the electrons in any given atom have an "address" no matter howI. ft particles *• a particlesFIGURE 4. Radiations from radioactive subtances.FIGURE 5. Particles derived from nuclei.FIGURE 6. The contribution of Moseley.BRASSNUCLEAR STRUCTURE ANDRADIOACTIVITYprotonneutronelectron(e'or/? )+ proton energyalpha (a) + energy!/) 'loss of 2 in atomic number and 4in atomic massNo. 1 TEACHING GENERAL CHEMISTRY 19large the atomic number. It is pointed out that no two electrons can have thesame address or state. This is a way of stating to general chemistry studentsthat no two electrons in the same atom can have the same four quantum numbers(fig. 13).XII. The development of the first transition series follows at this point afterhaving developed the structures of the elements through calcium (fig. 14). Theapparent inconsistency of chromium is not stressed; though the 4s1, 3d5 arrange-THE NEUTRONOnneutronTHE POSITRONboron alpha•K®7" 0 nI nitrogen neutronRUTHERFORD GOLD LEAFEXPERIMENT(1911)Aui leaf20,000plate or,screen~l.6 xJ.Q. inor4 X 10 o fcmFIGURE 7. The neutron.FIGURE 8. The positron.FIGURE 9. The Rutherford gold foil experiment.CARBON ARCAS HEATSOURCEPENCIL OFLIGHTyRED LIGHT A~\SPECTRALLINESHEATEDATOMSSLITPRISMVIOLETLIGHTPHOTOGRAPHICPLATEELECTRON ARRANGEMENTELECTRONIN NORMALSTATE> LIGHT EMITTED WHEN ELECTRON RETURNSTO NORMAL POSITION ON COOLINGHIGH ENERGY LIGHT EMITTED(VIOLET)ELECTRON ELEVATED OR EXCITED ON HEATINGLARGE AMOUNT OF HEAT ABSORBEDPOSITIVELYCHARGED NUCLEUS/ LIGHT OF ,LOWER ENERGY/EMITTED(REO)EXCITED\ STATE\LESSER AMOUNTOf HEAT \- ABSORBED1\NORMAL, - ) STATEEXCITEDELECTRONPRINCIPAL SHELL DESIGNATIONSUBSHELLDESIGNATIONMAXIMUM NO.OF ELECTRONSIN SUBSHELLSFIGURE 10. Schematic diagram of a spectrograph.FIGURE 11. The principal orbits or shells of the atom showing electrons innormal and excited states.FIGURE 12. The electron subshells of the atom.maximum impairing in a given atom. The magnetic susceptibility measurementfollows with no additional theory. Dia- para- and ferromagnetism of elementsand their compounds are easily explained in terms of unpaired electrons.XIII. The importance of filled and half-filled energy levels is then stressedNo. 1 TEACHING GENERAL CHEMISTRY 21to account for the chemical properties of certain elements. Thus, the chemicalproperties of zinc and copper atoms can be contrasted with those of chromiumand manganese. This is brought out in the following observations:1. The zinc atom has a constant oxidation state of 2 because all its sub-shells are filled (fig. 16).2. The oxidation state of Cu(I) is to be expected. The evidence for theCu(II) state must be accounted for by the loss of a single 3d electron (fig. 16).The blue color of the Cu(II) ions as due to an unpaired electron can be mentioned.However it is reasonable to assume that a Cu(III) state is hardly to be expected.p= lor K1= s2-orL 3orMP dcffi 110 0 -10+1 0 -I 0*1 -2-1 0+H2..i = * ( n - Ocmd 0m s=* l /2and- l /2ORBI TALSTHIRD AND FOURTH QUANTUM NUMBERSTHE FIRST TRANSITION SERIES3 Ms p d fcmd \n\u\nm*\ I I I I I I%T3CoVM|1 1 I I I I3d crfcitaisFIGURE 13. The third and fourth quantum numbers.FIGURE 14. The first transition series.3. The manganese atom serves as a good example for a number of pointsof chemical interest. It's relative inertness as a metal follows from filled andhalf filled subshells as well as the small size, in contrast to calcium with its strongelectron donating potential (fig. 15). As the magnitude of positive oxidationstate increases it is logical to believe that there will be less ionic nature to thebonds formed. The half filled orbitals make reasonable a certain amount ofdifficulty in attaining state higher than (II); however, all states up to and including(VII) are possible. Compounds in states above (II) are likely oxidizing agents,and auto-oxidation and reduction reactions can occur.XIV. The use of ion-electron or half cell equations is initiated at this pointfor simple cases in the discussion of the loss or gain of electrons by atoms in the22 CHEMISTRY SYMPOSIUM Vol. 59formation of the ionic bond. A good example of the application of these ideasis the ionization of sodium with the resultant diminution in size from 1.86 to0.95 A (fig. 17). The student is then ready to understand the relative inertnessof the sodium ion as due to the stability of filled orbital systems or paired electronsin the s and p subshells. This is much more meaningful to the student than anexplanation based upon the octet rule. Thus a considerable body of informationis derived through individual reasoning by the student, which is a distinct aid toThe Chromium AtomThe Manganese AtomThe Zinc AtomThe Copper AtomFIGURE 15. Chromium and manganese as transition elements.FIGURE 16. Comparison of the structures of atoms of zinc and copper.the morale of the students. The subject of gaseous ionization potential is intro-duced at this point as a fairly quantitative and useful method for comparing thetendencies of metals to lose electrons and the reluctance of nonmetals to do so.XV. The importance of the sizes of atoms in explaining chemical and physicalproperties of the elements is now pointed out by calling attention to the decreasingNo. 1 TEACHING GENERAL CHEMISTRY 23size of atoms with increase in atomic number, say in the third series of the periodictable (fig. 18). The student has no difficulty in understanding the effect of theincreasing nuclear charge on the attraction for the electrons which are all distributedwithin three main levels. From this time on the student is urged to resort con-tinually to the use of the Campbell chart of the periodic table which shows therelative sizes of atoms and ions.XVI. The stabilization of higher oxidation states is briefly touched upon atthis point using the concept of available d orbitals. This topic as well as bondingin coordination compounds is not presented in any great detail at this time.However, the way has been paved for a fuller treatment in the third quarter inconnection with qualitative analysis. In the third quarter, the student is alsopresented the concept of hybridization of orbitals. This is nicely introducedin discussing the structure of the carbon atom in methane, wherein a 2s electron ispromoted to the 2p state permitting sp3 bonding. The dependence of the structureor molecular aggregates are also discussed at this time. Examples are given ofsquare coplanar (dsp2 bonding), tetrahedral (sp3 bonding) and octahedral (d2sp3bonding) configurations.IONIZATION OF SODIUMN o * ' — * No + le"2 e*IFIGURE 17. The ionization of sodium and a comparison of the sizes of thesodium atom and ion.XVII. Acid-base relationships are explained right from the beginning interms of the structural concepts of the central atom utilizing both Lewis andBronsted theories (fig. 19). Thus, small metal ions with large positive chargesmay be shown to act as strong electron pair acceptors and therefore acidic oramphoteric in their behavior. The strong attraction of the metal ion for the oxideand resulting repulsion of the proton (as part of the hydroxide ion) gives a partialexplanation of the capacity of such metal hydroxides to be amphoteric (or amphi-protic in this case). The same principles are carried through to the nonmetalcentral group (as S(VI), N(V), P(V) ) when these atoms (in varying oxidationstates) are associated with the OH group. The large positive charge and thevery small size will further enhance the acidic behavior. The converse situation24 CHEMISTRY SYMPOSIUM Vol. o9with the large metal ions with small positive charge and resulting poor electronpair accepting property leads to the strong basic behavior or very weak acidicbehavior.As has been indicated before, essentially all the material outlined in ParagraphsI to XVI inclusive is presented as a total unit early in the course in general chem-istry. This leads to a working model for studying the elements and their com-pounds which is applied throughout the rest of the course. Thus, the presenceof a polar bond in water molecules is easily explained and the student is ready toaccept and apply the concept of angular molecules polar in character (fig. 20).13Al06OA lo43A,I7A16*>IOA >04ACONTRACTION OF ATOMS17Cl0 0 9 9 ANONMETALACID-BASE STRENGTHSIZE aCHARGE RELATIONweak acid strong acid0Strong baseWeak baseandWeak acidFIGURE 18. The effect of nuclear charge upon atomic size within sameseries in periodic table.FIGURE 19. Acid-base strength relations to size and charge on central atom.Students find this much more meaningful than to refer to water merely as HOHor H2O. As a specific example of the many consequences of the polar characterof the water molecule, consider the great differences in the behavior of methane(a typical nonpolar covalent molecule) and sodium chloride (a typical ioniccompound) in water (fig. 21). Methane is a gas and but slightly soluble in water.Sodium chloride is a crystalline substance of high melting point but quite solubleNo. 1 TEACHING GENERAL CHEMISTRY 25in water due to solvation effects dependent upon the polar character of the watermolecule. The resulting solution, of course, is very different from the solutionof methane in water. It is the feeling of our teaching staff that students, forexample, can approach the study of the colligative properties of solutions withgreater understanding using our working model than using the traditional physico-chemical approach employing Racult's law.In conclusion, it is the prime purpose of this type of introduction to atomicarchitecture to give the student a working model, which it must be admitted isderived with many over simplifications, but which appears to have real teachingl» 2>2p Is) )) + *(2e2e4e le2 H — >THE WATER MOLECULE POLAR BONDSOLVATION a SOLUTIONH°H°C°H°H°no so|va,tionno solutionno conductionNoBlHtlifconductionFIGURE 20. The shape and polar character of the water molecule.FIGURE 21. Solution effects dependent upon the polar character of water molecules.potential. An essentially identical model has been given to high school juniors,with sufficient success so that the feeling exists that we do not need to wait till asenior college course to develop a reasonably precise working model from thechemical point of view. Students going on in chemistry are in a good position toaccept the more advanced postulates and theories of quantum mechanics, fieldtheory, and other related topics. These not continuing for a major but takingmore chemistry courses and even terminal students, do not find such a presentationtoo difficult. The few extra periods devoted to structure at this early point paylarge dividends in later portions of the course.

Atomic Architecture for the General College Chemistry Student

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Atomic Architecture for the General College Chemistry Student

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