Return to the Spring 1999 CONFCHEM.

A TENTATIVE COMPENDIUM OF PROPOSALS FOR CHANGE IN THE INTRODUCTORY CHEMISTRY CURRICULUM

7th Edition. Last updated 28 January 1999

Compiled by Stephen J. Hawkes
Department of Chemistry
Oregon State University
Corvallis, Oregon, 97331-4003.

Proposals do not necessarily represent the opinions of the compiler. The compiler disagrees with some of them, and some are mutually contradictory. They are included to ensure that the compilation includes many viewpoints and does not merely reflect the compiler's prejudice.

Proposals are worded as debatable propositions, rather than tentative considerations, in the hope that serious debate will follow, leading to drastic reform. After debate, amendment, deletion, extension and adoption, texts that adopt contrary approaches should state why they have done so in footnotes or introduction or in the instructor's manual. National and international examinations should test the resulting curriculum and not the present standard (though non-official) curriculum.

When proposals and comments have come to the compiler's attention from conversations or unpublished communications, no reference is supplied and the source is not revealed.

Further proposals are earnestly invited. So are corrections of errors or misunderstandings in this compendium.

Sweeping Generalizations

  1. Introductory chemistry is so out of touch with reality and with the needs and interests of students that some such listing as this should be constructed and endorsed by influential bodies such as ACS.
            Comment: Publishing individual recommendations or corrections of errors is not sufficient to motivate reform. Some of those listed below are half a century old.
            Comment: The excellent text used in our chemistry major's course still contain the errors that diffusion of gas through the wall of a balloon is governed by Graham's Law, and that in solutions of CaCO3, [Ca2+]=[CO32-], ignoring HCO3- and the three-fold error it causes. The latter error also appears in environmental texts, underlining the deception we have practiced on our students.

  2. A course that is the only chemistry course a student will take must contain everything that is needed to understand subsequent courses for which it is prerequisite and to meet the needs of a citizen. If there is time to include other material that students find interesting, that is a bonus.
            Comment: However, obsolete concepts should not be taught solely because teachers of later courses will use them. We must exercise some leadership.

  3. The introductory course should be the same for all students, including the tiny minority of chemistry majors.
            Comment: Otherwise, a separate chemistry major's course will be regarded as the serious course that is required to judge which students are worthy of admission to competitive programs such as medicine.
            Comment: The actual needs of most students can be met in a one-semester course. Chemistry majors can have the additional material that they are presumed to need in a second semester. Where sadism is required, students can be required to take this second semester. However, these needs are unproven, even for chemistry majors, and often need to be unlearned in later courses.
            Comment: The spirallized system used at Willamette University would accommodate this. Brink,C.P., Goodney,D.E.,Hudak,N.J.,Silverstein,T.P. J.Chem.Educ. 72(1995)530-532.

  4. The introductory course is loaded with calculations that most students will never use, many of which give dramatically incorrect answers. They should be omitted.
            Comment: Too often, students learn algorithms for solving the calculations without learning the chemistry or the meaning of the numbers they produce. This is attested by Lin,H.S., J.Chem.Educ.75(1998)1326 quoting DeBerg,K. Sci.Educ. 73(1989)115-134.

  5. Wherever possible, everyday issues, environmental issues, and practical/industrial chemistry should be among the processes used to exemplify chemical principles, so that the connection between chemistry and real life is made clear. These issues should not be relegated to "optional reading" in textbooks.

  6. The content of general chemistry should reflect a balanced view of the entire scope of chemical phenomena, and should not reflect merely some active frontiers of research. Even though certain aspects of chemistry might be transitorily fashionable for investigations in a few esoteric laboratories or for arcane arguments among mutual admirers at conferences, there continue to be myriad instances of valid chemical phenomena in domestic and industrial contexts; their discussion is proper and essential to prepare a student for an understanding of the chemical aspects of his life and his environment or for a career in which chemical properties and processes play a major role.

Laboratory

  1. Introductory Chemistry, as it is presently taught, focuses on the artificial world of laboratory chemistry. Most students need to understand the very different chemistry of the real world. The laboratory should be de-emphasized or eliminated.
            Comment: Chemistry has been and remains a strongly experimental science. General chemistry courses should include practical work (or experimental classes) in the teaching laboratory as an integral component of a course, about as many hours per week actually at a laboratory bench as in the lecture hall. There should be a close relation between practical work and the total content of lectures within a course, although not necessarily an exact correlation within any particular week (for instance). Even though demonstrations of chemical phenomena in a lecture hall or elsewhere can serve valuable purposes in teaching, these are no substitute for actual laboratory practice, preferably on an individual basis.

  2. The two viewpoints in the previous recommendation may be reconciled by providing introductory students with the laboratory skills they may reasonably expect to find useful in the professions they will adopt.
            Comment: These would include the use of volumetric glassware, specific ion electrodes, and balances. But not titrimetry.
            Comment: Chemistry majors would be provided more fundamental investigation of the relation between experiment and theory in later courses. Other students' will be instructed by their major departments in this relationship, as it applies to their own discipline.

Acidity and Basicity

  1. Acids should be defined as substances from which a proton can be removed, and bases as substances that remove them. While it must be taught that this inevitably produces H3O+ or OH- in aqueous solution this should not be a definition. Arrhenius is obsolete.
            Comment: Acids do not "donate" protons. They yield them to superior force. Bases do not "accept" protons. They tear them off acids. If it is held that the acids must be supposed to do something active (although their role is entirely passive) they may perhaps be said to "supply" protons rather than "donate" them. Bases must still be defined as "removing" rather than "accepting" protons.
            Comment: Similarly, double bonds do not "donate" electrons to electrophiles. The electrophiles tear them off.

  2. pH should be presented as a useful scale of acidity of a solution, and not defined as -log[H+]. The scale should be explained as the ability of a solution to supply H+ to a reaction or to an electrode, measured by the potential at a H+-sensitive electrode, calibrated with standard solutions. (see MacInnes,D.A., Science 108(1948)693 and Hawkes,S.J., J.Chem.Educ. 71, 1994, 747-749.)
            Comment: This is OK, as long as you then go on to say that pH can be operationally equated to -log[H+] in the normal pH range.
            Counter-Comment: But it can't. pH of 0.1M HCl is 1.1.
            Comment: It may perhaps satisfy the academic curiosity of a few students to add that 10--pHis an "effective concentration" (often called an "activity") that has been adjusted for the effects of other substances in the solution.
            Comment: the calculation of pH of solutions of acids, bases, salts and buffers using the methods in introductory texts give dramatically incorrect answers and are not used by a significant number of non-chemists or even chemists.

Equilibria

  1. Since most chemical systems in the real world are far from equilibrium, the treatment of equilibrium should be much reduced. Treatment should be semiquantitative, at the level of "which changes will push the reaction to the left and which to the right?" Calculations should be no more than plugging numbers into the equilibrium constant.
            Comment: There is little usefulness even in calculations this simplistic.
            Counter-Comment: This quotation from Fetter,C.W. "Contaminant Hydrogeology",2nd ed., Prentice-Hall 1999, p.147, suggests the contrary; "If reactions are reversible and proceed rapidly enough, the reaction can be described as being in local equilibrium (Walsh,M.P.S. et al, Amer.Inst.Chem.Eng.J. 30(1984)317-328. Are there enough near-equilibrium systems in the real world, and enough usefulness in calculating them, that it is worth a neophyte's time to learn the algorithms?

  2. Calculation of gaseous equilibria should not be taught.
            Comment: Gaseous reactions reach equilibrium only in the artificial environment of the laboratory. Calculating compositions at equilibrium is therefore rarely useful even to chemists.
            Comment: Instead, the relative rates of the forward and reverse reactions should be discussed. This can be extended by discussing the need to allow time in manufacturing systems to obtain an economic yield and showing how biological systems balance reaction rates to maintain the necessary concentration of a product. The special case of infinite time may be considered, showing that the rates of the forward and reverse reactions change as the concentrations change so that they eventually have equal speed and concentrations remain constant.

  3. Discussion of ionic equilibria should ensure that students understand the qualitative chemistry of such equilibria. This is more important at the introductory level than the ability to use quantitative algorithms.
            Comment: Calculations of ionic equilibria are of little value to students with only one year of chemistry. The algorithms that are usually taught give answers that would be dangerously false if they were used for any real situation. They should not be in the introductory curriculum. If they are nevertheless included, then they must be taught using computer software (or perhaps algorithms that use advanced hand calculators) that is programmed for real systems. The usefulness of such calculations should be shown.
            Comment: The issue is finding the right balance between qualitative descriptions and a sense of how you would go about doing a quantitative calculation.
            Counter-Comment: Why would students in Introductory Chemistry, the vast majority of whom will not take any more chemistry, need to know how to do calculations that they will almost certainly never need to do again?

  4. Values of Ka , Ksp , Kb , and Kform should be those in the NIST data collection (Smith, R.M., Martell, A.E., Motekaitis,R.J. "NIST Critical Stability Constants of Metal Complexes Database", U.S. Department of Commerce, National Institute of Standards and Technology, Standard Reference Data Program, Gaithersburg, MD 20899, 1993) or calculated from them, unless it is definitely known that a different value is better for some particular case.
            Comment: the tables in many texts are out-of-date by several decades. See Clark,Roy, Bonicamp,J.M. 75(1998)1182-1185.

  5. If the hydrolysis of positive ions is taught, it should be in the framework that most positive ions give significantly acidic solutions, and not just a few selected ions. A few positive ions give negligibly acidic solutions. See Hawkes, S.J. J.Chem.Educ. 73, 1996, 516-517.
            Comment: The reverse statement that most negative ions give significantly basic solutions is already included in most texts, so does not belong in this list of proposals for change.

  6. Solubility products, if they are mentioned at all, should be discussed only qualitatively. Simple calculations based on them give dramatically false answers that are seldom in the right order of magnitude, while correct calculations are beyond the scope of an introductory course (unless supported by computer programs written for such calculations). (Hawkes, S.J., CHEM13 News, 250,1996,p.1, and J.Chem.Educ. 75(1998) 1179-1181)
            Comment: I disagree. BUT the students must be made aware that these simple calculations oversimplify the complexity of the chemistry, and cannot be expected to agree with experiment, even within an order of magnitude.
            Counter-Comment: What is the point of doing oversimplified calculations that give the wrong result? A qualitative understanding is all that is required.

  7. The intermediate species in the formation of a complex ion must not be ignored in calculations involving the formation constant. See Banks,J.E. J.Chem.Educ. 38(1961)391.

  8. If equilibrium calculations are taught that involve H+, despite the above recommendation that equilibrium calculations not be taught, then 10-pH should be used rather than [H+]. So "Bronsted" or "practical" Ka (10-pH[base]/[acid]) should be used rather than the thermodynamic Ka or a definition of Ka involving only concentrations. See Hawkes,S.J., J.Chem.Educ. 71, 1994, 747-749 and Hawkes,S.J., Chemical Educator, 1997, 1(6): S1430-4171. URL:http//journals.springer-ny.com/chedr.
            Comment: The pH of 0.1 M KH2PO4 half-neutralised with 0.1MNaOH is 6.82, but the usual textbook procedure plugging concentrations into the thermodynamic Ka gives 7.20, an error of 0.38 unit. Using the Bronsted Ka gives 6.81, an error of only 0.01.

  9. It should not be stated or implied that salts are totally ionized in solution. A very few salts are almost totally ionized. See Hawkes,S.J. 73, 1996, 421-423
            Comment: I cannot see the point of this argument. Certainly, if you want to do quantitative calculations, you need to take into account ion pairing etc. If the focus is on qualitative understanding, why is this level of complexity needed? Also, at the quantitative level it could greatly confuse the students (e.g., writing net ionic equations to explain the reaction between NaOH and HNO3).
            Counter-Comment: Net ionic equations are useful, provided they are recognised as a convenient fiction. The interaction Cr6+ + 3Fe2+ Cr3+ + 3Fe3+ is a useful generalisation even though the Cr6+is always combined with other ions as in CrO42-. But we do not represent it to students as a sensible reality. It is a fallacy that when a salt is only partially ionized the non-ionised part consists of ion pairs: it may, but it is at least as usual that it is truly covalent.

  10. Discussion of vapor pressure over solutions should be limited to Henry's Law. Raoult's Law is so rarely applicable that it is deceptive to teach it to beginning students. See: Hawkes,S.J. J.Chem.Educ. 72 (1995) 204-205, 73(1996)41, Peckham,G.D. J.Chem.Educ. 75(1998)787-788.
            Comment: I agree completely, but think that Henry's Law should be taught, especially in the context of the solubilities of gases in water.

  11. If Raoult's Law is taught, despite the above recommendation, it must not be implied that it results from obstruction of the surface of a solution by solutes. The entropy explanation must be used. See Peckham, G.D., J.Chem.Educ. 75(1998)787-788.
            Comment: While the entropy treatment is correct, it is really a proof rather than an explanation. Few teachers will attempt a sound explanation of entropy such as that in Lowe,J.P. J.Chem.Educ.65(1988)403-406, and the usual explanation as positional disorder will not account for Raoult's Law.

  12. Phase equilibria are important, but should not usually be illustrated using phase diagrams. These disconnect the minds of many students from the physical reality, so that the diagram appears more important than the equilibrium. This distorts their comprehension of simple concepts.

  13. If phase diagrams are shown in introductory texts, despite the above, they should show no inflection at the triple point. The very slight deviation from a smooth curve is too small to show on any scale that can be used. See: Hawkes,S.J., J.Chem.Educ, accepted.
            Comment: Physical chemistry texts could show the area around the triple point on an enlarged scale that shows the discontinuity, if it is considered important.
            Comment: The discontinuity should not be exaggerated in introductory texts to show its existence. This would require a note in the caption explaining that it is exaggerated, which would give the matter undue importance.

  14. Phase diagrams, if they are used, should not show lines that attempt to demarcate a "supercritical fluid" region. Such boundaries have no physical meaning. There is no temperature at which a liquid under high pressure undergoes a change becoming a supercritical fluid.
            Comment: There is a semantic meaning. Practical applications of supercritical fluids are at temperatures and pressures both of which exceed the critical values. There is no theoretical reason that both values must be supercritical but practice seems to require it.

  15. Liquid-liquid partition equilibria should be discussed, showing how the equilibrium is affected by inter-molecular forces.

  16. Adsorption equilibria should be discussed qualitatively, showing how the equilibrium is affected by inter-molecular forces in both phases.

  17. Ion-exchange equilibria should be discussed qualitatively, showing how the equilibrium is affected by the chemistry of the two phases.

Periodic Table

  1. The periodic chart of chemical elements has an experimental basis; it is no result of a theoretical model, simple or complicated. Periodic trends in chemical and physical properties and transmutation of elements due to spontaneous or induced radioactive decay should be introduced in lectures through numerous instances. Laboratory exercises on reactions and properties of specific elements and their typical compounds, in which preparations and analyses can be undertaken simply on a small scale, are an essential adjunct to the lecture content. As particular instances of chemical reactions, industrial processes operating at sites in the vicinity of or region near a particular educational institution are recommended.

  2. If a physical basis is given for the periodic table, it should be related to ionization energies rather than quantum numbers. See: Gillespie,R.J., Spencer,J.N. Moog,R.S. J.Chem.Educ. 73 (1996) 622-627.

  3. Relativistic effects should be invoked to explain why some elements behave otherwise than would be predicted by their position in the periodic table. However, no explanation suitable for introductory students has yet been published. See last paragraphs of Scerri,E.R., J.Chem.Educ 75(1998)1384.

  4. If semimetals are shown in a periodic table they should not include B, Si, Po, or At but should include Se and Bi.

  5. It should not be taught that elements adjoining the staircase-shaped line across the periodic table separating metals from non-metals are semimetals. Some are, some are not. Neither should it be taught that all semimetals adjoin the line. Se and Bi are semimetals also.

Nuclear

  1. The political arguments now surrounding nuclear power and nuclear technology require that graduates have an understanding of the subject; in particular they should understand what levels of radiation can be considered negligible and why.
            Comment: Yes, but some universities and colleges may teach this topic in freshman physics, in which case the freshman chem course should be careful not to duplicate the material. Also, the topic of environmental radiation levels is quite a long stretch from the remainder of the chemistry curriculum, and it may be difficult to fit it in logically. Do you mean to include the definitions of the gray and the sievert?
            Comment: Nuclear power and nuclear technology are important but they are not part of chemistry - they should be taught in physics or engineering.

  2. Calculations of mass-energy relations and, by extension, mass defects and binding energy, are physics rather than chemistry and have no chemical consequence. They should not be part of an introductory chemistry course. This is also true of the zone of stability, magic numbers, intranuclear forces, nuclear engineering, and the effect of velocity on mass.

  3. If the inverse square law is taught in connection with nuclear radiation, it must be emphasized that it applies only to a point source in a vacuum or in a medium that is transparent to the radiation, which air usually is not. Probably it should not be taught in introductory chemistry.

  4. If nuclear weapons are discussed, it must be emphasized that their difference from other nuclear devices is that their energy is released over a few milliseconds and thus with great intensity. Many students do not understand what an explosion is, so do not comprehend how a nuclear explosion differs from a nuclear accident.

  5. Nuclear chemistry should include no more than a superficial treatment of nuclear structure. Its description as a collection of neutrons and protons is sufficient.

  6. Nuclear chemistry does not include the engineering necessary to convert nuclear energy to electricity, but may include the release of energy as less thermodynamically stable nuclei are transformed to more stable ones.

  7. Nuclear chemistry should include the ability of ionizing radiation to ionize or dissociate molecules.

Nomenclature

  1. Too much student time is spent memorising words such as "anion", "Henderson-Hasselbach", "Nernst", "Bessemer", "allotrope", "deliquescent", "isoelectronic", and "electrophilic". Most of these terms should be eliminated from the introductory course. In the cases where descriptive terms are needed, names should be found that convey their meaning directly, such as "negative ion" or "electron-attracting"
            Comment: Greek is no longer the mark of a scholar. Let's write in English.
            Comment: I agree with one reservation. If you expunge the terms cation and anion completely, the students will not be prepared to encounter them in other places. This is the same argument as rigidly using "propanone" instead of "acetone" when every bottle of the stuff is still called acetone. We have to avoid being "prim and proper" in the chemical sense.
            Comment: Agree, except that almost all students will need to know what the terms mean even if they do not use them. It will take many years to expunge the terms from the literature of chemistry and many related subjects.
            Counter-Comment: If we do not teach the words "anion" and "cation" so that students are unprepared to meet them, they will form the nucleus of dissatisfied chemists who will urge their abandonment.

  2. Cathodes and anodes should not be labeled or discussed as positive or negative electrodes, although the terminals of the external device to which they are attached may be so labeled. Discussion of cells should not involve positive or negative electrodes.
            Comment: Students are confused by a "positive" electrode releasing electrons into a galvanic cell, while electrons are released into a voltaic cell by a "negative" electrode.
            Comment: Whether an electrode is positive or negative depends on whether it is observed from the point of view of the solution or of the external device.

  3. "Atomic mass" should be used in preference to "atomic weight".

  4. The quantity conventionally measured in mole units should be called "chemical amount". See Gorin,G. J.Chem.Educ. 71(1994)114-116.
            Comment: In the SI definition of the mole unit, the corresponding quantity is called "amount of substance". But that designation does not make a clear distinction between the quantity measured in moles and either mass or volume, which also are measures of "amount" of matter. Chemical amount is the factor which relates mass to molar mass.
            Comment: "Chemical Amount" is accepted by IUPAC as "the alternative name for amount of substance", McNaught,A.D., Wilkinson,A., Compendium of Chemical Terminology, IUPAC Recommendations, 2nd ed. (1997) p.21.

  5. The ous/ic nomenclature for metallic ions should not be taught, notwithstanding the confusion that will inevitably result. See: Nomenclature of Inorganic Chemistry. Recommendations 1990, Blackwell, Oxford, 1990.
            Comment: We have taught the Stock nomenclature for half a century but the old system is still more widely used. If we teach only the Stock nomenclature, so that students enter the real world knowing only that there is an alternative obsolete nomenclature, they will provide the driving force for change. It must have taken some such strategy to change sal-ammoniac to ammonium chloride.

  6. 1H should not be described as "protium". The term has no wide currency.

  7. H2S should not be called "hydrosulfuric acid". This name has no wide currency.

  8. If "heavy metals" are mentioned they should be defined as the transition and post-transition metals. See Hawkes,S.J., J.Chem.Educ. 74(1997)1374

Units and Calculations

  1. Molality is used only in discussions of colligative properties. This concentration term is not needed elsewhere in the course and its introduction only confuses students.

  2. The quantity "n" should have the units (mol e-)(mol rxn)-1 and not merely mol e-. See: DeKock, J.Chem.Educ. 73(1996)955-956.
            Comment: The matter is not quite as DeKock represents it. Atkins 5th ed. 1994, p.331 and Alberty-Silbey, 1992 p.242 both give a correct formulation.

  3. Common units such as ppm and ppmv should be taught, and students should be able to relate them to "chemists' units" of moles per liter.
            Comment: The IUPAC nomenclature for ppm is g.m-3 while chemists compromise for mg/L. Similarly mg/m-3and µg/L for ppb. Ppmv is unitless.

  4. Physical quantities should be specified in the units recommended by IUPAC. (See:Mills,I.;Cvitas,T.;Homann,K.;Kallay,N. Quantities, Units and Symbols in Physical Chemistry; Blackwell 1988, p.101. Specific cases are:
    1. Traditional units of pressure should be abandoned and the IUPAC nomenclature using Pascals should be universally used.
              Comment: If other pressure units are mentioned at all, the mention should be in a footnote or fine print.
    2. Liter should be abandoned in favor of dm3.
              Comment: This is tilting at windmills. The liter is too popular to be replaced.
    3. Density should be expressed as kg.m-3 (which is g.cm-3 x 1000)
    4. "amu" should be used and the "dalton" or Da or d should not be mentioned.
    5. If "Dalton" is used, despite the above, then the symbol Da recommended by IUPAC should be used and the symbol d or D abandoned. See: McNaught,A.D., Wilkinson,A., Compendium of Chemical Technology: IUPAC Recommendations, p.101.
    6. Calories should be abandoned in favor of Joules.
              Comment: If traditional units are mentioned at all, it should be in a footnote.
    7. Ângström units should be abandoned.
              Comment: I have for years told students that the Ângström was so useful to chemists that it would never be replaced, being about the diameter of an atom. Current usage has proved me wrong, since most atomic lengths are now expressed in pm.


  5. "Dimensional Analysis" or "Factor-Label" methodology should be replaced by "quantity calculus". (See: White,M.A., J.Chem.Educ. 75(1998)607-609. and Mills,I.;Cvitas,T.;Homann,K.;Kallay,N. Quantities, Units and Symbols in Physical Chemistry; Blackwell 1988, p.101.
            Comment: but a less intimidating name than Quantity Calculus is needed.

  6. Tables of data should be headed in the IUPAC notation as quantity divided by units.
            Comment: This nomenclature is now required for publications in J.Chem.Educ. (see J.Chem.Educ. 75(1998)646-647) and by IUPAC for all its publications (IUPAC Quantities,Units and Symbols in Physical Chemistry, 2nd ed. Blackwell Scientific Publications, Oxford, 1993). Students should get used to it.

  7. The ASTM significant figure rule 5.4.3.2 leads to the belief that 0.99 x 1.05 = 1.0. When teaching significant figure algorithms, the rule should be followed by the caution "unless the least precise factor in the calculation begins with a 1 and the product or quotient begins with 5,6,7,8, or 9 (or vice versa) then give the quantity beginning with 1 one more significant figure than the other quantity." See: Fields,L.D., Hawkes,S.J. J.Coll.Sci.Tchg. 16(1986) No.1, p.30.

  8. When rounding the last digit of a number whose last digit is exactly 5, not followed by any other figures, students in U.S.A. should be taught to follow the ASTM rule 5.4.1.3 and round upwards if the result is an odd number and downwards if it is even. Japanese students should follow the "shi-sha, go-nyu" rule (which translates to "drop 4, enter 5").
            Comment: The difference is so minute- and the rule so debatable - that following the local authority is the wisest course.

  9. The limiting reactant problem has too few real-world applications to warrant study in introductory chemistry.

  10. Interconversion of Fahrenheit and Celsius temperatures does not belong in a chemistry course.
            Comment: chemistry uses only the Celsius and Kelvin scales.

  11. Equivalent weight and normality should no longer be taught.

  12. Combustion analysis calculations should not be included. See Owens,P.M.,Springer,D.S. CHED Newsletter, Fall 1996, 67-68.
            Comment: This implies that determination of empirical formulae is unnecessary. For the practical purposes of non-chemists this is true, but it is anti-intellectual. If it is taught for intellectual completeness, it should be presented as a reversal of the calculation of % composition from formulae, rather than the usual presentation as an algorithm. But then it is sufficient to deduce empirical formulae from percent composition without resort to combustion analyses.

Thermodynamics

  1. Gibbs energy (free energy) is poorly and often incorrectly introduced. Indeed much of thermodynamics is not correct as presented. Even the units on thermodynamic parameters for chemical reactions are incorrectly given. Coupled reaction concepts should not be presented but when they are, they're incorrect. (see Spencer,J.N. J.Chem.Educ. 69 (1992) 182-186)

  2. The word "spontaneous" is confusing and should not be used. See Ochs,R.S., J.Chem.Educ. 73(1996)952-954 and correspondence by Earl,B.L. and Ochs,R.S. J.Chem.Educ. 75(1998)658-659.

  3. Entropy includes non-randomness in the distribution of energy, as well as the usually taught non-randomness of position of atoms and bonds. See: Lowe,J.P., J.Chem.Educ. 65(1988) 403-406.
            Comment: Heating an inert gas at constant volume increases the entropy without any change in the positional randomness.
            Comment: This is explained in General Chemistry 2nd ed by Jean Umland and Jon Bellama p.643.

  4. Arguments about free energy must not finish with Go. Go and G must be carefully distinguished.

  5. Standard enthalpy, entropies, and free energies of formation place a barrier between the student and what we are trying to have students learn about reaction thermodynamics. There are better and simpler ways to get students to understand the thermodynamic parameters related to the making and breaking of bonds. (see Spencer,J.N., Moog,R.S., Gillespie,R.J., J.Chem.Educ.73(1996)631-6).

  6. Heat of formation should be abandoned in favor of heat of atomisation. See Spencer,J.N., Moog,R.S., Gillespie,R.J., J.Chem.Educ.73(1996)631-6.
            Comment:The concept conveys more simply the principle that reaction consists of the reorganization of bonds. If it is held that students need to relate this to an extensive set of published tables, such as the tables of heat of formation or combustion, then a footnote can be added showing how such data can be converted to heat of atomisation.

  7. Students should be introduced to computer software that will calculate heats of atomisation ab initio.

  8. "Standard" pressure should be defined as 105Pa in accordance with the IUPAC recommendation instead of 1 atm. (See McNaught,A.D., Wilkinson,A. "Compendium of Chemical Terminology, IUPAC Recommendations" 2nd ed., Blackwell 1997 p.391).
            Comment: It is a small correction (1 atm = 101325 Pa). Let the standard of 1 atm go the way of the obsolete definition of the liter as the volume of the standard liter (1.0008 dm3).
            Comment: The effect on equilibrium constants is discussed in Treptow,R.S. J.Chem.Educ. 76(1999)212-215.

Textbook Errors

  1. Graham's Law of Diffusion does not apply to any circumstance that a student is ever likely to meet. It is deceptive in that it is often applied to situations where it is false. Students are better served by not teaching it. See: Mason,E.A. J.Chem.Educ. 44 (1967) 742-744 and Hawkes,S.J. J.Chem.Educ. 74(1997)1069.

  2. If Graham's Law of Effusion is discussed it must be pointed out that it applies only when the hole through which effusion takes place is so small that the effusing molecules do not touch each other during passage through the hole.
            Comment: Otherwise the relation to kinetic energy makes no sense.
            Comment: Does the law of effusion apply to any circumstance that a student is ever likely to meet?

  3. As an illustration of the kinetic theory, Knudsen's Law should be used in preference to Graham's Laws.
            Comment: The separation of uranium isotopes using UF6 does not involve Graham's Laws but Knudsen's Law. This is usually true of separation of gaseous isotopes.
            Comment: Knudsen's Law governs flow through porous media where the mean free path of a gas is long by comparison with the mean diameter of the pores. Then the mol mass velocity is [(4/3)/(2piMRT)1/2][(d/l)p]. See: Villani,S., Isotope Separation, pub. American Nuclear Society, 1976 pp 66-70, and 155-160. The relation to M1/2 is then the same as in Graham's Laws. The kinetic arguments usually adduced in elementary texts to validate Graham's Laws, will validate the square root relation in Knudsen's Law, and lead to a practically significant result.

  4. The rate of diffusion of a gas through the wall of a balloon (or through any other plastic) is not dependent on its molecular mass but on its molecular diameter. The relation has a complex exponential character, far beyond the mathematics of introductory students. Graham's Laws have no relevance to this.

  5. The rule of thumb that reaction rate doubles for every 10K rise in temperature should not be taught. It is seldom even approximate.
            Comment: However, the belief is so widespread that it may be useful to refute the obsolete notion in a footnote.

  6. Students should not be taught that cold glass flows or that it crystallizes in historic or even in geologic time or that it is amorphous, or that it is a liquid. It is a solid with less molecular order than a crystal but more than the liquid phase of the same substance. (Hawkes,S.J. J.Chem.Educ., submitted)

  7. It should not be taught that colloids are stabilized by Brownian motion. Brownian motion accounts for the stability only of colloidal solutions with the smallest colloidal particles. For most sizes of colloidal particles, stabilization is by convection currents. See Mysels,K.J. J.Chem.Educ. 32(1955)319.

  8. It should not be taught that the solubility of gases necessarily decreases with increasing temperature. It may or may not. See Mysels,K.J. J.Chem.Educ. 32(1955)399.

  9. When the lubricating properties of graphite are taught it must be emphasized that pure graphite is not a lubricant. A layer of gas is needed between the planes of the graphite. See Lavrakas,V. 34(1957)240

  10. When the Hall process is discussed it should be pointed out that some of the Cryolite is consumed in the Hall process for production of aluminum, causing significant environmental pollution by fluorides. See Hendricks,B.C. J.Chem.Educ. 32(1955)97.

  11. Boron trichloride should not be characterized as a covalent compound. It is nearly as ionic as BeF2 or LiF. It is a gas rather than a solid at room temperature only because of its weaker crystal forces. See: Gillespie,R.J. J.Chem.Educ.75(1998)923-925.
            Comment: This must be part of a general misunderstanding about the relation between physical state and ionic character, which needs clarification.

  12. It should not be taught that semimetals are semiconductors. Some are, some are not.

Proposals to Include Subjects Presently Neglected or Underemphasised

  1. Students should be able to consult and understand a "Material Safety Data Sheet", including definitions of LD50, ppb etc

  2. Sufficient organic chemistry should be included to make polymer and biochemistry comprehensible. This will emphasize functional groups and de-emphasize hydrocarbons and their nomenclature.

  3. Biochemistry. See Moore,J.W. CHED Newsletter, Winter 1995, 29-31, Spring 1996, 5&63-64, Gillespie,R. CHED Newsletter, Fall 1996, 44-45, Hawkes,S.J. Fall 1996, 46-47.
            Comment: I have mixed views about this, principally because biological molecules are generally of complex structure. What can/should you teach to students who do not yet know about functional groups? Some high school biology courses attempt to include biochemistry, but the students' understanding at that level is often poor.
            Counter-Comment: Students need to know about functional groups. There should be more organic earlier in the course. There is no good reason for neglecting organic - that chemistry students in the US get organic in the 2nd year is not a good reason. Then they will be ready for a small amount of Biochem.
            Comment: ACS's Education Division is developing a freshman chemistry textbook for science majors that places chemistry in a biological context. Chem. & Eng. News April 20,1998.
            Comment: This recommendation does not suggest just what aspects of biochemistry should be included in introductory chemistry. It may be possible to incorporate much of the biochemistry as illustrations of general principles.
            Comment: the usual biochemistry included in introductory texts includes (a) the structures of DNA and RNA (b) the mechanism by which the genetic code controls the structure of a protein (c) the chemical difference between starch and cellulose and the reason that starch is more reactive (d) the polymerisation of glucose (e) the energy needed to convert CO2 and H2O to glucose (f) the structure, hydrolysis and metabolisation of fats (g) the distinctions among saturated and unsaturated and polyunsaturated fats (h) the chemical structure of cell membranes (i) the penetration of cell membranes by ions or molecules (j) the mechanism of enzyme catalysis (k) the release of energy from glucose and its further release in muscle contraction. No research has been reported to show that these are the most useful parts of biochemistry, but they may be.

    No suggestion is yet made about what aspects of the following subjects should be included in an introductory chemistry course.

  4. The glassy state.

  5. Environmental chemistry.

  6. Chemical safety.

  7. Toxicology. (See Abstracts 263-270 of 214th ACS National Meeting at Las Vegas, Sept 1997).
            Comment:A chemistry course is the only course in which "chemical" can be defined and the alleged risks of chemicals can be discussed. The introductory chemistry course should therefore include some toxicology, even though it is not chemistry.

  8. Materials chemistry. (See Ellis,A.B., J.Chem.Educ. 74 (1997) 1033-1040)
            Comment: Chemistry is a science of not only molecules but also materials. The several states of aggregation including liquid-crystalline substances and mixtures, plastic crystals and supercritical fluids should be described and illustrated with instances and applications of particular chemical compounds. Not only simple crystalline compounds such as metallic salts and elements in their varied allotropic forms but also polymeric substances should be described and explained as instances of the diversity of chemical nature. Practical exercises involving appropriate materials serve to illustrate these phenomena.

  9. Polymer chemistry. See Kybett,B.D. J.Chem.Educ. 59(1982)724-725

  10. Analytical Chemistry.
            Comment: Students taking post-BS courses for industrial chemists "are predominantly interested in taking analytical, polymer and bioorganic chemistry courses" Chem & Eng News, April 20,1998, p.73.
            Comment: The analytical techniques described in "Environmental Chemistry; Essentials of Chemistry for Engineering Practice" by The Fu Yen, Prentice-Hall 1999 are titrimetry, gamma-ray Spectrometry, X-ray Diffraction, Mass Spectrometry, Ultraviolet/Visible and Infrared Spectrophotometry, NMR, and Separation Science. They are all dealt with very briefly - 30 pages in all. This lists what analytical techniques are important to engineers in the opinion of one writer whose opinion is respected.

  11. Surface Chemistry

  12. Corrosion

  13. Prevention of Chemical waste.

  14. Cleanup of chemical waste.

Atomic Structure

  1. The Bohr atom should not be taught.

  2. Atomic orbital shapes are not needed. Electron domain theory adequately handles atomic geometry. See Gillespie,R.J., Spencer,J.N., Moog,R.S. J.Chem.Educ. 73(1996)622-627.

  3. The "Common Sense Science" approach to atomic structure which considers subatomic particles to be charged rings should replace the quantum models. See: Bergman,D.L. Galilean Electrodynamics 2(March/April 1991) 30-32, Gulko,A.G.,Bergman,D.L., ibid 4,No.5, pp98-100., http://www.cormedia.com/css/info.htm
            Comment: The arguments proposed for the charged ring model assert that it is closer to reality and gives better prediction of experimental data. However, the quantum model has been thoroughly developed and gives excellent predictions in many useful spheres. The charged ring model has not been published in major scientific journals. Until it is so published or its usefulness in predicting chemical structures and reactions has been proved, it is unsuitable for an introductory chemistry course.

  4. The uncertainty principle need not be subject to lengthy discourse.
            Comment: This needs elaboration. What does deserve to be included, if anything?

  5. The Gankin calculations on chemical bonding, which do not involve quantum mechanics, should be taught in introductory chemistry. See: "How Chemical Bonds Form and Chemical Reactions Proceed", Gankin,V.Y., Gankin,Y.V. pub. Institute of Theoretical Chemistry, Shrewsbury, Mass, 1998.
            Comment: But see remarks on "Common Sense Science". in the previous item.

  6. Discussion of the properties of isolated atoms should be carefully distinguished from those of a solid element, so that students do not believe, for example, that the ionization energy of the carbon atom is also that of graphite. See Rustad,D., J.Chem.Educ. 75(1998)542-543.

  7. As a corollary to the previous recommendation students should be shown the ionization energies of some solids and shown that they are more descriptive of electrodes than are the ionization energies of atoms.
            Comment: these are usually called "work functions" but students should not be required to burden their memories with this curious name. It may be included in a footnote.

  8. Quantum numbers have little purpose in introductory chemistry and should not be introduced. (Ed; see Ruis,S.P. J.Chem.Educ. 65(1988)720, and Gillespie,R.J., Spencer,J.N., Moog J.Chem.Educ. 73(1996)622-627).
    For a counter-argument see Richman,R.M. J.Chem.Educ. 75(1998)536.

  9. No attempt should be made to explain why electron shells exist, or why they have the structure that they do. These can be explained only by quantum theory that is beyond the scope of introductory chemistry. They should be presented as experimental reality illustrated by the periodic variation of physical properties. (See Gillespie,R.J., Moog,R.S., Spencer,J.N. J.Chem.Educ(1998)539-540)

  10. It is especially useful to show the variation of first ionization energies with atomic number, because this shows how readily an electron is removed from the atom, and hence why there is a relation between position in the periodic table and the chemical properties of the element. (See Gillespie,R.J., Moog,R.S., Spencer,J.N. J.Chem.Educ(1998)539-540 and Gillespie,R.J., Spencer,J.N., Moog,S.M. J.Chem.Educ. 73(1996)617-622)

The Chemical Bond

  1. Discussion of the chemical bond should be careful of the distinctions among descriptions of the bond, the forces that cause it, the forces that control electron distribution, and steps in the modeling of the bond-forming process that are used for mathematical computation.

  2. It should not be stated or implied that the attraction of nuclei for the small electron density between two atoms is sufficient to cause bonding. See Backsay,G.B., Reimers,J.R., Nordholm,S. J.Chem.Educ. 74(1997)771-776.

  3. Introductory chemistry should not include models of the chemical bond that exist only as steps in mathematical models and have no objective reality. These include (a) sigma and pi bonding (b) hybridization (c) molecular orbitals in general.
            Comment: MO theory is too abstract and of too little utility to students. Once the few examples that the students are capable of understanding are given, there are no further useful applications.
            Comment: Merely reporting the results of calculations based on theories too complicated to describe, or--even worse--giving an alleged explanation of structure based on some preparative stage of a prospective calculation is intellectually unjustifiable and heuristically unsound.

  4. Explanations of crystalline or molecular structure should not depend on classifying or distinguishing electrons, which is physically impossible.

  5. If the chemical bond is qualitatively described, the description should include the importance of the kinetic energy of the electrons. However, no description of this effect has yet been published that could be presented to beginning students. See:Rioux,F., DeKock,R.L., J.Chem.Educ.75(1998)537-539.

  6. Structures of crystals of elementary substances and simple compounds should be presented as an illustration of the variety of nature deduced from experiments, not as a consequence of artifacts of particular physical or mathematical models.

  7. The electrochemical "activity series" is not sufficiently useful to be included in introductory chemistry, and is indeed deceptive.
            Comment: The activity series puts titanium (Eo Ti Ti2+ + 2e- =1.63V) close to magnesium (Eo Mg Mg2++2e- = 2.37V) in nobility, but in practice it is almost as noble as platinum and is described in CRC Handbook as having "excellent corrosion resistance". Conversely, the activity series puts lead (Eo Pb Pb2++ 2e- = 0.13V) as more noble than tin (Eo Sn Sn2++ 2e- = 0.14V) while in practice the reverse is true, tin being "used to coat other metals to prevent corrosion or other chemical action"(CRC Handbook). See Pourbaix,M. "Atlas of Electrochemical Equilibria in Aqueous Solutions" pub. National Association of Corrosion Engineers, Houston, Texas, 2nd English Ed. 1976, p.80. The activity series is deceptive without much more chemistry than is usual in an introductory text.

  8. Models of the chemical bond should be limited to those models that can be used by beginning students to predict something that is useful to graduates with one year of chemistry. This leads to the proposals immediately following.

  9. Multiple bonds should be described by "bond domains". See:Gillespie et al, "Atoms, Molecules and Reactions" p.111.

  10. Students should be shown electron density maps of several different molecules to show the real distribution of the electron cloud in various kinds of bond and the unreality of lone pairs. (See "Teaching Chemistry with Electron Density Models", Shusterman,G.P., Shusterman,A.J., J.Chem. Educ. 74(1997)771-776).
            Comment: Should they perhaps be introduced to computer software that will produce those electron density maps, and use them to predict the reactions of compounds?

  11. The electron domain model described by Gillespie et al is the most suitable that is presently available for presentation to introductory students and should be universally used until a better is offered. It may perhaps be followed by the Lewis model as a shorthand representation. See Gillespie,R.J., Spencer,J.N., Moog,R.S., J.Chem.Educ. 73(1996)622-627.

  12. The Paling scale of electronegativity should be replaced, for the introductory course, with Average Valence Electron Energies. See Gillespie,R.J., Spencer,J.N., Moog,R.S., J.Chem.Educ. 73(1996)622-627.

  13. Resonance should not be discussed in introductory courses.
            Comment: Describing a bond as some intermediate between two Lewis structures is no longer useful.
    Reality is better described by electron density diagrams.

  14. It should not be stated or implied that salts are totally ionized in crystals. (Sanderson,R.T. J.Chem.Educ. 63(1986)845, Ogilvie,J. in "Conceptual Perspectives in Quantum Chemistry" ed. J.-L.Calais, E.,Kryachko, pub. Kluwer Academic Publishers, 1997 pp 127-142.
            Comment: What do you mean by salts not being completely ionized in crystals? To a very good approximation they consist of ions, which are very nearly spherical in the case of monatomic ions. Of course there is a very small amount of shared electron density between the ions but this level requires no more than a mention.

  15. Students should be shown the ionization energies of some solids and shown that they are more descriptive of electrodes than are the ionization energies of atoms.
            Comment: these are usually called "work functions" but students should not be required to burden their memories with this curious name. It may be included in a footnote.

  16. The concept of partial charge should be taught as a predictor of the chemical behavior of a substance.
            Comment: While partial charge is discussed by Gillespie et al (Gillespie,R.J., Spencer,J.N., Moog J.Chem.Educ. 73(1996)622-627) its chemical importance awaits discussion. See also Gillespie,R.J. "Covalent and Ionic Molecules:Why are BeF2 and AlF3 High Melting Solids Whereas BF3 and SiF4 are Gases?" J.Chem.Educ. 75(1998)923-925.

  17. Ligand field theory should not be presented. See Owens,P.M., Springer,D.S. CHED Newsletter, Fall 1996, 67-68.

  18. Paramagnetism should not be presented.
            Comment: The phenomenon is of little interest to non-chemists.
            Counter-Comment: But it accounts for birds' ability to navigate by Earth's magnetic field, and for the existence of the magnetic field of rocks, and its use in dating geological events.
            Counter-counter-Comment: Students will not appreciate these phenomena any better after they have learned to predict paramagnetism of simple compounds using electron pairing in orbital diagrams.

Miscellaneous

  1. Students need not be taught to balance oxidation-reduction reactions in acid and basic media. There is almost no use for rote memorization of some procedure to follow to balance an equation. (see Moore,J.W., J.Chem Educ. 74 (1997) 1253)
            Comment: A balanced equation is seldom useful. Stoichiometric relations depend on changes in charge, or oxidation number, or the number of acid or basic groups in a molecule. Balancing the equation is an unnecessary step.

  2. Deviations from the ideal gas law should be taught qualitatively, not offering a non-ideal equation.

  3. LeChatelier's Principle is complex and need only be presented in a most elementary way.

  4. Atomic spectra require so much explanation that the time needed is not well spent. There are simpler and better ways to introduce energy levels in atoms and quantization.

  5. The Nernst equation cannot be adequately introduced in the classroom. If this particular construct is desired, the laboratory is a better place to introduce it.
            Comment: Why not? And why?

  6. "Hydrophobicity" should be taught as an entropy effect. (See Silverstein, T.P., J.Chem.Educ. 75 (1998) 116-118).
            Comment: This needs to include an explanation why dissolving hexane in water, which has negative H and so is favored, causes a sufficient decrease in entropy to offset the enthalpy. Many chemists believe incorrectly that the solution is prevented by a positive H.
            Comment: it should be emphasised that "hydrophobicity" is a misnomer. Water is not ever repelled. The misnamed effect is caused by water's self-love.

  7. Radiation as an energy source should be included, along with E = hc/, and the role of sunlight as an atmospheric energy source. This will be useful for those instructors who want to develop the theme of alternative energy sources.

  8. Electrochemistry should include discussion of Volta potentials.

  9. Electrochemistry should include discussion of membrane potentials (both biochemical and synthetic).

  10. Colligative properties should be discussed only qualitatively.
            Comment: There is seldom need to calculate the effects, either for chemists or non-chemists. They are no longer commonly used to determine molecular mass, despite the regular use of such determinations in laboratory classes.
            Comment: A century ago, the quantitative variations in colligative properties were important in the development of chemical theory. Introductory chemistry courses do not and should not trace this development, so the theory is not needed for intellectual completeness.

SUSPICIONS by the compiler that do not amount to recommendations.

Any opinions offered on these suspicions will be earnestly considered and researched.

  1. Is the reason that ionic crystals split along a plane, really that movement by the width of one ion causes repulsion right through the plane? Do molecular crystals like iodine, which have no such alternating charge structure, split along a plane? Graphite does. Should we believe that the mechanism that causes molecular crystals to split along a plane does not apply to ionic crystals? Or is the charge repulsion only a contributing factor?

  2. Textbooks often calculate the wavelength of baseballs using the de Broglie equation. Does a wavelength of 10-34m have any meaning in this context? Does any useful learning result from plugging numbers into the de Broglie equation, even if meaningful? How small must a particle be for the deBroglie wavelength to have meaning?

  3. It has been asserted (Ogilvie,J. In "Conceptual Perspectives in Quantum Chemistry" ed. J.-L. Calais, E. Kryachko, pub. Kluwer Academic Publishers, 1997 pp 127-142.) that the oxygen ion in solid MgO has a charge of about -1 rather than the -2 that is traditionally taught. This is consonant with the electron affinity of oxygen which shows that 02- is unstable. This further accords with the observation that when alkali metals burn in air they do not normally form M20 but peroxides and superoxides O22- or O2- in which oxygen atoms have an average charge of -1 or less. Too much can be extrapolated from this, but is there any evidence that 02- ever exists, other than in vacuum tubes? Does this say anything about S2-?

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