The traditional Representation of chemical compounds

1. The traditional Representation of chemical compounds Goals: Name the compounds (XVIII century) Describe them: Geometry and electron distribution Global formulae, planar representation, representation in space Single or multiple representations different models, different ways of “ counting electrons “

2. Name and Describe Naming is important: God’s power in genesis And God said, Let there be light: and there was light. Name and Describe In chemistry we want to describe at the same time. Ideograms are more descriptive than names 化学 Huàxué : Chemistry Science of transformation

3. The domain of Chemistrythat of Democritos Δημόκριτος , born 460 bc Most of the understanding of chemistry goes back before the knowledge of electrons, including the discovery of the periodic classification. What are the forces in physics? Which one concerns Theoretical Chemistry?

4. The domain of Theoretical Chemistry of J. J. Thomson electron discovery in 1887 • Domain of electrostatic(electromagnetic) forces. Coulomb Law • Nuclei will be treated as positively charged particles. This charge is equal to the atomic number of the atom, Z. It is a multiple of the charge of the proton, + e (e = 1.6 10-19 C). The electrons are particles with a charge -e. Between them exercised only electrostatic forces.

5. The atomic number, Z Mass number Atomic number Isotopes are used by chemists for characterization (physical chemistry) They should not be distinguished for chemistry itself (reactivity, exchanges of atoms)

6. Chemical formula: AxByXz • This gives composition: stoichiometry • Law of the defined proportions (1807) or the law of constant composition) • Defined does not mean unique: NO2, NO and N2O • Law of multiple proportions. The relative number of different atoms is always a simple integer This is a quantum law. • Few exceptions: CuyS. • As defined originally the ratio of atomic “weights” was approximate due to isotopic mixures. • A consequence about connectivity: in AB2, B has twice neighbors than A (assuming only A-B bonds) "Stoichiometry" is derived from the Greek words στοιχειον (element) and μετρον (measure.)

7. Molecular and structuralformula • This system for writing chemical formulas was invented by the 19th-century Swedish chemist Jöns Jakob Berzelius. • A chemical formula supplies information about the types and spatial arrangement of bonds in the chemical, though it does not necessarily specify the exact isomer. • For polymers, parentheses are placed around the repeating unit. For example, a hydrocarbon molecule that is described as: CH3(CH2)50CH3, is a molecule with 50 repeating units.

8. developedformula • This gives the connection between atoms. • Isomers of position: CH2Cl-CH2-CH3 and CH3-CHCl-CH3 • Skeleton: sequence without monovalent atoms. • Coordination: number of first neighboring atoms.

9. valence or valency number William Higgins (1763-1825), an Irish chemist • a measure of the number of chemical bonds formed by the atoms of a given element. • This definition is ambiguous but historically very helpful. • The International Union of Pure and Applied Chemistry (IUPAC) has made several attempts to arrive at an unambiguous definition of valence. The current version, adopted in 1994,[4]: • The maximum number of univalent atoms (originally hydrogen or chlorine atoms) that may combine with an atom of the element under consideration, or with a fragment, or for which an atom of this element can be substituted. • This definition reimposes a unique valence for each element at the expense of neglecting, in many cases, a large part of its chemistry.

10. valence or valency number

11. Which product is more likely ? HOCl or HClO ?

12. Unsaturation: Multiple bonds to respect valency numbers • H2C=O a double bond is necessary • N2 a triple bond is necessary Never write a quadruple bond! (except for Cr2) • The chemical formula of a saturated hydrocarbon (alcane) is CnH2n+2. For others, the number of unsaturations may be found from y in the formula CnH2n+2-2y. • This is easily generalized: • Replace every monovalent atom by H • Ignore divalent atoms • Replace NH by CH2

13. -H 2 H H H C -H 2 Ring Closure H H H H H H H H H H C C C C Double bond H H H H H H H H H C C C CH CH HC = CH 2 2 Unsaturations: multiple bonds or rings

14. Planar representation of the molecule A developed formula is inadequate to describe the electronic environment of the atoms in the molecule. Even before we talk about electron, from 1845, Auguste Laurent wrote "a system of such formulas is too absolute and if adopted, it would prevent finding a wealth of valuable reports. We do not know how atoms are really, but we already know that in such a compound the atoms are arranged like in another ... I can say that compounds have the same or different structures without knowing anything of their structure."

15. Planar representation: Lewis structures 1916 The representation of Lewis gives the connection between atoms and precise of every valence electron. • Valence electron are described by • a dot (single electron) • a segment (electron pair or bonds) • a small square may indicate a vacant place available for extra electrons Gilbert Newton Lewis, Berkeley (Octobre 23,1875 - Mars 23,1946)

16. History:Lewis (1916), Lamgmuir (1919) The first representation of the Lewis structure used “cubic atoms” well adapted for the octet rule. The binding did result from a pairing of cubes through vertices, edges or faces according to Abegg law.

17. Valence electrons, core electrons • Valence electrons are the least stable ones, core electrons the most stable ones. • Ionization potential for Valence electron is in the range of 10 eV or less, that for core electron in the range of 100 eV or more Why to be interested in the least stable electrons? …because they are involved in chemistry; the others don’t; they are too stable.

18. Number of valence electrons • For main atoms: the last digit of the number of the column in the periodic table. • For Transition metals, the number of the column. • The number of core electronsis the number of electrons of the last rare gas atom before. +2 +8 +8 +18 +18+32 0 → 2 → 10 → 18 → 36 → 54→ 86 The main values correspond to seps by 2, 8 and 18.

20. What makes a formula a stable one? A compound is stable when two rules are satisfied: • the octet rule • and the electric neutrality. If it not possible to satisfy both, obey first the octet rule unless excessive deviation of charge neutrality. Molecules better have only electron pairs (closed shell molecules). No dots.

21. Choosing a central atom • There are no general recipe • For compounds containing multiple elements with only one atom in the formula, the central atom is the least electronegative single atom that is not hydrogen. For instance, in thionyl chloride (SOCl2), the sulfur atom is the central atom. • Do not connect similar atoms between them except C or Si (see Pauling remarks: A-B stronger than (A-A+B-B)/2)

22. The octet rule2 (doublet rule), 8 or 18 (eighteen electron rule) • We are interested in the local environment of a given atom. How many electrons lie around on atom? • Warning! Be careful that the sum of the electrons for all the atoms is larger than the total number of valence electrons. Some atoms belong to two neighborhood and are counted twice! A molecule is stable when for all of its atoms, the number of electrons in the atomic environment is 2, 8 or 18 (the number of valence electrons of the rare gas atom that follows). Two is for hydrogen. The eighteen electron rules apply for transition metals.

23. Notice to find appropriate Lewis structures • Do not couple dots. • Start counting the total number of valence electrons. • Try to build a structure respecting the octet rule • If you fail, start from there and move electrons to adjust.

24. Lewis structures: exceptions • Deficient atoms. BeH2, AlCl3… Lewis acids. We may indicate the lack of electrons by a square. • Hypervalent compounds. Rare gas atoms, large atoms, compromise for avoiding excessive charge (For H2SO4, charge -2 would appear).

25. Building Lewis structure • Start counting the total number of valence electrons. • Try to build a structure respecting the octet rule • If you fail, start from there and move electron to help.

26. Mesomery, Resonance • One structure may be not enough! • Symmetry imposes equivalence between atoms; this is not possible using a single Lewis structure. The neutral formula does not obey the octet rule. One Lewis formula respecting the octet rule does not verify symmetry; several are then needed. The arrow indicates mesomery. Both representation contain vluable (different) information.

27. p Electron • A requirement for conjugation is planarity. • Conjugation takes place for double bonds (or double bonds and electron pairs) that are on adjacent atoms (neither on the same atom nor on atoms separated by saturated atoms).

28. Conjugation is informative even when a unique Lewis formula correctly describes the ground state of a molecule. It indicates potential electron localization occurring for reactivity. The various Lewis structures for a given structure globally describe a molecule in the Valence Bond approach, VB theory. This theory will attribute a weight to each structure.

29. + + + C l C l C l C l C l Conjugation is informative even when a unique Lewis formula correctly describes the ground state of a molecule. It indicates potential electron localization occurring for reactivity. + + +

30. benzene naphtalene 6 p electrons 6 p electrons within each rin 10 p electrons total (two are shared) Symmetry is the important concept Benzene is D6h: an hexagon with an edge of 1.40 Å.

31. s and p separation It is strictly defined using symmetry (QM) • Molecule Atoms orbitals of H2 porbitals of ethene

32. p orbitals of benzene

33. Butadiene p orbitals

34. Allene: Why there is no conjugation between double bonds on the same atom. Hyperconjugation

35. Mesomery: summary • In some cases, several structures are necessary to represent together a molecule. • If these structures are equivalent (symmetry related) the account for the molecule together. • If these structures are different, one is more appropriate; however the others are informative

36. Formal Charges At variance with the octet rule, the counting of formal charge is a partition of electron. The sum of the electrons should be the total number of the valence electrons of the system. The sum of the charge should be the total charge of the system (0 for molecules, A value for ions). These are two informative visions of the electron distribution.

37. Formal Charges, a democratic splitting When distributing electrons on the atoms, one has to split the contribution of electron pairs in A-B bonds between A and B. In the “formal charge” approach, the deal is one electron each even if A and B differ. There are other models also useful that we will see later on. The electronic density on each atom is d=Se The formal charge on each atom is q-d, q being the atomic number of valence electron of the atom. This definition matches the “Mulliken charge” definition in theoretical chemistry, which is a standard.

38. Formal Charges, a democratic splitting A formal charge is a partial charge on an atom in a molecule assigned by assuming that electrons in a chemical bond are shared equally between atoms, regardless of relative electronegativityor in another definition the charge remaining on an atom when all ligands are removed homolytically.

40. ¨ :C≡O: or :C=O ¨ The case of carbon monoxide, CO • The total number of valence electron is 4+6=10. • One Lewis structure satisfies the octet rule and not satisfies the electron neutrality whereas the other does the contrary. • The dipole moment is very weak; it corresponds to the charged formula. Since O is more attractive to electrons than C, the formal charges however come out very small.

41. Several possible Lewis structures ¨ ¨ ¨ ¨ :Cl-Mg-Cl: or Cl=Mg=Cl ¨ ¨ ¨ ¨ Neutral formula Charged formula Satisfies the octet rule Indicates p conjugation Shows Lewis acidity does not

42. Electron count • This is very important. However there are many ways that differ; all are informative • Isoelectronic system Isolobal • HClO HOCl • N2 CO • 3 main ways of distributing electrons: • Formal charges, democratic • Oxidation numbers, not democratic: everything for the most powerful atom • Electronegativity, intermediate

43. The isolobal analogy; electron count Roald Hoffmann Cornell, Ithaca NY Nobel 1981

44. Electronegativity • Partial charges. If A is stronger than B (more electronegative) A has more electrons. • Not to consider in cases of presence of formal charges. • There are many different scales: Pauling, Mulliken, Allred and Rochow, Sanderson, Allen… • Decomposition of dipole moments may generate partial charges on the atoms (however decomposition is also an oversimplification)…

45. Pauling Electronegativity

46. Pauling Electronegativity 1932 The covalent bond between two different atoms (A–B) is stronger than would be expected by taking the average of the strengths of the A–A and B–B bonds. The difference in electronegativity between atoms A and B is given by: where the dissociation energies, Ed, of the A–B, A–A and B–B bonds are expressed in eV, the factor (eV)−½ being included to ensure a dimensionless result. Linus Pauling, Stanford Nobel Chemistry 1954 Peace 1952

47. Mulliken Electronegativity the arithmetic mean of the first ionization energy and the electron affinity should be a measure of the tendency of an atom to attract electrons. c= (PI+EA)/2 As this definition is not dependent on an arbitrary relative scale, it has also been termed absolute electronegativity with eV. However, it is more usual to use a linear transformation to transform these absolute values into values which resemble the more familiar Pauling values. For ionization energies and electron affinities in electronvolts c= 0.187(PI+EA)+0.17 Robert Sanderson Mulliken Nobel 1962

48. Oxidation numbers It is also a partition of electronic density. The electrons from the bonds are attributed to the most electronegative atom. The model is extreme and purely ionic. We can count electrons on the atoms in this doing so and oxidation numbers are the corresponding charge, q-d.

49. Oxidation numbers, N; history They have been defined relative to O (F would have been more appropriate) 4 rules: • N is 0 in elements whatever the allotropic form is. • For O, N=-2 (except when bound to itself: O2, O3,HOOH, peroxides) • H is +1 or -1 according to electronegativities • N for others is deduced so that the total sum is nil (molecules) or equal to the total charge (ions)

50. Octave rule • nmax-nmin = 4-(-4) = 8 for C • nmax-nmin = 5-(-3) = 8 for N • nmax-nmin = 6-(-2) = 8 for S • nmax-nmin = 7-(-1) = 8 for Cl The oxidation numbers vary in a range of 8: By losing electrons, they reach the number of electrons of the preceding gas rare atom (number of core electrons), by gaining them they acquire the number of electrons of the following gas rare atom (number of core electrons+ valence electrons); the difference is eight (number of valence electrons).