Bonding Models. Bonding Models (Lewis) Bonding Models (Lewis) Resonance Structures. Section 2 (Chapter 3, M&T) Chemical Bonding

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1 Bonding Models Section (Chapter, M&T) Chemical Bonding We will look at three models of bonding: Lewis model Valence Bond model M theory Bonding Models (Lewis) Bonding Models (Lewis) Lewis model of bonding is simple: electrons are paired (where possible) Radical species have electrons left over after pairing In forming Lewis structures, atom symbols are drawn with their valence electrons (dots) and unpaired electrons are paired with those of other atoms to form covalent bonds Electrons are added as lone pairs and bonds around atoms to give each atom an octet. In certain cases, some atoms can have more than an octet of electrons, some less. The Lewis model is useful for: Showing connectivity between atoms Showing lone pairs A starting point for valence bond models and VSEPR theory Lewis structures are not used to convey anything about the shape of a molecule/ion Lewis Structures Multiple Bonds Resonance Structures In some cases (e.g. C, C ), multiple bonds must exist in order to create an octet In some cases where multiple bonds exist, there may be more than one possible structure. or example, in carbonate ion, C - Structures which differ only in the placement of electrons (position of atoms doesn t change) are called resonance structures. Resonance structures increase the stability of an arrangement of atoms Each resonance structure contributes to the overall bonding picture in a molecule or ion; none individually is correct Bond lengths in C - are all found to be 9 pm. C- bond is ~ 4 pm and C= ~ 6 pm C - C - C - *correct Lewis structure also shows lone pairs

2 ormal Charge ormal charge is a comparison of the number of electrons around an atom in a Lewis structure with what you d find for an isolated atom of that element ormal charges can be used to assess different resonance structures (to see which is most favorable) and to provide a guess at a site s electron richness/poorness ormal Charge = (group # of atom) - (# of electrons around atom in structure) ormal charges are not actual charges, just a method of book keeping ormal Charge In counting electrons, lone pairs count as two electrons and bonds as one (only one of the electrons in each bond is assumed to reside on the atom) Consider the NCS - ion; which is most correct? Most correct structure has: a. lowest formal charges and b. most negative formal charge on most electronegative element N C S N C S N C S *again, correct Lewis structures also show lone pairs Expanded Shells The Lewis structure that seems to agree best with experimental findings is the one that has expanded octets and lowest formal charges Many common organic compounds obey the octet rule; however, many inorganic ones incorporating heavier elements* (e.g. rd period) may possess more than eight electrons (an expanded octet ) Examples include S 6, P 5, P Rem: correct Lewis structures will also show non-bonding electrons P.. Br.. S 0e - 0e - e - *possibly due to the availability of low lying, empty d-orbitals in these elements Less than an octet or group II and group III elements such as Be (in Be ) and B (in B ), it is not possible to obtain an octet without creating unreasonable formal charges B + - B or these cases, experimental evidence seems to indicate several possibilities are important + B - VSEPR Theory Valence Shell Electron-Pair Repulsion theory is used to predict the shape of molecules based on electron-pair repulsions Pairs of electrons include bonding pairs and non-bonding pairs (lone pairs) eirarchy of repulsions Lone pair-lone pair Lone pair-bonding pair Bonding pair-bonding pair

3 VSEPR Theory The repulsions that exist between electron pairs on the central atom dictates the arrangement of the other atoms The number of pairs of electrons around the central atom are sometimes called electron domains/steric number Each of the following count as one electron domain: Single bond Double bond Triple bond Lone pair Effect of Lone Pairs Increasing the numbers of lone pairs around the central atom increases the degree of repulsion experienced by the bonding pairs The net effect is a closing of the -X- bond angle This effect is seen in other geometries also non-bonding pair A bonding pair B Notice that while these molecules have different molecular geometries, they all have the same electron domain geometries (tetrahedral) Molecular Geometry tetrahedral trigonal pyramid bent 5-Coordinate (Trigonal bipyramid) Geometry Two lone pairs on a trigonal bipyramidal EDG The trigonal bipyramidal geometry involves two kinds of environments (axial and equatorial) Introducing a lone pair into this geometry creates two possibilities eirarchy of repulsions (angle): 90 o >>0 o >80 o strong repulsion weak repulsion strongest repulsion least favorable angle 6 total 90 o repulsions in each x lp-lp x lp-bp x bp-bp the difference 4 x lp-bp x bp-bp Effect of Multiple Bonds Electronegativity and Size Effects Multiple bonds confine a greater amount of electron density between two atoms than do single bonds. Their effect on geometries is similar to lone pairs (though not as strong) bserved bond angles in molecules are influenced by both the size and electronegativity of atoms The influence of size and electronegativity depends on the position of the atom in the molecule (central or peripheral) Multiple bonds Larger groups

4 Periodic Electronegativity Trends Periodic Electronegativity Trends Electronegativity trend Increases L R Decreases top bottom These electronegativities are derived from bonddissociation enthalpies (Pauling) igure (previous slide) shows Pauling electronegativities, χ P, but the trend is the same as for Mulliken (χ M ) values Bond dissociation enthalpy, D: (g) (g) D = 4 kj. mol - (g) (g) D = 40 kj. mol - Averaging would describe ½(D + D ) = 6 kj. mol - a situation where the bonding electrons are equally shared Pauling electronegativity is related to the difference in the average of the homonuclear bond dissociation enthalpies and the value for the heteronuclear (experimental) bond enthalpy Experimental bond dissociation energy (g) (g) + (g) D = 48 kj. mol - Difference results from unequal sharing of electrons in - bond Electronegativity and Size Effects Peripheral atoms The greater the size of the outer atom, the larger the bond angle The greater the electronegativity of the outer atom, the smaller the bond angle Most times, these effects are cooperative Bond angle as a function of outer atom electronegativity Effect of size/electronegativity of outer atom is smaller and more electronegative than Electronegativity and Size Effects Bond angle as a function of center atom electronegativity Where the electronegativity or size of the central atom is concerned, Bond angle increases as the electronegativity of the central atom increases Decreases as the size of the central atom increases N P As Sb 0. o 97.8 o 96. o 87.o or group 5 elements electronegativity: N > P > As > Sb 4

5 Polar Molecules Bonds between non-identical atoms (heteronuclear bonds) will often involve unequal sharing of electrons The situation is one of relative electron deficiency (Q+) at one end and electron excess (Q-) at the other (a dipolar bond) an electric dipole The extent of the polarity of a bond is expressed in units of the dipole moment, m µ = Q r Where Q is the apparent charge (C) of each atom separated by a distance, r (in m). The dipole moment is a measure of the charge separation (bigger the disparity in charge at each end, bigger the dipole moment) Polar Molecules Entire molecules may be polar as a consequence of their polar bonds Molecules with polar bonds may be non-polar as a consequence of their symmetry Sometimes the electronegativity of the atoms involved makes prediction of the dipole difficult N N.47D.85D 0.D Valence Bond Theory (A bonding model that uses combinations of valence atomic orbitals to explain shapes) Consider a molecule. If we try to explain the bonds of this molecule as overlap of valence atomic orbitals of (s) and (s, p), we cannot get a picture that describes a 04.5 o -- bond angle Valence Bond Theory ne way of describing a bonding picture in polyatomic molecules is through valence bond theory (a localized bonding model) This model describes the σ-bonding framework of polyatomic molecules through overlap of hybrid orbitals rbital hybridization: mixing of atomic orbitals ybrid orbital: orbital that is created by mixing two or more atomic orbitals ybrid rbitals Directional hybrid orbitals are obtained by mixing characters of different atomic orbitals of similar energies Use sp hybrid orbitals to describe bonding in linear molecules like Be (g) Two linear combinations of valence atomic orbitals sp ybrid rbital Scheme In this scheme, there are two different combinations of the s orbital and the p x orbital of Be, yielding two different sp hybrid orbitals irst combination adds the p x orbital to the s orbital. The overlap region where the phases (signs) are similar is reinforced, producing a bigger function here. Where the signs are opposite, the lobe becomes smaller (cancellation) = ( s + px ) A coefficient which describes the contribution of each orbital to the hybrid 5

6 sp ybrid rbital Scheme rbital Energy Diagram The second combination involves the subtraction of the p x orbital from the s orbital this implies overlap of the p x -orbital drawn pointing in the opposite direction This hybrid is degenerate (same energy) with the one discussed on the previous slide = ( s px ) Atomic orbitals ybrid orbitals Singly occupied hybrid orbitals can overlap with other singly occupied hybrid orbitals sp ybridization sp ybridization or trigonal planar molecular geometries, there must be a hybrid orbital picture that describes three equivalent orbitals available for overlap with the orbitals (atomic or hybrid) of other atoms sp hybrids involves mixing of an s and two p orbitals. There are three combinations considered sp sp sp _ = hybrid = hybrid = hybrid s + s 6 s 6 px x y + p p x y p p Coefficients which describe the contribution of each atomic orbital to the hybrid orbital The combination of n atomic orbitals will yield n hybrid orbitals Bonding in B Each bond is described by the overlap of a s orbital with a sp hybrid orbital of B The bonds in this picture are c, e - bonds sp ybridization or tetrahedral molecules, we need a scheme that involves four equivalent hybrid orbitals (s, p x, p y, p z ) = = = = ( ) s px ( + ) s px ( + ) s px ( + ) s px p y p y p y p y pz pz pz pz 6

7 N ybrid rbital Picture ther ybridization Schemes.. N N-atom uses sp -hybrid orbitals -atoms use s-atomic orbitals Use sp -hybrid orbitals to describe atoms having four electron domains N 5-coordinate 6-coordinate ther orbital hybridization schemes should involve different combinations of atomic obitals Examples: sp d: trigonal bipyramidal sp d : square-base pyramid sp d : octahedral Molecules having multiple bonds Ethene, C 4 ; ask the questions What is the central atom? C What is the geometry (electron domain) around the central atom? What orbital hybridization scheme will enable this geometry? Are there orbitals left over? Which? ow are they involved? C Unhybridized p z orbitals are available for π-bond formation overlap yields π-bond A Quick Discussion of Bond Symmetry ydrogen Cyanide, CN Sigma bonding symmetry implies that the orbital overlap picture which described the bond is symmetric with respect to a 80 o rotation about the internuclear axis Symmetric no sign change Same questions: What is the central atom? What is the geometry (electron domain) around the central atom? What orbital hybridization scheme will enable this geometry? Are there orbitals left over? Which? ow do they become involved? Sign change for 80 o rotation: π-bond 7

8 s -C sp overlap C sp -N sp overlap unhybridized p-orbitals per atom, available for formation of π-bonds Boron Trifluoride, B Central atom e - domain geometry rbital hybridization scheme (do for also) Unperturbed orbitals what do they do? What is the other sp-orbital on nitrogen for? B is trigonal planar Requires sp hybrid orbitals for boron luorine could have sp or sp Empty p z -orbital on boron can accept lone pair from an adjacent -atom to form a π-bond 8

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