9What happens if B is a lot more electronegative than A? In this case, the electron pair is dragged over to B's end of the bond.Ions have been formed.
10Ionic BondingTypically formed between metals and nonmetals (High difference in electronegativities)Electrons are transferred from one atom to another (from metal to nonmetal) resulting in the formation of positive and negative ions.The electrostatic attractions between the positive and negative ions hold the compound together.
14Calculation of enthalpy of formation The loss of an electron from an elementIonization energyNa(g)Na+(g) + 1e- ∆H = 496 kj/molThe gain of an electron by a nonmetalElectron affinityCl(g) + 1e-Cl-(g) ∆H = -349 kj/molAttraction of cation and anionLattice energyNa+(g) + Cl-(g)NaCl(s) ∆H = -788 kj/mol
15Ionic Bonding Energetics of Ionic Bond Formation Lattice energy: the energy required to completely separate an ionic solid into its gaseous ions.k is a proportionality constant(depends on solid structure & e- configs of ions),Q1 and Q2 are the charges on the ionsd is the distance between ions.ChargeDistanceEl
17What happens if two atoms of equal electronegativity bond together? If the atoms are equally electronegative, the bonding electrons are evenly shared
18What happens if B is slightly more electronegative than A? B will attract the electron pair rather more than A does.Uneven sharing results in one side of the bond being more negative than the other (polarity)
19Covalent Bonding Typically between two or more nonmetals No, or low, difference between electronegativitiesPositive nucleus is attracted to negative electron cloud of other atom
21Nonpolar Covalent Both atoms have the same electronegativity. Usually two identical atoms, for example, H2 or Cl2 moleculesThis sort of bond could be thought of as being a "pure" covalent bond - where the electrons are shared evenly between the two atoms.
25Covalent Bonding Lewis Structures Covalent bonds can be represented by the Lewis symbols of the elements:In Lewis structures, each pair of electrons in a bond is represented by a single line:
26Covalent Bonding Multiple Bonds It is possible for more than one pair of electrons to be shared between two atoms (multiple bonds):One shared pair of electrons = single bond (e.g. H2);Two shared pairs of electrons = double bond (e.g. O2);Three shared pairs of electrons = triple bond (e.g. N2).Generally, bond distances decrease as we move from single through double to triple bonds.
27Covalent Bonding Multiple Bonds Generally, bond distances decrease as we move from single through double to triple bonds.Generally, bond energies increase as we move from single through double to triple bonds.
28In order to determine# bonds, we need to learnhow to draw molecules!
29Bond Energy and Enthalpy The enthalpy of a reaction depends on the strength of the bonds of the molecules involved in the reaction.-A reaction with tightly bound reactants will require a higher input of energy to make the reaction proceed than one with loosely bound reactants.-Likewise, the amount of energy required to form the bonds of the products affects the overall enthalpy.
30Bond Energy and Enthalpy The enthalpy of a reaction is given by:∆ H = Σbond energies of reactants -Σbond energies of productsNote that bond energies are always positive quantities.
31Bond Energy and Enthalpy Example:Using bond energies, calculate the enthalpy of the following reaction. C3H8(g ) + 5O2(g ) 3CO2(g ) + 4H2O(l )H-C J/mol C-C J/mol H-O J/mol C=O J/mol C-O J/mol O=O J/mol O-O J/mol
32Bond Energy and Enthalpy Σ bond energies of reactants =8(414) + 2(347) + 5(502) = 6516 kJ/molΣ bond energies of products =6(730) + 8(464) = 8092 kJ/molThe enthalpy change is therefore kJ/mol.
33Drawing Lewis Structures Add valence electrons.Write symbols for the atomsShow initial bondings.Try to complete the octets (8e-)If there are not enough electrons try multiple bonds.
34Drawing Lewis Structures NF3Add valence electrons.Draw skeleton structure-put atom with lowest electronegativityin middle (except hydrogen)4. Show initial bondings.5. Try to complete the octets (8e-)6. If there are not enough electrons try multiple bonds..N5 + (3x7) = 26e-.F.F.F.F.F.F.N.
35Polyatomic Ions Same procedure except: •Take charge into account -add electrons for negative charge-subtract electrons for positive charge•Brackets around structure with charge shown in upper right[ structure ]charge
36Resonance StructuresThe Lewis structure of ozone (O3)
37Resonance StructuresHowever... known facts about the structure of ozone:The bond lengths between the central oxygen and the other two oxygens are identical:
38Resonance StructuresWe would expect that if one bond was a double bond that it should be shorter than the other (single) bondSince all the atoms are identical (oxygens) which atom is chosen for the double bond?
39Resonance Structures-These Lewis structures are equivalent except for the placement of the electrons-Equivalent Lewis structures are called resonance structures, or resonance forms-The correct way to describe ozone as a Lewis structure would be:
40Resonance StructuresThe important points to remember about resonance forms are:The molecule is not rapidly oscillating between different discrete forms
41Resonance StructuresThere is only one form of the ozone molecule, and the bond lengths between the oxygens are intermediate between characteristic single and double bond lengths between a pair of oxygensWe draw two Lewis structures (in this case) because a single structure is insufficient to describe the real structure
42Exceptions to Octet Rule Less than an octet:•”Wimpy” atoms bonding with highly electronegative atoms-typical of B, Be, Al
43Exceptions to Octet Rule Greater than an octet:(central atom must have a d sublevel)-more than 4 atoms around central atom (PCl5)-extra pairs of valence electrons (I3-)
44Formal ChargeSometimes when writing a Lewis structure you come across two different ways to write the molecule, both which look fine.In this case, you should use formal charge to decide which structure is correct for the molecule.
45Formal ChargeThe formal charge is the difference in the number of valence electrons in the atom and the number of valence electrons in the Lewis structure.
46Formal Charge Cf = Ev - (Eu + 1/2Ep) The equation for the formal charge of any atom in a Lewis structure isCf = Ev - (Eu + 1/2Ep)
47Formal Charge Cf is the formal charge whereCf is the formal chargeEv is the number of valence electrons in the bare atomEu is the number of electrons in lone pairs on the atom in the Lewis structureEp is the number of electrons in bonded pairs on the atom in the Lewis structure
48Formal ChargeTo decide if a given structure is correct, check the formal charge on some atoms in all possible structures. In general the most likely Lewis structure has:-all formal charges as close to zero as possible-negative formal charges on electronegative atoms like halogens or oxygen.
49Formal ChargeFor example, consider the methanol molecule CH3OH. This can be written two different ways:In both cases the octet rule is satisfied for all of the atoms in the structure. Which is correct?
50Formal Charge For the leftmost structure The carbon has four bonds, each worth 2 electrons, for a total of eight. It has no lone pairs. Thus, Cf = 4 - (0 + 1/2*8) =0The oxygen has two bonds, each worth 2 electrons, for a total of four. It has two lone pairs. Thus, Cf = 6 - (4 + 1/2*4) =0
51Formal Charge For the rightmost structure The carbon has three bonds, each worth 2 electrons, for a total of six. It has one lone pair. Thus, Cf = 4 - (2 + 1/2*6) = -1The oxygen has three bonds, each worth 2 electrons, for a total of six. It has one lone pair. Thus, Cf = 6 - (2 + 1/2*6) = +1
52Formal ChargeThe leftmost structure has the formal charges closer to zero, and thus is probably the correct structure.
53A water molecule is angular or bent. Molecular Geometry……is simply the shape of a molecule.Molecular geometry is found using the Lewis structure, but the Lewis structure itself does NOT necessarily represent the molecule’s shape.A water molecule is angular or bent.
54VSEPRValence-Shell Electron-Pair Repulsion (VSEPR) is a simple method for determining geometryBasis: pairs of valence electrons in bonded atoms repel one another.These mutual repulsions push electron pairs as far from one another as possible.B BBAB A BB
55Electron GeometriesAn electron group is any collection of valence electrons, localized in a region around a central atom, that repels other groups of valence electrons.The mutual repulsions among electron groups lead to an orientation of the groups that are called electron geometry.These geometries are based on the number of electron groups:2linear3trigonal planar4tetrahedral5trigonal bipyramidal6octahedral
56A Balloon Analogy Each electron group may be: -an unshared pair of electrons, or-a bond (single, double, triple bonds are each counted as one electron group).
57VSEPR NotationIn the VSEPR notation used to describe molecular geometries, the central atom in a structure is denoted as A, terminal atoms as X, and the lone pairs of electrons as E.Example: ClF3 is designated AX3E2. It has three groups (atoms in this case) around the Cl atom, and two lone pairs of electrons on the Cl (draw the Lewis structure to see).
62Polar Molecules And Dipole Moments A polar bond is a bond with separate centers of positive and negative charge.A molecule with separate centers of positive and negative charge is a polar molecule.The dipole moment (m) of a molecule is the product of the magnitude of the charge (d) and the distance (d) that separates the centers of positive and negative charge.m = ddA unit of dipole moment is the debye (D).One debye (D) is equal to 3.34 x C m.
63Polar Molecules In An Electric Field An electric field causes polar molecules to “line up” but has no effect on nonpolar molecules.
64Bond Dipoles And Molecular Dipoles A polar covalent bond has a bond dipole; a separation of positive and negative charge centers in an individual bond.Bond dipoles have both a magnitude and a direction (they are vector quantities).A molecule can have polar bonds, but may be a nonpolar molecule – IF the bond dipoles cancel.
65Bond Dipoles And Molecular Dipoles CO2 has polar bonds, but is a linear molecule; the bond dipoles cancel and it has no net dipole moment (m = 0 D)No net dipoleThe water molecule has polar bonds also, but is an angular molecule.The bond dipoles do not cancel, so water is a polar molecule.Net dipole
66Molecular Shapes And Dipole Moments Molecular polarity can be predicted based on the following three-step approach:Use electronegativity values to predict bond dipoles.Use the VSEPR method to predict the molecular shape.From the molecular shape, determine whether bond dipoles cancel to give a nonpolar molecule or combine to produce a resultant dipole moment for the molecule.
67DETERMINING MOLECULAR POLARITY 1. Molecules are non-polar or polar2. Non-polar molecules have an even (symmetrical) distribution of charge (+ or – )If all atoms are the same in a 2-atom molecule (non-polar bonds)(H2 , N2 , Br2 )If there are no lone pairs on the central atom and the attached atoms are all the same(CO2 , BCl3 , CH4)3. Polar molecules have an uneven (assymetrical) distribution of charge. The molecule has a dipole (+ side and a – side, like a bar magnet)If the outer atoms are different from each other(HCl , H2CO , CH3F)ORIf there are lone pairs on the central atom(H2O , NH3 , SO2 , O3)
68Localized Electron Model A molecule is composed of atoms that are bound together by sharing pairs of electrons using the atomic orbitals of the bound atomsMolecular Structure (Lewis Structures/VSEPR)Bonds within the moleculeValence Bond TheoryHybridization of orbitals
69Atomic Orbital Overlap Valence Bond (VB) theory states that a covalent bond is formed when atomic orbitals (AOs) overlap.In the overlap region, electrons with opposing spins produce a high electron charge density.Overlap region between nuclei has high electron densityIn general, the more extensive the overlap between two orbitals, the stronger is the bond between two atoms.
70The measured bond angle in H2S is 92°; good agreement Bonding In H2SThe measured bond angle in H2S is 92°; good agreementHydrogen atoms’ s-orbitals can overlap with the two half-filled p- orbitals on sulfur.
71Points of VB TheoryMost of the electrons in a molecule remain in the same orbital locations that they occupied in the separated atoms.Bonding electrons are localized in the region of AO overlap.For AOs with directional lobes (such as p-orbitals), maximum overlap occurs when the AOs overlap end to end.VB theory is not without its problems…
72Hybridization Of Atomic Orbitals VB theory: carbon should have only two bonds, and they should be about 90° apart.Reality: carbon has four bonds, which (singly bonded) are about 109° apart.We can hybridize the four orbitals; mathematically combine the wave functions for the 2s orbital and the three 2p orbitals on carbon.The four AOs combine to form four new hybrid AOs.The four hybrid AOs are equivalent, and each has a single electron (Hund’s rule).Four equivalent hybrid orbitalscan now form four bonds
73sp3 HybridizationHybridizing an s-orbital with three p-orbitals gives rise to four hybrid orbitals called sp3 orbitals.The number of hybrid orbitals is equal to the number of atomic orbitals combined.The four hybrid orbitals, being equivalent, are about 109° apart.
74The sp3 Hybridization Scheme Four AOs……form four new hybrid AOs.
75Methane and Ammonia Four sp3 hybrid orbitals: tetrahedral Four electron groups: tetrahedralCoincidence? Hardly…In ammonia, one of the hybrid orbitals (top) contains the lone pair that is on the nitrogen atomIn methane, each hybrid orbital is a bonding orbital
76sp2 HybridizationThree sp2 hybrid orbitals are formed from an s-orbital and two p-orbitals.The empty p-orbital remains unhybridized. It may be used in a multiple bond.The sp2 hybrid orbitals are in a plane, 120o apart.This distribution gives a trigonal planar molecular geometry, as predicted by VSEPR.
77The sp2 Hybridization Scheme in Boron A 2p orbital remains unhybridized.Three AOs combine to form……three hybrid AOs
78sp HybridizationTwo sp hybrid orbitals are formed from an s-orbital and a p-orbital.Two empty p-orbitals remains unhybridized; the p-orbitals may be used in a multiple bond.The sp hybrid orbitals are 180o apart.The geometry around the hybridized atom is linear, as predicted by VSEPR.
80Hybrid Orbitals Involving d Subshells This hybridization allows for expanded valence shell compounds.By hybridizing one s, three p, and one d-orbital, we get five sp3d hybrid orbitals.This hybridization scheme gives trigonal bipyramidal electron-group geometry.By hybridizing one s, three p, and two d-orbitals, we get five sp3d2 hybrid orbitals.This hybridization scheme gives octahedral geometry.
82Predicting Hybridization Schemes In the absence of experimental evidence, probable hybridization schemes can be predicted:Write a plausible Lewis structure for the molecule or ion.Use the VSEPR method to predict the electron-group geometry of the central atom.Select the hybridization scheme that corresponds to the VSEPR prediction.Describe the orbital overlap and molecular geometry.
83Hybrid Orbitals and Their Geometric Orientations
84Hybrid Orbitals And Multiple Covalent Bonds Covalent bonds formed by the end-to-end overlap of orbitals, regardless of orbital type, are called sigma (s) bonds.All single bonds are sigma bonds.A bond formed by parallel, or side-by-side, orbital overlap is called a pi (p) bond.A double bond is made up of one sigma bond and one pi bond.A triple bond is made up of one sigma bond and two pi bonds.
85VB Theory for Ethylene, C2H4 σ-bond: overlap of s-orbital of hydrogen and sp2 hybrid orbital.π-bond has two lobes (above and below plane), but is one bond. Side overlap of 2p–2p.σ-bond: sp2 - sp2 overlap
87VB Theory: Acetylene σ-bond: s - sp overlap σ-bond: sp - sp overlap Two π-bonds (above and below, and front and back) from 2p–2p overlap…σ-bond: sp - sp overlap…form a cylinder of π-electron density around the two carbon atoms.