17/11/2010. Lewis structures

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1 Reading assignment: As you read ask yourself: How can I use Lewis structures to account for bonding in covalent molecules? What are the differences between single, double and triple bonds in terms of bond distance and strength? When is it useful to assess the formal charges on atoms in a Lewis structure? How does a resonance hybrid structure differ from a regular Lewis structure? How do resonance structures affect the predicted bond distances? What are the situations when the octet rule is disobeyed? How will I recognize molecules that disobey the octet rule? How can I make use of average bond enthalpies? Chem Lewis structures 1. Find the total number of valence electrons (account for any charges) = TOTAL = decide connection between atoms, draw a line to represent 1 electron pair for each connection, count the electrons SHARED = 4 3. calculate the remaining electrons = TOTAL SHARED, assign these to the terminal atoms to make octet (or 2 for H atom) 10-4=6 4. any electrons left? put them on the central atom Choose central atom correctly least electronegative atom (not H) oxygen rarely bonds to itself more than 1 central atom (e.g. N 2 O 4 )? make a symmetric arrangement 5. if central atom doesn t have an octet, make multiple bonds from nonbonded electron pairs on terminal atoms Chem

2 Formal charges often can make more than one Lewis structure bookkeeping of electrons which one is correct? calculate the charge on atom IF all bonding electrons shared equally assign to the atom all unshared (nonbonding) electrons + ½ of all bonding (shared) electrons formal charge = number of valence electrons total assigned electrons Evaluate Lewis structures more stable if there are small (or no) formal charges the most electronegative atom has the most negative formal charge Chem Resonance actual structure is neither A or B but a resonance hybrid structure hybrid is intermediate between the two parent structures Resonance has impact on bond lengths and strengths Chem

3 Exceptions to the octet rule 1. less than an octet small atoms that are too crowded with an octet 2. odd number of electrons octet impossible, small number of stable molecules 3. more than an octet elements in third period (and beyond) can expand valence shell can expand valence shell to make a Lewis structure with lower formal charge Chem Covalent bond strength stability of molecule is related to strength of covalent bonds energy change when a particular bond is broken in one mole of gaseous substance is bond enthalpy, ΔH HCl(g) H(g) + Cl(g) H =431kJ bond enthalpies are always positive the greater the ΔH, the stronger the bond depend on atoms in the bond and type of bond (single, double or triple) For polyatomic molecules the bond enthalpies are average values Chem

4 Estimate enthalpy change of a reaction overall change is the difference between bonds broken (it takes energy to break bonds) and bonds formed (forming bonds releases energy) Chem ΔH is negative (rxn. is exothermic) when weak bonds are broken and strong bonds are formed Chem

5 Bond lengths also depend on nature of atom and type of bond also calculated as averages trends : shorter bonds are stronger N N N N kj/mol kj/mol N N kJ/mol Chem Reading: Chapter 9, sections As you read these sections ask yourself: How does the number of bonds and nonbonded dedpairs of electrons affect the shape of a molecule? How is the geometry of a molecule defined? Why is the repulsion between two domains of nonbonded pairs of electrons greater than between two domains of bonded pairs of electrons? Chem

6 Chapter 9 Molecular Shapes Lewis structures provide info on number and types of bonds no info on shape of molecule in 3D shape determined by bond angle: angle between bonds joining nuclei Chem The Valence Shell Electron Pair Repulsion Model VSEPR Assumptions: Shape determined by numbers of valence electrons around the atoms bonded and nonbonded pairs Electron domains arranged to minimize repulsion between domains of electrons VSEPR predicts shape of electron arrangement Chem

7 VSEPR model consider electron domains around central atom leads to approximate molecular shape, based on electron locations domains may be single bond double bond each of these counts triple bond as 1 domain nonbonded pair number of domains around central atom determines the shape of electron domains arranged to minimize the repulsion Chem Optimal shape determined by number of domains Two domains: linear, bond angle = 180 Three electron domains: trigonal planar bond angles = 120 Four electron domains: tetrahedron bond angles = Five electron domains: trigonal bipyramid two bond angles 90 or 120 Six electron domains octahedral all bond angles = 90º Chem

8 VSEPR steps Draw Lewis structure count electron domains: single, multiple, nonbonded Determine electron domain geometry NOT molecular geometry Molecular shape is determined by position of bonded atoms example: NCl 3 electrons = (7 3) + 5 = 26 4 domains electron domain shape is tetrahedral Chem Nonbonding electrons, multiple bonds affect bond angles nonbonded electrons occupy more space than bonded pairs in selected VSEPR geometries: there will be favoured positions for nonbonded electron pairs multiple bonds also have larger electron domains nonbonded electrons and multiple bonds compress bond angles due to greater repulsion Chem

9 sp hybrid sp 2 hybrid sp 3 hbid hybrid Chem sp 3 d hybrid Chem

10 sp 3 d 2 hybrid Chem Reading: Chapter 9, sections As you read these sections ask yourself: How can a molecule with polar bonds be nonpolar? Why do we need theories of bonding that differ from VSEPR? How does Valence Bond theory differ from the Lewis concept of chemical bonding? How does molecular orbital theory differ from valence bond theory? How does a hybrid orbital differ from a pure atomic orbital? How are hybrid orbitals related to the VSEPR shapes you learned earlier? How do sigma and pi bonds differ from each other? Chem

11 Molecular shape and polarity Can now take into consideration the effect of bond polarity on the overall molecule Bond polarity is a measure of unequal sharing of electrons in a bond depends on differences in electronegativity dipole moment defined as a measure of the separation and magnitude of charge of a polar molecule Chem The overall molecular dipole moment depends on polarity of the individual bonds (electronegativity difference and direction) overall geometry of the molecule bond dipoles are vector quantities molecular dipole moment is the vector sum of the bond dipoles Chem

12 To determine if a molecule is polar draw Lewis structure and determine geometry determine if the bonds are polar determine if the polar bonds add together (based on geometry) to form a net dipole moment Chem Bonding theories VSEPR models and Lewis structures good for predicting molecular shapes do not explain why bonds exist do not explain how the electron s atomic orbitals are involved in bonding Need a theory that combines the idea of two electron bonds with the theory of atomic orbitals Chem

13 Valence Bond Theory valence electrons are in the localized atomic orbitals of isolated atoms these are the s,p,d,f, orbitals bond is formed from overlap of half-filled valence orbitals, spin-pairing of valence electrons if the interactions lower energy, a bond is formed shape of molecule determined by geometry of overlapping orbitals Chem Make new orbitals orbitals in a molecule don t have to be the same as in an atom orbitals are (wave) functions can make math combinations to form new orbitals has the effect of mixing the orbitals Hybrid orbitals - have shapes that match actual electron distribution in bonded atoms - number of hybrid orbitals = number of atomic orbitals mixed - central or interior atoms have the greatest tendency to hybridize Chem

14 Bonding schemes: 1. Draw the Lewis structure 2. Determine the electron domain geometry using VSEPR 3. Choose hybrid orbitals for central/interior atoms based on VSEPR shape Chem Multiple Bonds The sp and sp 2 hybrid orbitals have unused p orbitals on the central atom These p orbitals are perpendicular to the hybrid orbitals The p orbitals can overlap sideways Two kinds of overlap π bond σ bond overlap along line between nuclei overlap above and below the line between nuclei Chem

15 single bonds are double bonds are triple bonds are always sigma (σ) bonds one sigma bond and one π bond one σ bond and two π bonds Chem Resonance in valence bond theory two or more resonance structures with pi bonds can not be described with localized bonding the pi bonding in resonance structures t is delocalized li d electrons are delocalized (smeared out) over more than 2 atoms all atoms with delocalized π bonding must be in the same plane Chem

16 Bonding schemes: 1. Draw the Lewis structure 2. Determine the electron domain geometry using VSEPR 3. Choose hybrid orbitals for central/interior atoms based on VSEPR shape 4. sketch molecule starting with central atom and its orbitals, show overlap 5. label all bonds using the σ and π notation Chem Molecular orbital (MO) theory (only section 9.7) Need better theory to understand excited states and the properties of some molecules MO theory describes the electrons in molecules with wave functions called molecular orbitals wave function over entire molecule is constructed from the atomic orbitals of all the atoms in the molecule MO can hold a maximum of 2 electrons MO has a definite energy MO has an electron-density distribution Chem

17 MO diagram (energy level diagram) antibonding MO raises energy bonding MO lowers energy bonding electrons Chem He He He 2 predict that He 2 is unstable and does not form Chem

18 Bond order stability of the bond depends on the relative number of bonding and antibonding electrons Bond order = ½ {no. bonding electrons no. antibonding electrons} antibonding electrons Bd.ord. = ½ Bd.ord. = 0 Chem Metal bonding (Sections 23.5 and 12.2) Electron sea model array of metal cations in a sea of electrons electrons are mobile and uniformly distributed model explains conductivity, malleability and ductility Chem

19 MO model combines atomic orbitals to make MO over entire molecule each MO can hold 2 electrons number of MO = number of atomic orbitals combined in general: lowest energy MO are the most bonding highest h energy MO are most antibonding as no. of atoms increases energy separation between MOs decreases bands Chem bands are not independent and can be represented as one set of energy levels Bonding MOs are called the valence band In the band structure roughly half of the MOs (or energy levels) are BONDING Antibonding MOs are called the conduction band the upper half (high energy half) are ANTIBONDING Conduction arises when electrons are promoted into unoccupied MOs or energy levels Chem

20 Bonding strengthens as electrons are added to bonding orbitals strength of bonding in the transition metals increases until the band structure is roughly half-full roughly 6-7 electrons strength decreases with more than 6-7 valence electrons because some electrons are in antibonding orbitals valence electron configurations 3B ns 2 (n-1)d 1 6B ns 1 (n-1)d 5 1B ns 1 (n-1)d 10 Chem Metals, insulators and semiconductors differ in the size of the gap between the valence and conduction bands metals have partially filled band with no gap insulators have a large gap, filled valence band and empty conduction band semiconductors have a small gap, that electrons can be induced across Chem

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