* Your assessment is very important for improving the workof artificial intelligence, which forms the content of this project
Download Midterm 2 Review slides from November 15
Survey
Document related concepts
Woodward–Hoffmann rules wikipedia , lookup
Photoelectric effect wikipedia , lookup
Rutherford backscattering spectrometry wikipedia , lookup
X-ray photoelectron spectroscopy wikipedia , lookup
Homoaromaticity wikipedia , lookup
Metastable inner-shell molecular state wikipedia , lookup
Auger electron spectroscopy wikipedia , lookup
Electron scattering wikipedia , lookup
Heat transfer physics wikipedia , lookup
Physical organic chemistry wikipedia , lookup
Aromaticity wikipedia , lookup
Atomic orbital wikipedia , lookup
Atomic theory wikipedia , lookup
Molecular orbital wikipedia , lookup
Transcript
17/11/2010 Reading assignment: 8.5-8.8 As you read ask yourself: How can I use Lewis structures to account for bonding in covalent m 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 101 1 Lewis structures 1. Find the total number of valence electrons (account for any charges) = TOTAL 1 + 4 + 5 = 10 2. decide connection between atoms, draw a line to represent 1 electron pair for each connection, count the electrons SHARED 2+2=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 Choose central atom correctly least electronegative atom (not H) oxygen rarely bonds to itself more than 1 central atom (e.g. N2O4)? make a symmetric arrangement 4. any electrons left? – put them on the central atom 5. if central atom doesn’t have an octet, make multiple bonds from nonbonded electron pairs on terminal atoms Chem 101 2 1 17/11/2010 Formal charges often can make more than one Lewis structure which one is correct? bookkeeping of electrons 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 101 3 Resonance actual structure is neither A or B but a resonance hybrid structure hybrid is intermediate between the two “parent” parent structures Resonance has impact on bond lengths and strengths Chem 101 4 2 17/11/2010 Exceptions to the octet rule 1. less than an octet small atoms that are too crowded with an octet 2. odd number of electrons 3. more than an octet octet impossible, small number of stable molecules elements in third period (and beyond) can expand valence shell can expand valence shell to make a Lewis structure with lower formal charge Chem 101 5 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 = 431 kJ 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) (single, For polyatomic molecules the bond enthalpies are average values Chem 101 6 3 17/11/2010 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 101 7 ΔH is negative (rxn. is exothermic) when Chem 101 weak bonds are broken and strong bonds are formed 8 4 17/11/2010 Bond lengths also depend on nature of atom and type of bond also calculated as averages trends : shorter bonds are stronger N N 1.47 163 kJ/mol N N 1.24 418 kJ/mol N N 1.10 941kJ/mol Chem 101 9 Reading: Chapter 9, sections 9.1-9.2 As you read these sections ask yourself: How does the number mbe of bonds bo ds and a d nonbonded o bo ded pairs pai s 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 101 10 5 17/11/2010 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 101 11 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 101 12 6 17/11/2010 VSEPR model leads to approximate molecular shape, based on electron locations consider electron domains around central atom domains may be single bond double bond triple bond nonbonded pair each h off these th counts t as 1 domain number of domains around central atom determines the shape of electron domains arranged to minimize the repulsion Chem 101 13 Optimal shape determined by number of domains Two domains: linear, bond angle = 180° Three electron domains: Four electron domains: Five electron domains: Six electron domains Chem 101 trigonal planar tetrahedron bond angles = 120° bond angles = 109.5° trigonal bipyramid octahedral two bond angles 90° or 120° all bond angles = 90º 14 7 17/11/2010 VSEPR steps Draw Lewis structure count electron domains: single, multiple, nonbonded Determine electron domain g geometry y NOT molecular g geometry y Molecular shape is determined by position of bonded atoms example: NCl3 electrons = (7 × 3) + 5 = 26 4 domains electron domain shape is tetrahedral Chem 101 15 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 101 16 8 17/11/2010 sp hybrid sp2 hybrid sp3 hybrid h b id Chem 101 17 sp3d hybrid Chem 101 18 9 17/11/2010 sp3d2 hybrid Chem 101 19 Reading: Chapter 9, sections 9.4-9.7 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 d H does a h hybrid b id orbital bit l diff differ ffrom a pure atomic t i orbital? How are hybrid orbitals related to the VSEPR shapes you learned earlier? How do sigma and pi bonds differ from each other? Chem 101 20 10 17/11/2010 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 p polar molecule Chem 101 21 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 Chem 101 is the vector sum of the bond dipoles 22 11 17/11/2010 To determine if a molecule is polar… draw Lewis structure and determine geometry determine if the bonds are polar p determine if the polar bonds add together (based on geometry) to form a net dipole moment Chem 101 23 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 101 24 12 17/11/2010 Valence Bond Theory valence electrons are in the localized atomic orbitals of isolated atoms these are the s,p,d,f ,p, , 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 101 25 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 101 26 13 17/11/2010 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 3. Chem 101 27 Multiple Bonds The sp and sp2 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 p σ bond – overlap along line between nuclei Chem 101 π bond – overlap above and below the line between nuclei 28 14 17/11/2010 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 101 29 Resonance in valence bond theory two or more resonance structures with pi bonds can not be described with localized bonding th pii b the bonding di iin resonance structures t t iis d l 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 101 30 15 17/11/2010 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 3. 4. sketch molecule starting with central atom and its orbitals, show overlap 5. label all bonds using the σ and π notation Chem 101 31 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 101 32 16 17/11/2010 MO diagram (energy level diagram) antibonding MO raises energy bonding g MO lowers energy bonding electrons Chem 101 33 He He He2 predict that He2 is unstable and does not form Chem 101 34 17 17/11/2010 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 Chem 101 Bd.ord. = ½ Bd.ord. = 0 35 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 p conductivity, y, malleability and ductility Chem 101 36 18 17/11/2010 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 hi h t energy MO are mostt antibonding highest tib di as no. of atoms increases energy separation between MOs decreases bands Chem 101 37 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 101 38 19 17/11/2010 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 67 valence electrons because some electrons are in antibonding orbitals valence electron configurations 3B ns2 (n-1)d ( 1)d1 6B ns1 (n-1)d5 1B ns1 (n-1)d10 Chem 101 39 Metals, insulators and semiconductors differ in the size of the gap between the valence and conduction bands metals t l insulators h have partially ti ll fill filled d band b d with ith no gap have a large gap, filled valence band and empty conduction band semiconductors Chem 101 have a small gap, that electrons can be induced across 40 20