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Unit 7 Lewis diagrams molecular geometry bond and molecular polarity IMFAs.

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1 Unit 7 Lewis diagrams molecular geometry bond and molecular polarity IMFAs

2 Lewis dot diagrams add up the total number of valence electrons for all atoms in the molecule arrange the atoms to pair up the separate atoms’ single electrons as much as possible confirm that:  the total number of electrons exactly matches the total valence electrons of the original atoms, and  each atom has an octet of electrons (8), except  H and He have a duet of electrons (2)

3 structural formulas also called “Lewis structures” or “Lewis diagrams” (but not “Lewis dot structures”) replace each shared pair of electrons with a solid line representing a covalent bond consisting of two shared electrons continue to show the lone pairs of electrons (which are unshared) double-check that the lone pairs plus bond pairs still add up to the correct total number of valence electrons

4 multiple bonds additional bonds may need to be added to a Lewis structure if  single electrons remain  atoms do not have octets in simple cases, you may be able to pair up single electrons on adjacent atoms to form additional bonds, e.g.  CO 2 N2N2 C2H4C2H4

5 multiple bonds in other cases, you cannot strictly keep electrons with their original atoms; the electrons are free to move elsewhere in the molecule as needed to complete octets, e.g.  carbon monoxide, CO  ozone, O 3 in these cases, atoms may not form their “normal” number of bonds but the total number of valence electrons must not change; they are just rearranged

6 multiple bonds computational approach you can also calculate exactly how many bonds are in a molecule in the following way  add up the valence electrons that the atoms in the molecule actually have  separately add up the valence electrons those atoms need in order to have noble gas configurations  calculate the difference, need – have that difference is the number of shared electrons the molecule must have every 2 shared electrons make one bond

7 multiple bonds computational approach O2O2 after building the basic skeleton with bonds  add remaining electrons as needed to complete octets  double-check that the total number of electrons is exactly the number of valence electrons (“have”) have: 6 + 6 = 12 need: 8 + 8 = 16 O 4 shared e - thus 2 bonds CO have: 4 + 6 = 10 need: 8 + 8 = 16 C O 6 shared e - thus 3 bonds

8 general hints for Lewis structures if a given molecule can be drawn with both symmetrical and asymmetrical structures, the symmetrical one is more likely to be correct central atoms are often  written first in the formula  the least electronegative element  the element that can form the most bonds hydrogen and halogens  only form one bond, thus are terminal atoms  are generally interchangeable in molecules

9 exceptions to octet “rule” most atoms have octets (8 valence electrons) when in molecules, but there are exceptions groupnumber of electrons number of bondsexamples column 1duet (2)1H 2, LiH column 2quartet (4)2BeH 2, MgI 2 column 3sextet (6)3BH 3, AlCl 3 columns 4- 8 octet (8) 4 bonds 3 bonds + 1 lone pair 2 bonds + 2 lone pairs 1 bond + 3 lone pairs CH 4 NH 3 H 2 O HCl

10 molecular shapes: VSEPR model valence shell electron-pair repulsion groups of electrons naturally find positions as far apart from each other as possible different molecular shapes result based on how many groups of electrons are present each of the following counts as one “set” of electrons around the central atom  a lone pair  a single bond (2 shared e - )  a double or triple bond (4 or 6 shared e - )

11 VSEPR model—central atom with: 2 sets of e – linear e.g. BeF 2 3 sets of e – trigonal planar e.g. BF 3 4 sets of e – tetrahedral e.g. CF 4 5 sets of e – trigonal bipyramidal e.g. SF 5 6 sets of e – octahedral e.g. XeF 6

12 electron geometry vs. molecular shape each set of electrons occupies a position around the central atom the number of sets defines the electron geometry but lone pairs are essentially transparent even though they are invisible, lone pairs make their presence known by distorting the positions of the bonds around them (since lone pairs repel the electrons in the bonds) this results in several related molecular shapes within each general class of electron geometry

13 tetrahedral electron geometry 4 electron sets bondslone pairsmolecular shapeexample 4 single0tetrahedralCH 4 3 single1triangular pyramidNH 3 2 single2bent (~109°)H2OH2O 1 single3 linearHCl

14 tetrahedral electron geometry


16 triangular planar electron geometry 3 electron sets bondslone pairsmolecular shapeexample 3 single0triangular planarBH 3 2 single + 1 double0triangular planarCH 2 O 1 single + 1 double1bent (~120°)O3O3

17 linear electron geometry 2 electron sets bondslone pairsmolecular shapeexamples 2 single0linearBeH 2 2 double0linearCO 2 1 single + 1 triple0linearHCN in addition, any diatomic molecule must be linear (since any two points lie on a line)

18 triangular planar and linear electron geometry

19 bond polarity two electrons shared between two atoms form a covalent bond  if those electrons are shared equally (or nearly equally), it is a non-polar covalent bond  if one atom attracts the electrons much more strongly than the other atom, it is a polar covalent bond  if one atom completely removes an electron from the other atom, the result is an ionic bond

20 bond polarity the electronegativity difference between the two atoms determines how polar a bond is Cℓ2Cℓ2 HC ℓ LiC ℓ bond type Δ EN, electronegativity difference non-polar polar ionic 0.0 – 0.4 0.5 – 1.7 > 1.7

21 dipole moment is the actual measureable quantity related to bond polarity the size of the dipole moment is affected by  electronegativity difference  bond length we will focus on Δ EN and a qualitative sense of bond polarity bond polarity

22 molecular polarity the overall polarity of a molecule depends on the combined effect of the individual polar bonds individual bonds polar overall molecule nonpolar overall molecule polar

23 molecular polarity what allows bond dipoles to cancel?  geometric symmetry of the molecule  having identical terminal atoms (or atoms with the same electronegativity) what prevents bond dipoles from canceling?  geometric asymmetry (due to lone pairs)  having different terminal atoms

24 molecular polarity

25 inherently symmetrical shapes (if all surrounding atoms are the same)  tetrahedral  triangular planar  linear inherently asymmetrical shapes  bent  triangular pyramid even symmetrical shapes become asymmetrical if different terminal atoms are attached

26 IMFA: intermolecular forces of attraction “bricks”— individual atoms, ions, or molecules of a solid “mortar”— holds the separate pieces together (the IMFA)

27 IMFA: intermolecular forces of attraction

28 types of IMFA strongest weakest London forces dipole-dipole attraction hydrogen bond metallic bond ionic bond covalent network occurs between non-polar molecules polar molecules ultra-polar molecules (those with H – F, H – O, or H – N bonds) metal atoms cations and anions (metals with non-metals in a salt) atoms such as C, Si, & Ge (when in an extended grid or network) van der Waals forces

29 consequences of IMFAs melting points and boiling points rise with  strength of IMFA  increasing molar mass substances generally mix best with other substances having the same or similar IMFAs  ”like dissolves like”  non-polar mixes well with non-polar  polar mixes well with polar  (polar also mixes well with ultra-polar and ionic) other physical properties such as strength, conductivity, etc. are related to the type of IMFA

30 predicting melting points, boiling points stronger IMFAs cause higher m.p. and higher b.p.  when atoms/ions/molecules are more strongly attracted to each other, temperature must be raised higher to overcome the greater attraction more polar molecules have higher m.p. and b.p. atoms and molecules that are heavier and/or larger generally have higher m.p. and higher b.p.  larger/heavier atoms (higher molar mass) have more e –  larger e – clouds can be distorted (polarized) more by London or dipole forces, causing greater attraction strategy to predict m.p. and b.p.  first sort atoms/molecules into the six IMFA categories  then sort those in each category from lightest to heaviest

31 same IMFA: sort by molar mass thus at room temperature:  F 2 (g)  C ℓ 2 (g)  Br 2 ( ℓ )  I 2 (s) °C°C –250 –200 –150 –100 –50 0 +50 +100 +150 ex: halogen family all are non-polar (London force) lowest to highest m.p. and b.p. matches lightest to heaviest – 219.62 F 2 (38) meltboil – 101.5 C ℓ 2 (71) – 7.2 Br 2 (160) +113.7 I 2 (257) – 182.95 F 2 (38) – 34.04 +58.8 +184.4 C ℓ 2 (71) Br 2 (160) I 2 (257)

32 same mass: sort by IMFA type °C°C –50 0 +50 +100 +150 ex: organic molecules all are ~60 g/mol different types of IMFA – 0.5 butane (non-polar) +10.8 methyl ethyl ether (slightly polar) +56.2 acetone (more polar) +97.4 1-propanol (ultra-polar = H-bonds) +198 ethylene glycol (can form twice as many H-bonds) the stronger the IMFA, the higher the boiling point

33 isomers (and an isobar) n- and neo pentane glycerol and 1-propanol 1-propanol and methyl ethyl ketone butane and 2-methylpropane 1-propanol and 2-propanol

34 details about each IMFA strongest weakest London forces dipole-dipole attraction hydrogen bond metallic bond ionic bond covalent network

35 London (or dispersion) forces non-polar molecules (or single atoms) normally have no distinct + or – poles how can they attract each other enough to condense or freeze? they form temporary dipoles electron clouds are slightly distorted by neighboring molecules  sort of like water sloshing in a shallow pan

36 London dispersion forces in action non-polar molecules, initially with uniform charge distribution 1. temporary polarization due to any random little disturbance δ+δ+ δ-δ- 2. induced polarization caused by neighboring molecule 3. induced polarization spreads 4. induced polarization reverses

37 dipole-dipole attractions polar molecules have permanent dipoles the molecules’ partial charges ( δ +, δ -) attract the oppositely-charged parts of neighboring molecules this produces stronger attraction than the temporary polarization of London forces  therefore polar molecules are more likely to be liquid at a temperature where similar non-polar molecules are gases

38 dipole-dipole attractions δ+δ+ δ-δ-

39 hydrogen bonding (or ultra-dipole attractions) H—F, H—O, and H—N bonds are more polar than other similar bonds  these atoms are very small, particularly H  F, O, and N are the three most electronegative elements  these bonds therefore are particularly polar molecules containing these bonds have much higher m.p. and b.p than otherwise expected for non-polar or polar molecules of similar mass the geological and biological systems of earth would be completely different if water molecules did not H-bond to each other

40 hydrogen bonding (or ultra-dipole attractions) non-polar molecules (lower boiling points) ultra-polar molecule (much higher boiling point) hydrogen bonds (between molecules, not within them)

41 hydrogen bonding (or ultra-dipole attractions) HH O HH O HH O HH O Beware!! These are not hydrogen bonds. They are normal covalent bonds between hydrogen and oxygen. These are hydrogen bonds. They are between separate molecules (not within a molecule).

42 metallic bonding structure  nuclei arranged in a regular grid or matrix  “sea of electrons”—delocalized valence electrons free to move throughout grid  metallic “bond” is stronger than van der Waals attractions but generally is weaker than covalent bond since there are not specific e – pairs forming bonds resulting properties  shiny surface  conductive (electrically and thermally)  strong, malleable, and ductile alloy = mixture of metals

43 ionic bonding (salts) structure: orderly 3-D array (crystal) of alternating + and – charges made of  cations (metals from left side of periodic table)  anions (non-metals from right side of periodic table) properties  hard but brittle (why?)  non-conductive when solid  conductive when melted or dissolved

44 why are salts hard but brittle? 1. apply some force 2. layer breaks off and shifts 3. + repels + – repels – 4. shifted layer shatters away from rest of crystal

45 covalent networks strong covalent bonds hold together millions of atoms (or more) in a single strong particle properties  very hard, very strong  very high melting temperatures  usually non-conductive (except graphite) examples  carbon (two allotropes: diamond, graphite)  pure silicon or pure germanium  SiO 2 (quartz or sand)  other synthetic combinations averaging 4 e – per atom: SiC (silicon carbide), BN (boron nitride)

46 m.p. = ~1600°C m.p. = 3550°C C 60 buckminsterfullerine “bucky ball”

47 summary of properties strongest weakest London dipole hydrogen metallic ionic network strength soft and brittle strong, malleable, ductile hard but brittle extremely hard van der Waals forces m.p. & b.p. low medium to high very high conductive? no very (delocalized e – ) if melted or dissolved (mobile ions) usually not

48 soaps and emulsifiers some molecules are not strictly polar or non-polar, but have both characteristics within the same molecule non-polar region polar region this kind of molecule can function as a bridge between molecules that otherwise would repel each other oil water soap or emulsifier

49 soaps and emulsifiers with a soap or emulsifier present to surround it, a drop of non-polar oil can mix into polar water

50 IMFAs

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