1. Polar bonds and molecular polarity Degrees between ionic and covalent

2. Sharing two electrons effectively doubles the count  In the molecule F 2 :  Each atom wants 8  Alone each has seven  Together they have eight  Single covalent bond

3. Covalent bonds between unlike elements  Oxygen requires octet – shares two electrons with H atoms (one with each)  Hydrogen requires two – each atom shares one electron with O O H H H H O

4. Lewis dot structures  In going from G4 – G7, a H atom is replaced by a lone pair of electrons  The total number of electrons is equal to the sum of all the valence electrons  The total number of electrons remains the same – 8  Each atom has a complete octet

5. Multiple bonds are a feature  O 2 and N 2 do not achieve octets by sharing two electrons  Must share more electrons  O 2 has double bond (four electrons shared)  N 2 has triple bond (six electrons shared) – one of the strongest in chemistry  N 2 is very stable and unreactive – also the major product from explosives NN OO

6. Properties of covalent compounds  Gases, liquids and solids at room temperature  May be hard or soft (diamond is a covalent solid)  May be soluble in polar or non-polar solvents  Solutions and melts do not conduct electricity  Most covalent compounds are molecular

7. Polar bonds  The ionic bond and the equally shared covalent bond are two extremes  Complete transfer of charge to equal sharing of charge  Many bonds fall in between: atoms of different elements have different attraction for electrons

8. Electronegativity  The degree to which an atom attracts electrons towards itself in a bond with another atom  highly electronegative atom attracts electrons  weakly electronegative atom does not

9. Table of electronegativity Most electronegative Least electronegative

10. Increasing electronegativity difference increases polarity Non- polar Polar C:CC:NC:OC:F 00.490.891.43 Electronegativity difference

11. The gamut of bonding types Ionic Polar covalent Nonpolar covalent Na:FH:FF:F Sodium fluoride Hydrogen fluoride Fluorine

12. Polar bonds and polar molecules  Any bond containing different elements will be polar to some degree  For a molecule to be polar will depend upon how the bonds are arranged  A molecule may contain polar bonds and be itself non-polar  We need to understand the molecular structure

13. Lewis dot structures: doing the dots  Molecular structure reduced to simplest terms showing only the arrangements of the valence electrons as dots in a 2-dimensional figure  Show only valence electrons  Electrons are either in: bondsbonds lone pairs (stable molecules do not contain unpaired electrons – with very few exceptions)lone pairs (stable molecules do not contain unpaired electrons – with very few exceptions)  Octet rule is guiding principle for distribution of electrons in the molecule

14. Lewis dot structures made easy  Start with the skeleton of the molecule  Least electronegative element is the central atom  S = N - A  N = number of electrons required to fill octet for each atom (8 for each element, except 2 for H and 6 for B)  A = number of valence electrons  S = number of electrons in bonds

15. Applying the rules  Calculate N for the molecule  Calculate A (all the dots), including charges where appropriate (add one for each –ve charge and subtract one for each +ve charge)  Determine S from S = N – A  Satisfy all octets and create number of bonds dictated by S (may be multiples)  NF 3  N = 8(N) + 3 x 8(F) = 32  A = 5(N) + 3 x 7(F) = 26  S = 32 – 26 = 6 N FF F NFF F

16. Two tests for dot structures  Are the number of dots in the molecule equal to the number of valence electrons?  Are all the octets satisfied?  If both yes structure is valid  If either no then back to the drawing board

17. Example of sulphur dioxide  N = 24 (3 atoms @ 8)  A = 18 (S = 6, O = 2 x 6 = 12 valence electrons)  S = 6 (3 two-electron bonds)  12 non-bonded electrons (6 pairs) SO O

18. Expansion of the octet  Elements in second row invariably obey the octet rule  The heavy congeners regularly disobey it  Consider:  OF 2 but SF 6  NCl 3 but PCl 5  Octet expansion is a consequence of the availability of vacant 3d orbitals to the third row, where there are no 2d orbitals in the second row

19. Investigate with dot structures  Proceed with same S = N – A strategy  Octet expansion is indicated by the inability to obtain a reasonable solution

20. Consider SF 4  N = 40, A = 28 + 6 = 34  S = 6  6 bonding electrons and 4 bonds! Means excess electrons  Make bonds and complete octets on peripheral atoms  Add the excess to the central atom

21. PCl 5  N = 48, A = 5 x 7 + 5 = 40  S = 8  8 bonding electrons and 5 bonds  Proceed as before  In this case the octet expansion involves a bonded atom rather than a lone pair

22. Resonance: short-comings of the dot model  The dot structure of O 3 (or SO 2 ) can be drawn in two equivalent ways  Neither is correct in of itself  The “true” structure is an average of the two “resonance hybrids”  Lewis model considers bonds as being between two atoms  In many molecules, the bonding can involve 3 or more atoms  This phenomenon is called delocalization  In O 3 the bonding electrons are delocalized over all three O atoms