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Polar bonds and molecular polarity

Polar bonds and molecular polarity . Degrees between ionic and covalent. 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. H. H. O. O. H. H.

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Polar bonds and molecular polarity

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

  2. Sharing two electrons effectively doubles the count • In the molecule F2: • Each atom wants 8 • Alone each has seven • Together they have eight • Single covalent bond

  3. H H O O H H 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

  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. O O N N Multiple bonds are a feature • O2 and N2 do not achieve octets by sharing two electrons • Must share more electrons • O2 has double bond (four electrons shared) • N2 has triple bond (six electrons shared) – one of the strongest in chemistry • N2 is very stable and unreactive – also the major product from explosives

  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 Least electronegative Most electronegative

  10. Increasing electronegativity difference increases polarity

  11. The gamut of bonding types

  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: • bonds • 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. 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) NF3 N = 8(N) + 3 x 8(F) = 32 A = 5(N) + 3 x 7(F) = 26 S = 32 – 26 = 6 F N F F N F F F Applying the rules

  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. S O O 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)

  18. Expansion of the octet • Elements in second row invariably obey the octet rule • The heavy congeners regularly disobey it • Consider: • OF2 but SF6 • NCl3 but PCl5 • 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 SF4 • 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. PCl5 • 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 O3 (or SO2) 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 O3 the bonding electrons are delocalized over all three O atoms

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