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Chemical Bonding

Chemical Bonding. Chemical Bonding. is all about the electrons . in most of our discussions we will concentrate on the outer ‘s’ and ‘p’ electrons. These are called the valence electrons. since there are 4 of these orbitals in any quantum it requires 8 electrons (an octet) to fill them.

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Chemical Bonding

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  1. Chemical Bonding

  2. Chemical Bonding • is all about the electrons. • in most of our discussions we will concentrate on the outer ‘s’ and ‘p’ electrons. These are called the valence electrons. • since there are 4 of these orbitals in any quantum it requires 8 electrons (an octet) to fill them. • the tendency of atoms to try to fill out the outer ‘s’ and ‘p’ shells is the octet rule.

  3. Chemical Bonding • Elements will combine in order to fill their valence shell with electrons, like the noble gases. This can happen in one of two ways: • To share electrons - • The outer orbitals of 2 atoms overlap so that each atom is in the vicinity of a full set of valence electrons. • This type of bonding is called covalent bonding.

  4. Chemical bonding • Atoms gain or lose electrons to arrive at a full set of valence electrons. • When atoms gain or lose electrons they become ions. • Ions are attracted to ions of opposite charge and repelled by ions of the same charge. • This type of bonding is called ionic bonding

  5. How Many Bonds ??? • the number of bonds made by atoms in the ‘s’ and ‘p’ blocks of the Periodic Table is determined by how many electrons they are away from an octet:

  6. How Many Bonds ???

  7. most of our discussion will be centred on covalent bonding.

  8. Valence level expansion • Some compounds occur which cannot be easily explained: PF5, SF6, ClF7, ArF8 • in each case the number of chemical bonds is equal to the number of valence electrons. • this can only happen if electrons are promoted to a higher energy level. In this case it is the adjacent ‘d’ orbital.

  9. PF5 • the normal orbital diagram looks like this: • with valence level expansion it looks like this:

  10. Valence level expansion • includes elements of groups 15 to 18, from period 3 down; periods 1 and 2 do not have a ‘d’ orbital to promote to. • Please note that valence level expansion is the exception, not the rule.

  11. Polar-Covalent Bonding • Covalent bonding implies equal sharing of electrons. • If sharing is not equal, the electrons in a bond will spend more time with one atom than the other.

  12. Polar-Covalent Bonding • The atom where the electrons spend more time will have a net negative charge, while the atom at the other end of the bond will be positive. • This type of bond is polar-covalent.

  13. Electronegativity • is a measure of how strongly an atom is holding on to its valence electrons. • If an atom loses an electron fairly easily it has a low electronegativity (and tends to be a cation). • If an atom tends not to lose electrons, but tends to steal them from other atoms (and become an anion) it has a high electronegativity.

  14. To determine what type of bond exists between two atoms you subtract their respective electronegativities: • if the electronegativity difference is 0.2 or less, the bond is covalent • if the electronegativity difference is 1.7 or greater the bond is ionic.

  15. Electronegativity • If the electronegativity difference between two atoms is between 0.3 and 1.6 the bond is polar-covalent. • The greater the electronegativity difference the greater the ionic character of the bond:

  16. Assignment • Determine the electronegativity difference for each chemical bond. If the bond is polar covalent draw an arrow in the direction of the dipole, from positive to negative: C - H END = | 2.5 - 2.1 | = 0.4 polar covalent bond N - H B - F S - O P - H Si - Cl Cu - Br N - I Br - Cl O - H C - Cl C - O

  17. N - H • END = |3.0 – 2.1| = 0.9 • polar covalent bond: N - H B - F • END = |2.0 – 4.0| = 2.0 • ionic bond: [B]3+[F]1- S – O • END = |2.5 – 3.5| = 1.0 • polar covalent bond: S - O

  18. P - H END = |2.1 – 2.1| = 0.0 covalent bond: P – H Si - Cl END = |1.8 – 3.0| = 1.2 polar covalent bond: Si - Cl Cu - Br END = |1.9 – 2.8| = 0.9 polar covalent bond: Cu - Br

  19. N – I • END = |3.0 – 2.5| = 0.5 • polar covalent bond: N - I Br – Cl • END = |2.8 – 3.0| = 0.2 • covalent bond: Br – Cl • O – H • END = |3.5 – 2.1| = 1.4 • polar covalent bond: O – H

  20. C – Cl • END = |2.5 – 3.0| = 0.5 • polar covalent bond: C – Cl C – O • END = |2.5 – 3.5| = 1.0 • polar covalent bond: C – O

  21. Covalently-Bonded Structures • we now have to consider molecules made of several atoms. • most of the following discussion will concern itself with molecules made with covalent or polar-covalent bonds. • ionic bonds (and others) will return later in the unit.

  22. Lewis Structures • Lewis structures allow us to predict how atoms will come together to make molecules. • Lewis structures are representations of molecules showing all electrons, bonding and nonbonding.

  23. Find the sum of valence electrons of all atoms in the polyatomic ion or molecule. If it is an anion, add one electron for each negative charge. If it is a cation, subtract one electron for each positive charge. PCl3 Writing Lewis Structures 5 + 3(7) = 26

  24. Writing Lewis Structures • The central atom is the least electronegative element that isn’t hydrogen. Connect the outer atoms to it by single bonds. Keep track of the electrons: 26  6 = 20

  25. Writing Lewis Structures • Fill the octets of the outer atoms. Keep track of the electrons: 26  6 = 20  18 = 2

  26. Writing Lewis Structures • Fill the octet of the central atom. Keep track of the electrons: 26  6 = 20  18 = 2  2 = 0

  27. Writing Lewis Structures • If you run out of electrons before the central atom has an octet… …form multiple bonds until it does.

  28. Exceptions to Octet Rule • Electron Promotion • group 2 central atom will have 4 electrons • group 13 central atom will have 6 electrons • Valence Level Expansion • group 15 central atom will have 10 electrons • group 16 central atom will have 12 electrons • group 17 central atom will have 14 electrons • group 18 central atom will have 16 electrons

  29. Writing Lewis Structures • Write Lewis Structures for the following molecules:

  30. Coordinate Covalent bonds • is defined as a covalent bond where one atom provides both of the electrons:

  31. Polyatomic Ions • we can use coordinate covalent bond theory to explain most ions:

  32. Resonance This is the Lewis structure we would draw for ozone, O3. + -

  33. Resonance • But this is at odds with the true, observed structure of ozone, in which… • …both O—O bonds are the same length. • …both outer oxygens have a charge of 1/2.

  34. Resonance • One Lewis structure cannot accurately depict a molecule such as ozone. • We use multiple structures, resonance structures, to describe the molecule.

  35. Resonance Just as green is a synthesis of blue and yellow… …ozone is a synthesis of these two resonance structures.

  36. Resonance • In truth, the electrons that form the second C—O bond in the double bonds below do not always sit between that C and that O, but rather can move among the two oxygens and the carbon. • They are not localized, but rather are delocalized.

  37. Resonance • The organic compound benzene, C6H6, has two resonance structures. • It is commonly depicted as a hexagon with a circle inside to signify the delocalized electrons in the ring.

  38. Molecular Shapes • The shape of a molecule plays an important role in its reactivity. • By noting the number of bonding and nonbonding electron pairs we can easily predict the shape of the molecule.

  39. What Determines the Shape of a Molecule? • Simply put, electron pairs, whether they be bonding or nonbonding, repel each other. • By assuming the electron pairs are placed as far as possible from each other, we can predict the shape of the molecule.

  40. Valence Shell Electron Pair Repulsion Theory (VSEPR) “The best arrangement of a given number of electron domains is the one that minimizes the repulsions among them.”

  41. Electron-Domain Geometries • All one must do is count the number of electron domains in the Lewis structure. • The geometry will be that which corresponds to that number of electron domains.

  42. Molecular Geometries • The electron-domain geometry is often not the shape of the molecule, however. • The molecular geometry is that defined by the positions of only the atoms in the molecules, not the nonbonding pairs.

  43. Molecular Geometries Within each electron domain, then, there might be more than one molecular geometry.

  44. Group 2 Geometries • In this domain, there is only one molecular geometry: linear. • NOTE: If there are only two atoms in the molecule, the molecule will be linear no matter what the electron domain is.

  45. Group 13 Geometries • because there are no lone pair electrons the molecular is planar (flat).

  46. Group 14 Geometries

  47. Summary for Groups 2, 13 & 14 • because there are no lone pair electrons: • central atom bonded to 1 atom: linear • central atom bonded to 2 atoms: linear • central atom bonded to 3 atoms: trigonal planar • central atom bonded to 4 atoms: tetrahedral

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