4.2 Covalent Bonding

1. 4.2 Covalent Bonding 4.2.1 Describe the covalent bond as the result of electron sharing. 4.2.2 Draw the electron distribution of single and multiple bonds in molecules 4.2.3 Deduce the Lewis structures of molecules and ions for up to 4 electron pairs on each atom. 4.2.4 State and explain the relationship between the number of bonds, bond length and bond strength. 4.2.5 Predict whether a compound of two or more elements would be covalent from the position of the elements in their periodic table or from their electronegativity values. 4.2.6 Predict the relative polarity of bonds based on electronegativity values 4.2.7 Predict the shape and bond angles for molecules with four charge centres on the central atom. 4.2.8 Predict molecular polarity based on bond polarity and molecular shape. 4.2.9 Describe and compare the structure and bonding in the 3 allotropes of carbon (diamond, graphite and C60 fullerene) 4.2.10 Describe the structure of and bonding in silicon and silicon dioxide

2. Pure covalent bonds • Sharing of electrons between two or more of the same type of non-metal atoms. • HOBrFINCl elements are all covalently bonded. • H2, O2, Br2, F2, I2, N2, Cl2

3. Pure covalent bonds • Equal sharing of electrons when forming the bond • H2(g) forms a single bond (shared pair)

4. Polar covalent bond • Unequal sharing of electrons. • One atom will have a higher electronegativity than the other, so it will “pull” the shared electrons closer to itself making that atom slightly more negative than the other. • The Cl (3.00) is more negative than the H (2.20)

5. Naming simple molecules • Must memorize the prefixes • RULES: if there is only one of the first atom than don’t use a prefix, otherwise use a prefix. • Ex: CO = carbon monoxide • Ex: P2O4 = diphosphorous tetroxide

6. Chemical structures • Need to show the structure of a molecule. • Will use Lewis structures (electron dot diagrams) to show where there are lone pairs (filled orbitals) and bonding pairs (places where bonds most likely occur)

7. Drawing Lewis Structures • Look at valence electrons of all atoms • Pick a central atom (least electronegative usually, has most bonding sites) • Align all atoms so that each have their ideal amount of valence electrons achieved through sharing.

8. Carbon tetrachloride • Carbon is the central atom. • It has 4 bonding pairs. • Chlorine wants to share one bonding site each. • Need 4 chlorines for every one carbon (Cl has 3 lone pairs and 1 bonding pair)

9. Some examples

10. Practice drawing and naming Lewis Structures • H2O • CH2O

12. Tricky ones! • Try ozone O3

13. What about ions? • Count up all valence electrons that you are allowed to place. • Still pick the central atom. • Still have the correct number of electrons around each atom (usually 8, except for H and He) • Add extra electrons if an anion and take away electrons if a cation

14. Practice with a cation

15. Practice with an anion Oxygen has an unshared pair of electrons, but since this is an anion it receives an extra electron which will fill up the outer orbital.

16. Coordinate covalent bonds (dative) • A covalent bond that occurs between two atoms in which both electrons shared in the bond come from the same atom. • Both electrons from the nitrogen are shared with the upper hydrogen • Ammonium has 3 polar covalent bonds and 1 coordinate (dative) covalent bond.

17. Hydronium (H3O+) Carbon monoxide (CO) Examples

18. Free Radicals • A molecule with an odd amount of electrons, or a broken bond causing a particle with an uneven amount of electrons • Free radicals are very unstable and react quickly with other compounds, trying to capture the needed electron to gain stability, but causing a new free radical to form in the process. • It’s a chain reaction which usually involves the destruction of living cells • Vitamin E (fat soluble) and C (water soluble)are antioxidants which are able to neutralize the damage by ‘donating’ an electron causing the chain to stop

19. Free Radicals • NO is usually a slow reaction with nitrogen and oxygen gases, but can occur more quickly in the presence of a catalyst or high temperatures • NO is a common free radical that is primarily found due to internal combustion engines (car exhaust). • Cars have catalytic converters to reverse the reaction (decompose NO) • It reacts to form nitric acid, causing more problems with acid rain, and reacts with ozone to produce NO2

20. VSEPR • Valence shell electron pair repulsion theory • Bonding pairs and lone pairs around an atom in a molecule adopt positions where their mutual interactions are minimized. • Electron pairs are negatively charged and will get as far apart from each other as possible. (Same charge = repulsion)

21. Bond angles • Lone pairs occupy more space than bonding electron pairs. • Double bonds occupy more space than single bonds. • LP-LP > LP-BP > BP-BP • Lone pairs are more repulsive than bonding pairs

22. Chemistry SL Shapes

23. Examples

25. SO2 SO3 [SO4 ]-2 AsCl3 SI2 CH3F CH2F2 NH4+ NO2- NO2+ H3O+ Practice Lewis structure and state the shape

26. Advanced structural drawings (3 D) • The dashed wedge = bond going back • Solid wedge = bond going forward • Unbroken line = plane of the paper

27. Polarity and shape • The shape of the molecule directly influences the overall polarity of the molecule. • If there is symmetry the charges cancel each other out, making the molecule non-polar • If there is no symmetry, then its polar

28. Polar bonds do not guarantee a polar molecule • Ex: CCl4 and CO2 both have polar bonds, but both are NON-POLAR molecules. They have a dipole moment of zero • The greater the dipole moment, the more polar the molecule

29. The symetry of the molecule Cancels out the “charges” Making this NON-POLAR No overall DIPOLE The bent shape creates an overall positive end and negative end of the molecule = POLAR

31. Summary of Polarity of Molecules • Linear: • When two atoms attached to central atom are the same, the molecule will be Non-Polar (CO2) • When the two atoms are different the dipoles will not cancel, and the molecule will be Polar (HCN) • Bent: • The dipoles created from this molecule will not cancel creating a net dipole moment and the molecule will be Polar (H2O)

32. Summary of Polarity of Molecules • Pyramidal: • The dipoles created from this molecule will not cancel creating a net dipole and the molecule will be Polar (NH3) • Trigonal Planar: • When the three atoms attached to central atom are the same, the molecule will be Non-Polar (BF3) • When the three atoms are different the dipoles will not cancel, resulting in a net dipole, and the molecule will be Polar (CH2O)

33. Tetrahedral When the four atoms attached to the central atom are the same the molecule will be Non-Polar When three atoms are the same, and one is different, the dipoles will not cancel, and the molecule will be Polar

34. Summary of Polarity of Molecules

35. Examples to Try • Determine whether the following molecules will be polar or non-polar • SI2 • CH3F • AsI3 • H2O2

36. Angular = bent triangular pyramid = pyramidal

38. Testing a liquid’s polarity • As the liquid is flowing bring a magnetically charged object close. • If the stream of liquid is attracted to the rod, it is polar • If the stream is unaffected, it is non-polar. • Can we explain why this would happen?

39. Why is molecular polarity important? • Polar molecules have higher melting and boiling points (for example the BP of HF is 19.5° C, and the BP of F2 is –188° C). • Polar solvents dissolve ionic and polar molecules more efficiently than non-polar solvents

40. Covalent bond strength • Two forces operating: • increased overlap of atomic orbitals (better sharing) brings atoms together • closer distance between nuclei increases positive-positive charge repulsion • balance of these forces = its bond length • Measured in pm (10-12 m) or Ǻ(10-10 m)

41. In a molecule as you increase the number of electrons shared between two atoms (from single to double to triple bond), you increase the bond order, increase the strength of the bond, and decrease the distance between nuclei. • Bond strength is measured by how much energy it takes to break the bond (kJ/mol)

42. Bond Length and Bond strength

43. Bond enthalpy (energy needed to break the bond as a gas)

45. Properties of molecules • The forces between discrete molecules are relatively weak (Intermolecular forces) so • Low boiling points and melting points • Quite soft if solid • Do not conduct electricity • Tend to be more soluble in non-polar solvents than polar solvents.

46. Allotropes of carbon • elements can exist in two or more different forms because the element's atoms are bonded together in a different manner • Carbon has 3 allotrophes • Diamond • Graphite • Fullerenes (C60)

47. Diamonds • carbon atoms are bonded together in a tetrahedral lattice arrangement (3D framework) • Giant covalent structure • Very strong, so they require a lot of energy to break them • M.P is 3820 K • Does NOT conduct electricity • 4x harder than any other natural mineral

48. Graphite • has a sheet like structure where the atoms all lie in a plane and are only weakly bonded to the sheets above and below. (2D framework) • Much softer, conducts electricity. • The C-C bonds are still quite strong.

49. Fullerene C60 • consists of 60 carbon atoms bonded in the nearly spherical configuration • C60 is highly electronegative, meaning that it readily forms compounds • it is a yellow powder which turns pink when dissolved in certain solvents such as toluene. • Also includes nanotubes (cylindrical)

50. Silicon • Has almost identical crystal structure to diamond