Form an ionic bond between the following

1. Form an ionic bond between the following Na and Cl Cl and F Na and Mg Mg and Br NH4+ OH- NH4+ SO42-

2. Covalent Bonding

3. The Octet Rule The octet rule says that atoms tend to gain, lose or share electrons so they have eight electrons in their outer shell.

4. Covalent Bonding Instead of gaining or losing electrons atoms can get stable by sharing electrons This is always between two non-metals. Two fluorine atoms, for example, can form a stable F2 molecule in which each atom has 8 valence electrons by sharing a pair of electrons. In covalent bonds they can share more than two electrons

5. Single, Double and Triple Bonds • With Covalent bonds the elements can share two or more electrons • A Single Bond is when 2 electrons are shared they are represented by a single line in bond diagrams • A Double bond is when 4 electrons are shared they are represented by two lines in bond diagrams • A Triple bond is when 6 electrons are shared they are represented by three lines in bond diagrams

6. Oxygen has six valence electrons and is seeking an octet. Therefore it combines with two hydrogen atoms, each with one valence electron.

7. Representing covalent molecules • Chemists use two types of diagrams to represent molecules: 1. Electron dot formula (Lewis structure) – each outer shell (valence) electron is represented by a dot. 2. Valence structure – lines are used to represent electrons: one line represents one pair of electrons.

8. Lewis Dot Structure In order to draw covalent bonds you must be able to write Lewis Dot structure This shows the number of valence electrons around the symbol Put one electron on each side before doubling up. We can see that helium does not follow this rule exactly.

9. First, we must determine whether CO2 is a covalent or ionic bond. Because nonmetals form covalent bonds, CO2 will be covalent. Ionic bonds form when combining both metals and a nonmetals. Begin by drawing the Lewis dot structure of each individual atom. How many electrons does the carbon want? 4 How many electrons does each oxygen want? 2

10. Determine which is the central atom All three of these atoms want an octet. Carbon needs to “share” four electrons. Carbon will share two of its electrons with the oxygen on the right and two with the oxygen on the right.

11. The carbon dioxide Lewis Dot Structure: :C: • Take note of the boxed in electrons below. :C: • Now each atom has eight electrons. What type of bonds does CO2 have? Double bonds

12. Draw and complete the following table Include the following compounds CS2, NH3 , CF4.

13. Naming Covalent formulas First name generally stays the same. Second non metal name changes to –ide. If there is one or more of the second element it will have prefix in front e.g. Mono (means 1) or di (2) or tri (3) tetra (4) penta (5) e.g. CO2 Carbon Dioxide If there are two or more of the first element it must have a prefix e.g. di (2) or tri (3) tetra (4) penta (5) Dihydrogen sulfide (H2S)

14. Naming Covalent compounds • What is dihydrogen monoxide? • H2O • Name CO • Carbon Monoxide • Name P2S3 • Diphosphourous trisulfide

15. Valence electrons • In covalent bonds when electrons are shared the electrons spend most of their times between the two nuclei of the atoms. The electrons are called localised electrons. • The shared electrons are called the bonding electrons • The valence electrons not involved in bonding are called the non-bonding electrons. • Pairs of non-bonding electrons are called lone pairs. Lone pair Pair of bonding electrons

16. Your turn • How many bonding electrons does Cl2 have? • How many non-bonding electrons does Cl2 have • How many lone pairs does Cl2 Have?

17. Covalency • Chlorine needs to share one electron to give it a stable outer shell. Therefore its has a covalency of one. • The covalency of an atom is generally the number of electrons it shares when bonding with another non metal atom.

18. Draw and Complete the following table Include the following compounds HCl,CS2, NH3 , CF4.

19. Charge Cloud Model • This Model of the atom pictures the electrons moving around the nucleus in a region called an electron cloud. • The electron cloud is a cloud of varying density surrounding the nucleus. The varying density shows where an electron is more or less likely to be. • an 'electron cloud' may be a single, double or triple bond, or a lone pair of electrons

20. What is the Electron Cloud Model?Diagram 1:

21. What is the Electron Cloud Model?Diagram 2:

22. Shapes of molecules • The shape of a molecule describes the way in which the atoms are arranged • Molecules can be linear, angular v-shaped or bent, triangular pyramid and tetrahedral. • There are more shapes but you only need to the know the above http://www.youtube.com/watch?v=FhVkCH9COZo http://www.youtube.com/watch?v=q6g7BUruUK0&feature=related http://www.youtube.com/watch?v=ra0A8Ni1DAA&feature=related

24. Activity Trigonal Planar 0 120° sulfur trioxide The molecule is all in the same plane and is two dimensional. • Use Balloons and string to make the shape of a molecule you have been assigned • Use the molecular models to make all the other shape molecules

25. x

26. ValenceShellElectronPairRepulsionTheory Trigonal Planar Tetrahedral Trigonal bipyramidal Octahedral

27. The valence shell electron pair repulsion (VSPER) model • VSPER is used for predicting the shapes of individual molecules, based upon the extent of electron-pair repulsion (to push away). • a lone pair of electrons is a non-bonding pair of electrons • 'electron clouds' are negatively charged since the electrons are negatively charged, so electron clouds repel (push away) from one another and try to get as far away from each other as possible.

28. The valence shell electron pair repulsion (VSPER) model RULES 1. Lone electron pairs in the outer shell of an atom in a molecule will repel each other the most due to their negative charge 2. lone pairs of electrons exert a greater repelling effect than bonding pairs do 3. lone pair-bonding pair repulsion is greater than bonding pair-bonding pair repulsion lone pair-lone pair repulsion > lone pair-bonding pair repulsion > bonding pair-bonding pair repulsion

29. How to determine shapes of molecules 1. Decide which atom is the central atom 2. Determine the number of valence electrons for each atom 3. Arrange the electrons to include bonding 4. Draw the bonding electrons and lone pairs so they are as far apart as possible 5.Count the total number of electron pairs around the central atom, including both the bonding pairs and lone pairs. Double and triple bonds are counted as on bonding pair. 6. Use the VSEPR geometry table to predict the geometry of the molecule (it can only be one of 5 shapes!!).

30. How to determine shapes of molecules • Sometimes the molecules are represented by AXY, A is the central atom, X is the 2nd atom and Y is the # of peripheral atoms • AX2 = linear or bent • AX3 = trigonal planar or trigonal pyrimadial • AX4 = tetrahedral (tetra = 4 faces) • AX5 = trigonal bipyramidal (2 pyramids) • AX6 = octahedral (octa = 8 faces)

31. Which shapes do you recognise? • x

32. Yourturn • Use Lewis structures, VSPER and the molecular models to help you determine the shape of the following molecules. • Carbon tetrachloride • Silicon disulfide • Dicarbon dihydride • Boron trihydride • Phosphorous trifluoride

33. What is Electronegativty? • Electronegativity: is a measure of an elements ability to attract electrons • The electronegativity of elements tends to increase from left to right across a period. • Electronegativities decrease from top to bottom within a group. • Fluorine is the most electronegative atom

34. Electronegativity Values • The electronegativity values can be found in the periodic table • The higher the value the higher the electrotnegativity • The Pauling scale is used to measure electronegativity. It is a relative scale running from 0.7 to 4.0 (hydrogen = 2.2). • The units for electronegativity are Pauling units.

35. x

36. There are two types of covalent bonding 1. Non-polar bonding with an equal sharing of electrons. 2. Polar bonding with an unequal sharing of electrons. The number of shared electrons depends on the number of electrons needed to complete the octet.

37. 1. Non Polar Bonds • What does polarised/polarity mean? • Opposites • Non polar bonding results when two identical non-metals equally share electrons between them. • One well known exception to the identical atom rule is the combination of carbon and hydrogen in all organic compounds. • Non polar bonding exists between carbon and hydrogen

38. Iodine (I2) and oxygen (O2) forms a diatomic (di – two, atomic – atoms) non-polar covalent molecule.

39. x

40. 2. Polar Bonds • Polarised bonds occur in covalent molecules. One atom has greater electronegativity and therefore has a partial (slight) negative charge because it attracts the bonding electrons more strongly (remember electrons have a –ive charge). • The other atom has a slight positive charge.

41. Partial positive is represented by this symbol Partial negative is represented By this symbol

42. x

43. Which is more electronegative H or Cl? • Cl has greater electronegativity • The bonding electrons therefore spend more time near the Cl atom • So the Cl side of the molecule becomes slightly negative (partial negative charge) and the H side becomes slightly positive making a polarised bond • What about HF?

44. Polar molecules • Polar molecules are those containing polar bonds that do not cancel each other out • Molecular polarity is dependent on the difference in electronegativity between atoms in a compound and the asymmetry of the compound's structure. • For example, a molecule of water is polar because of the unequal sharing of its electrons in a "bent" structure, whereas methane is considered non-polar because the carbon shares the hydrogen atoms uniformly.

45. Polar • The covalent bond between HCl is said to be polarised and the molecule is called a polar molecule • It is also said to be a dipole • because is has two charged ends

46. Non polar molecules • Non polar molecules are those that contain only non-polar bonds OR • Molecules in which polar bonds cancel each other out

47. Youtube clip • http://www.youtube.com/watch?v=GQwpb5jyMq8

48. Quiz • What happens to the electrons in a polar bonds? • How do you know if a bond is polar? • What is a non polar molecule? • What is a dipole? • What is the symbol for slightly negative? • Can non polar molecules have polar bonds?

49. How to determine polar molecules There are two important factors 1. The polarity of the individual bonds in the molecule; 2. The shape or geometry of the molecule. Steps to take • Determine if a given individual bond is polar, Look at the difference between electronegativity of the atoms in the perioidc table. If the difference is: 0.3 < non polar 0.3 -1.7 = polar

50. b) Determine the shape of molecule. Then there are 3 option i) if all bonds are non-polar, then the whole molecule is non-polar regardless of its shape. ii) If there is symmetry in the molecule so that the polarity of the bonds cancels out, then the molecule is non-polar. (symmetry arround the central atom) iii) If there are polar bonds but there is no symmetry the overall molecule is polar.