1. 4.2 Covalent Bonding

2. Assessment Statements 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.

3. Assessment Statements 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 C 60 fullerene) 4.2.10 Describe the structure of and bonding in silicon and silicon dioxide

4. Assessment Statement 4.2.1 Describe the covalent bond as the result of electron sharing.

5. Covalent compounds Covalent compounds are formed when non-metal atoms react together. As these atoms come near their outer electrons are attracted to the nucleus of both atoms and become shared by the atoms. The shared electrons count towards the shells of both atoms and therefore help fill up incomplete electron shells.

6. Covalent bonds Covalent compounds are held together by this sharing of electrons. covalent bond. A pair of electrons shared in this way is known as a covalent bond. It is sometimes represented in full bonding diagrams (see figure 1). Often these bonds are just shown as a pair of electrons (xx) or even just a line (see figure 2). F XXXX F F F - Figure 1Figure 2

7. Small covalent structures Sometimes just a few atoms join together in this way. This produces small covalent molecules – often known as simple molecular structures. a simple molecular structure covalent bonds

8. Giant covalent structures Sometimes millions of atoms are joined together by covalent bonds. giant lattice. This produces a rigid 3-D network called a giant lattice. a giant lattice covalent bonds

9. Covalent bonding and electron structures The driving force for covalent bonding is again the attainment of outer electron shells that are completely full. This is achieved by sharing electrons where the shared electrons count towards the outer shells of both atoms. Sometimes this is achieved with equal numbers of each type of atom. Sometimes it is not! Cl C H HH H N H HH H

10. Assessment Statement 4.2.2 Draw the electron distribution of single and multiple bonds in molecules

11. Covalent bonding in chlorine Chlorine (2.8.7) needs 1 more electron to attain a full electron shell. Cl (2,8,7) Cl ( 2,8,7 ) Cl (2,8,8) Cl (2,8,8) Cl-Cl

12. Both fluorine and chlorine needs 1 more electron to attain a full electron shell. Cl (2,8,7) F ( 2,7 ) Copy this diagram and add the electron arrangements that could exist in fluorine chloride (FCl). Cl (2,8,8) F ( 2,8 )

13. Covalent bonding in hydrogen chloride Both hydrogen (1) and chlorine (2.8.7) needs 1 more electron to attain a full outer shell. H (2) Cl (2,8,8) H-Cl Cl (2,8,7) H (1)

14. Covalent bonding in water Hydrogen (1) needs 1 more electron but oxygen (2.6) needs 2 more. Therefore, we need 2 hydrogens. O H H O H H O H H

15. Hydrogen (1) needs 1 more electron. How many does nitrogen (2.5) need? How many hydrogens per 1 nitrogen? Draw bonding diagrams for ammonia. N H H H N H H H 3 3

16. Hydrogen (1) needs 1 more electron. How many does carbon (2.4) need? How many hydrogens per 1 carbon? Draw bonding diagrams for methane. 4 4 C H H H H C H H H H

17. H H O O H H O O Copy the atoms below. Complete the diagram showing how each atom can achieve full shells.

18. Covalent bonding - multiple bonds Mostly electrons are shared as pairs. There are some compounds where they are shared in fours or even sixes. This gives rise to single, double and triple covalent bonds. Again, each pair of electrons is often represented by a single line when doing simple diagrams of molecules. Cl-Cl Single bond O=O Double bond N=N Triple bond

19. Covalent bonding in oxygen Oxygen (2.8.6) needs 2 more electrons to attain a full electron shell. O O O=O O O 4 electrons

20. Nitrogen (2.8.5) needs 3 more electrons to attain a full electron shell and forms a triple bond. Draw a bonding diagram of nitrogen. 6 electrons N N N N N=N

21. 1. Hydrogen fluoride (HF) 2. Hydrogen sulphide (H 2 S) 3. Ethane (C 2 H 6 and the carbons are joined by a single covalent bond) 4. Carbon dioxide (CO 2 and the carbon oxygen bonds are double bonds) H F H H S H H H H H H CC CO O Draw ‘dot and cross’ type bonding diagrams for each of the following:

22. Pure covalent bonds Sharing of electrons between two or more of the same type of non-metal atoms. H O Br F I N Cl elements are all covalently bonded. H 2, O 2, Br 2, F 2, I 2, N 2, Cl 2

23. Pure covalent bonds Equal sharing of electrons when forming the bond H 2 (g) forms a single bond (shared pair)

24. Assessment Statement 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

25. Another Term you need to Know Electronegativity (symbol χ)χ chemical property that describes the tendency of an atom to attract electrons towards itself. chemical propertyatomelectrons An atom's electronegativity is affected by both its atomic number and the distance that its valence electrons reside from the charged nucleus.atomic number valence electrons The higher the associated electronegativity number, the more an element or compound attracts electrons towards it.

27. Predicting Ionic or Covalent When the difference between the electronegativities of the elements in a compound is relatively large, the compound is best classified as ionic. When the electronegativities of the elements in a compound are about the same, the atoms share electrons, and the substance is covalent. Inevitably, there must be compounds that fall between these extremes. (Which means that the electrons will be shared UNEQUALLY)

28. 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)

29. 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: P 2 O 4 = diphosphorous tetroxide PrefixNumber Mono1 Di2 Tri3 Tetra4 Penta5 Hexa6 Hepta7 Octa8 Nona9 Deca10

30. Assessment Statement 4.2.3 Deduce the Lewis structures of molecules and ions for up to 4 electron pairs on each atom.

31. 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)

32. Drawing Lewis StructuresLewis Structures 1. Look at valence electrons of all atoms 2. Pick a central atom (least electronegative usually, has most bonding sites) 3. Align all atoms so that each have their ideal amount of valence electrons achieved through sharing.

33. Writing Dot Structures Writing Dot structures is a process: 1. Determine the number of valence electrons each atom contributes to the structure 2. The number of valence electrons can usually be determined by the column in which the atom resides in the periodic table 33

34. Writing Dot Structures 3. Add up the total number of valence electrons 4. Adjust for charge if it is a poly atomic ion Add electrons for negative charges Reduce electrons for positive charges Example SO 3 2-  1 S = 6 e  3 0 = 6x3 = 18 e  (2-) charge = 2 e --------- Total = 26 e 34

35. Electron Dot Structures 5. Make the atom that is fewest in number the central atom. 6. Distribute the electrons so that all atoms have 8 electrons. 7. Use double or triple pairs if you are short of electrons 8. If you have extra electrons put them on the central atom 35

36. Electron Dot Structures Example 2: SO 3  1 S = 6 e  3 O = 6x3 = 18 e  no charge = 0 e --------- Total = 24 e Note: a double bond is necessary to give all atoms 8 electrons 36

37. Electron Dot Structures Example 3: NH 4 +  1 N = 5 e-  4 H = 4x1 = 4 e-  (+) charge = -1 e- --------- Total = 8 e- Note: Hydrogen atoms only need 2 e- rather than 8 e- 37

38. Example -- Carbon Dioxide CO 2 1. Central atom = 2. Valence electrons = 3. Form bonds. 4. Place lone pairs on outer atoms. This leaves 12 electrons (6 pair). 5.Check to see that all atoms have 8 electrons around it except for H, which can have 2. C 4 e - O 6 e - x 2 O’s = 12 e - Total: 16 valence electrons

39. Carbon Dioxide, CO 2 There are too many electrons in our drawing. We must form DOUBLE BONDS between C and O. Instead of sharing only 1 pair, a double bond shares 2 pairs. So one pair is taken away from each oxygen atom and replaced with another bond. C 4 e- O 6 e- X 2 O’s = 12 e- Total: 16 valence electrons How many are in the drawing?

40. Violations of the Octet Rule Violations of the octet rule usually occur with B and elements of higher periods. Some common examples include: Be, B, P, S, and Xe. BF 3 SF 4 Be: 4 B: 6 P: 8 OR 10 S: 8, 10, OR 12 Xe: 8, 10, OR 12

41. 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)

42. Some examples

43. Practice drawing and naming Lewis Structures H 2 O CH 2 O

45. Tricky ones! Try ozone O 3

46. Lewis Dot Stucture for SO 3 The diagram below shows the dot structure for sulfur trioxide. The bonding electrons are in shown in red and lone pairs are shown in blue. 46

47. 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

48. Practice with a cation

49. 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.

50. Try ammonia (NH3) What about if I add a hydrogen ion to make the ammonium ion (NH4+)? What do you notice has to happen? (think about where the electrons are coming from to make this bond).

51. 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.

52. Examples Hydronium (H 3 O + ) Carbon monoxide (CO)

53. 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

54. 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 NO 2

55. Assessment Statement 4.2.4 State and explain the relationship between the number of bonds, bond length and bond strength. 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.

56. 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)

57. The Shapes of Molecules The shape of a molecule has an important bearing on its reactivity and behavior. The shape of a molecule depends a number of factors. These include: 1.Atoms forming the bonds 2.Bond distance 3.Bond angles 57

58. Bonding Electrons and Lone Pairs In a molecule some of the valence electrons are shared between atoms to form covalent bonds. These are called bonding electrons. Other valence electrons may not be shared with other atoms. These are called non- bonding electrons or they are often referred to as lone pairs. 58

59. 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

60. VSEPR: Predicting the shape Once the dot structure has been established, the shape of the molecule will follow one of basic shapes depending on: 1. The number of regions of electron density around the central atom 2. The number of regions of electron density that are occupied by bonding electrons 60

61. VSEPR: Predicting the shape The number of regions of electron density around the central atom determines the electron skeleton The number of regions of electron density that are occupied by bonding electrons and hence other atoms determines the actual shape 61

62. Basic Molecular shapes The most common shapes of molecules are shown at the right 62

63. Linear Molecules Linear molecules have only two regions of electron density. 63

64. Angular or Bent Angular or bent molecules have at least 3 regions of electron density, but only two are occupied 64

65. Triangular Plane Triangular planar molecules have three regions of electron density. All are occupied by other atoms 65

66. Tetrahedron Tetrahedral molecules have four regions of electron density. All are occupied by other atoms 66

67. Trigonal Bipyramid A few molecules have expanded valence shells around the central atom. Hence there are five pairs of valence electrons. The structure of such molecules with five pairs around one is called trigonal bipyramid. 67

68. Octahedron A few molecules have valence shells around the central atom that are expanded to as many as six pairs or twelve electrons. These shapes are known as octahedrons 68

69. Chemistry SL Shapes Sets ( group of bonding pairs) Lone Pairs Shape 20Linear 180 o 22Bent 104.5 o 30Triagonal Planar 120 o 31Pyramidal 107.3 o 40Tetrahedral 109.5 o

70. Examples Arrangement of electron pairs on central atom Number of bonding electron pairs Example Linear2BeCl 2 Planar triangular3BCl 3 Tetrahedral4CH 4

72. Practice Lewis structure and state the shape SO 2 SO 3 [SO 4 ] - 2 AsCl 3 SI 2 CH 3 F CH 2 F 2 NH 4 + NO 2 - NO 2 + H 3 O +

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

74. Polar Bonding Not this kind…………….

75. Polarity 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

76. Molecular Polarity Molecular Polarity depends on: 1. the relative electronegativities of the atoms in the molecule 2. The shape of the molecule 3. Molecules that have symmetrical charge distributions are usually non-polar 76

77. Non-polar Molecules The electron density plot for H 2. Two identical atoms do not have an electronegativity difference The charge distribution is symmetrical. The molecule is non-polar. 77

78. Polar Molecules The electron density plot for HCl Chlorine is more electronegative than Hydrogen The electron cloud is distorted toward Chlorine The unsymmetrical cloud has a dipole moment HCl is a polar molecule. 78

79. Molecular Polarity To be polar a molecule must: 1. have polar bonds 2. have the polar bonds arranged in such a way that their polarity is not cancelled out 3. When the charge distribution is non- symmetrical, the electrons are pulled to one side of the molecule 4. The molecule is said to have a dipole moment. HF and H 2 O are both polar molecules. CCl 4 is non-polar 79

80. Polar bonds do not guarantee a polar molecule Ex: CCl 4 and CO 2 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

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

83. Summary of Polarity of Molecules Linear: When two atoms attached to central atom are the same, the molecule will be Non-Polar (CO 2 ) 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 (H 2 O)

84. 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 (NH 3 ) Trigonal Planar: When the three atoms attached to central atom are the same, the molecule will be Non- Polar (BF 3 ) When the three atoms are different the dipoles will not cancel, resulting in a net dipole, and the molecule will be Polar (CH 2 O)

85. 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

86. Summary of Polarity of Molecules

87. Examples to Try Determine whether the following molecules will be polar or non-polar SI 2 CH 3 F AsI 3 H 2 O 2

88. Angular = benttriangular pyramid = pyramidal

90. 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? Try This with water………….

91. 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 F 2 is –188° C). Polar solvents dissolve ionic and polar molecules more efficiently than non-polar solvents

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

93. 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)

94. Bond Length and Bond strength

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

97. 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.

98. Assessment Statement 4.2.9 Describe and compare the structure and bonding in the 3 allotropes of carbon (diamond, graphite and C 60 fullerene)

99. Giant covalent structures 1. Carbon atoms form giant structures. 2. What is interesting is that there is more than one possible arrangement for the atoms. 3. Although this does not affect the chemical properties it can make a huge difference to the physical properties such as hardness, slipperiness, melting point and density. Different arrangements of the same element are called allotropes. C

100. 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 (C 60 )

101. 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

102. 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.

103. Giant covalent structures: graphite A more common form of carbon is graphite. Millions of carbon atoms are bonded into a giant structure but within this structure the layers are only weakly joined.

104. Fullerene C 60 consists of 60 carbon atoms bonded in the nearly spherical configuration C 60 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)

105. Giant covalent structures: carbon footballs! During the last 20 years new forms of carbon have been discovered some of which have “closed cage” arrangements of the atoms. These are large but are not really giant molecules. One of them contains 60 carbon atoms and bears remarkable similarities to a football!

106. Giant covalent structures: diamond One form of carbon is diamond. Each diamond consists of millions of carbon atoms bonded into a single giant structure. very It is very hard. Diamond strong covalent bonds carbon atoms

107. Assessment Statement 4.2.10 Describe the structure of and bonding in silicon and silicon dioxide

108. Silicon Has almost identical crystal structure to diamond

109. Silicon dioxide Sometimes called silica Occurs as quartz and sand Oxygen atoms bridge the silicon atoms

110. Bibliography and good sites http://www.chemguide.co.uk/atoms/bonding/dative.h tml http://www.chemguide.co.uk/atoms/bonding/dative.h tml http://en.wikipedia.org/wiki/Coordinate_covalent_bo nd http://en.wikipedia.org/wiki/Coordinate_covalent_bo nd http://en.wikipedia.org/wiki/Diamond Use links to find out about fullerenes