2. Bonding K Warne ClH X ++ --

3. Bonding Objectives: At the end of this unit you should be able to:- Explain how metallic bonding determines the prosperities of metals State/explain (understand) the significance of valence electrons State the conditions for covalent bonding. Explain the properties of substances (simple and giant covalent) in terms of their bonding and structure. Know (state) conditions for ionic bonding. Name chemical compounds correctly. List the characteristics of different states of matter.

4. Covalent bond A shared PAIR of electrons.  Formed between.........................  Common Ions KNOW formulae; eg sulphate ion SO 4 2-  Diatomic Molecules;.... 2,.... 2,.... 2,.... 2,.... 2,..... 2,...... 2,  Pure covalent bonds have.....................SHARING of the electrons. In covalent substances all the electrons are strongly held in the bonds and so the substance............................ conduct electricity. H Single hydrogen atom Hydrogen molecule A shared pair of electrons = a single covalent bond

5. Covalent bond A shared PAIR of electrons.  Formed between non metals.  Common Ions KNOW formulae; eg sulphate ion SO 4 2-  Diatomic Molecules; H 2, O 2, F 2, Cl 2, Br 2, I 2, N 2,  Pure covalent bonds have EQUAL SHARING of the electrons. HH In covalent substances all the electrons are strongly held in the bonds and so the substance will NOT conduct electricity. H x H H Single hydrogen atom Hydrogen molecule A shared pair of electrons = a single covalent bond This is a PURE COVALENT, SINGLE BOND!

6. O x x x x xx O x x x x xx O x x x x xx O x x x x xx O x x x x x x O x x x x x x “Dot Cross Diagrams” - Lewis & Couper Notation Lewis Diagrams O Couper Notation Chemical Formulae O Name: Oxygen Multiple Bonds: Two atoms can share more than one pair of electrons. Draw similar diagrams for all the other diatomic molecules.

7. O x x x x xx O x x x x xx O x x x x xx O x x x x xx O x x x x x x O x x x x x x “Dot Cross Diagrams” - Lewis & Couper Notation Lewis Diagrams O=OO=O Couper Notation Chemical Formulae O 2 Name:Oxygen Multiple Bonds: Two atoms can share more than one pair of electrons. Nitrogen (N 2 ) has a triple bond draw Lewis & Couper diagrams for nitrogen. Draw similar diagrams for all the other diatomic molecules.

8. N2N2

9. N2N2

10. Diatomic Molecules F 2, Cl 2, Br 2,

11. Diatomic Molecules F 2, Cl 2, Br 2,

12. Covalent Molecules CH 4, H 2 O, NH 3, CO 2, NH 4 +,

13. Covalent Molecules CH 4, H 2 O, NH 3, CO 2, NH 4 +,

14. O ● ● ● ● ● ● x x + 

15. O ● ● ● ● ● ● x x +  O ● ● ● ● ● ● x x

16. Methane - CH 4 All the bonds are identical and the molecule has a TETRAHEDRAL SHAPE

19. Fluorine oxide (OF 2 ) O X X X X X X F        Fluorine atom Oxygen atom

20.   F        F      O X X X X X X Fluorine oxide (OF 2 ) F F O By sharing pairs of electrons all bonding atoms now effectively have a full outer shell (8 electrons). O X X X X X X F        Lewis structure Couper Structure Fluorine atom Oxygen atom

21. Boron tri fluoride (BF 3 )

22. F F B By sharing pairs of electrons all bonding atoms now effectively have a full outer shell (8 electrons). Three shared pairs Trigonal Planar structure B X X X F        F          F      F

24.    Cl            P X X X X X        F        F        S F        F        F        F       

32. H CN

34. O x H x xx x Cl S xx x xx x O       O       S O O Two double bonds SO 2 Lewis structureSO 2 Couper structure Be F x F x F F Two shared pairs Linear shape

35. Triple bonds

37. Co-ordinate bonding Co-ordinate or Dative covalent bonding

38. Co-ordinate bonding Co-ordinate or Dative covalent bonding

39. Lewis acid & base Try and draw the other two and identify the coordinate bonds. H 3 NBF 3 Cu(NH 3 ) 4 + Cu(H 2 O) 6 2+

40. Lewis acid & base Try and draw the other two and identify the coordinate bonds. H 3 NBF 3 Cu(NH 3 ) 4 + Cu(H 2 O) 6 2+ Cu

41. Cu(NH 3 ) 4+

42. Electronegativity The ability/power to attract electrons in a bond.

43. Electronegativity in a Group H Li Na Group 1 Electronegativity ………………… from TOP to BOTTOM in a group as the number of ………… increase bonding electrons (outer) are ………… from nucleus and therefore ………… strongly attracted. Electronegativity DECREASES

44. Electronegativity in a Group H Li Na Group 1 Electronegativity DECREASES from TOP to BOTTOM in a group as the number of shells increase bonding electrons (outer) are further from nucleus and therefore LESS strongly attracted. Electronegativity DECREASES

45. Electronegativity Trends H He Li Be B C N OFNe Na Mg Al Si P S Cl Ar Electronegativity nucleus bonding electrons (outer) are more strongly attracted. Electronegativity ………………..from LEFT to RIGHT as the number of protons in the nucleus …………………….and bonding electrons (outer) are more strongly attracted. Group 123 4 5 6 7 8 Electronegativity bonding electrons (outer) are LESS strongly attracted. Electronegativity DECREASES from TOP to BOTTOM in a group as the number of shells increase bonding electrons (outer) are LESS strongly attracted.

46. Electronegativity Trends H He Li Be B C N OFNe Na Mg Al Si P S Cl Ar Electronegativity nucleus bonding electrons (outer) are more strongly attracted. Electronegativity INCREASES from LEFT to RIGHT as the number of protons in the nucleus INCREASES and bonding electrons (outer) are more strongly attracted. Group 123 4 5 6 7 8 Electronegativity bonding electrons (outer) are LESS strongly attracted. Electronegativity DECREASES from TOP to BOTTOM in a group as the number of shells increase bonding electrons (outer) are LESS strongly attracted.

47. Electronegativity

48. VALENCY – BOHR DIAGRAMS Valencyelectrons ……..….. ……….... Valency – ……………….. of electrons ……..….. or ……….... to have a FULL valence level. (Outer shell) H He Li Be B C N OFNe Valence electrons Valence electrons – those in ……………. shell. Na Mg Al Si P S ClAr

49. VALENCY – Bonds Formed Valencyelectrons lostgained Valency – number of electrons lost or gained to have a FULL valence level. (Outer shell) = number of bonds formed by an element. H He Li Be B C N OFNe Valence electrons Valence electrons – those in outer shell = group number. METALS NON - METALS

50. Polar Covalent Bond Each side of the molecule has a small charge due to the electrons being …………………………..SHARED.  Chlorine has a …………………..electro negativity than hydrogen. The “  ” symbol (delta) stands for small amount or small change. > This type of bonding exists when there is a relatively large …………………….. in electronegativity between the bonding atoms. A ……………(two poles) has been created. Electron density diagram - more electron density around the chlorine -- ClH X ++

51. Polar Covalent Bond Each side of the molecule has a small charge due to the electrons being UNEQUALLY SHARED.  Chlorine has a higher electro negativity than hydrogen. The “  ” symbol (delta) stands for small amount or small change. > This type of bonding exists when there is a relatively large difference in electronegativity between the bonding atoms. A dipole (two poles) has been created. Electron density diagram - more electron density around the chlorine -- ClH X ++

52. Bond Polarity in Water The oxygen atom has ………... electronegativity so it attracts the electrons more strongly than the hydrogen atoms. O H H -- ++ ++ The water molecule is a ………………………….- it has two oppositely charged “poles”. ++ -- O H O H H ++++ ---- H This unequal sharing of electrons creates a polar molecule has two …………… charged areas in it.

53. Bond Polarity in Water The oxygen atom has greater electronegativity so it attracts the electrons more strongly than the hydrogen atoms. O H H -- ++ ++ The water molecule is a DIPOLE - it has two oppositely charged “poles”. ++ -- O H O H H ++++ ---- H This unequal sharing of electrons creates a polar molecule has two oppositely charged areas in it.

54. O H H ++++ ---- ++++ H Cl ---- ++++ H H H N ---- ++++ ++++ ++++ B Cl Cl Cl ---- ++++ ++++ ++++ B Cl Cl Cl

55. Ionic Bonding Formed when there is a …………. of …………………... Formed between ………….. and …………………. Metals …………………….. and become ……………………... ions - CATIONS. Non metals …………………... and become …………………………. ions - ANIONS. …………………………… between oppositely charged ions bonds the ions together. Na... :Cl: -.. Na +. :Cl:.. Na. + : Cl: --> [Na] + [Cl] -. ELECTROSTATICATTRACTION

56. There is a transfer of electrons. Occurs when metals and non metals.bond Metals lose electrons and become positively charged ions - CATIONS. Non metals gain electrons and become negatively charged ions - ANIONS. Electrostatic attraction between oppositely charged ions bonds the ions together. Na... Cl: - :Cl: -.. Na +. :Cl:.. Na. + : Cl: --> [Na] + [Cl] -. ELECTROSTATICATTRACTION Ionic Bonding Sodium atom Sodium ion Smaller positive Chlorine atom Chloride ion Negative bigger

61. Ionisation Energy The ENERGY REQUIRED to REMOVE AN ELECTRON completely from an atom in the GAS PHASE. Sodium atom Sodium ion Whenever ionic bonding occurs this process must take place. Gas phase: The atoms are in the gas phase as the energy put in has melted and vapourised them.

62. ELECTRON AFFINITY The amount of ENERGY RELEASED when an electron is added to a gaseous atom. This always accompanies the formation of an ionic bond. e -

63. Bond Type vs Electronegativity WCED Boundaries Δ E neg 00<x<1 1 < x ≤ 2.1 2.1 < X Bond Type Pure Covalent (Non Polar) Covalent Covalent (weakly polar) Polar- Polar-covalent Ionic (In “General Chemistry” Linus Pauling writes: “The farther away two elements are from one another on the scale, the greater is the amount of ionic character of a bond between them. When the separation on the scale is 1.9 the bond has about 50% ionic character. If the separation is greater than this, it would seem appropriate to write an ionic structure for the substance, and if less, to write a covalent structure. No rigid adherence to such a rule is called for however.)

64. Bonding - Metallic Bonding - Exists between metal atoms. - Metal electrons are weakly held - therefore they become delocalized (move from one atom to another). - This leaves a lattice of positive ions - which become surrounded by a ‘sea’ of delocalized electrons. - A force of electrostatic attraction exists between the delocalized electrons and the positive ions which is the metallic bond. All the properties of metals can be explained in terms of this bonding. Since the electrons are weakly held metals CONDUCT electricity.

65. Formation of Ionic Bond A large amount of energy (lattice) is released when the gaseous ions bond together into the ionic crystal lattice. Ionic compounds are therefore very stable and require large amounts of energy to break the bonding. Ionic compounds have HIGH MELTING POINTS we say they are thermally stable. Na (s) + 1 / 2 Cl 2(g) NaCl (s) Na (g) + 1 / 2 Cl 2(g) Na (g) + Cl (g) Na + (g) + ….+ Cl (g) Na + (g) + Cl - (g) …………………Energy ……………….. Energy …………………….Energy Electron ………... ………Energy Born-Haber Cycle

66. Formation of Ionic Bond A large amount of energy (lattice) is released when the gaseous ions bond together into the ionic crystal lattice. Ionic compounds are therefore very stable and require large amounts of energy to break the bonding. Ionic compounds have HIGH MELTING POINTS we say they are thermally stable. Na (s) + 1 / 2 Cl 2(g) NaCl (s) Na (g) + 1 / 2 Cl 2(g) Na (g) + Cl (g) Na + (g) + e - + Cl (g) Na + (g) + Cl - (g) Ionisation Energy Dissociation Energy Sublimation Energy Electron Affinity Lattice Energy Born-Haber Cycle

67. Electrical conductivity anioncation No free moving charges in the solid state. The ions are free to move if acted upon by an electric field. Poslitive electrode Negative electrode IONIC SOLIDS DO NOT CONDUCT ELECTRICITY BUT IONIC LIQUIDS & SOLUTIONS DO IONIC SOLID IONIC LIQUID OR SOLUTION

68. Bonding Summary Covalent Non metals Shared electrons Molecules Ionic Metals + non metals +/- Ions - Lattice electrostatic attraction Metallic Metals “delocalised” electrons H x H Cl - Na + Properties Non - conducting (Electrons held in bond.) V Low or V High melting points Insoluble (H 2 O) Properties High Melting points Soluble (H 2 O) Conduct electricity when ions free to move(liquid or solution). Properties Good Conductors Malleable Ductile Luster (shiny). H-H Eg Hydrogen (H 2 )

69. Dissolution (dissolving) of an Ionic Solid Polar water molecules Dipoles on the water molecules are attracted to the ions in the ionic solid The ionic solid is broken apart by the water molecules

70. Bonding Summary Metallic – bonding between metals  Similar electronegativities (small) – delocalized electrons Covalent - equal sharing of electrons  Similar electronegativities (Large)  Δ E neg < 0.4 Polar covalent - unequal sharing of electrons, dipoles  Polar bonds - Δ E neg < 1.6  Polar molecules: Polar bonds & Asymmetrical shape Ionic - complete transfer of electrons, ions formed,  VERY different electronegativities. Increasing electronegativity DIFFERENCE.

71. Atomic Radius Atoms have linear dimensions. If one considers the atom as a sphere, we can define the radius of that atom as the smallest distance that this atom can approach another atom under a given bonding situation. The atomic radius is determined by the effective volume of the outermost electronic level, and not by the size of the nucleus. Values of atomic radii depend therefore on the binding state of the atom, and also on the method used to measure such radii.