2. Bonding & Structure K Warne

3. Bonding & Structure 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. Bonding O H H H2OH2O O H H H2OH2O O O H H H H Bonding takes place when atoms react to form molecules. Example: Two hydrogen molecules and one oxygen molecule react to form two water molecules.

5. Atomic Radius Note the trends in atomic radius across the periodic table.

6. VALENCY 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 SClAr

7. VALENCY Valencyelectrons lost gained Valency – number of electrons lost or gained to have a FULL valence level. (Outer shell) H He Li Be B C N OFNe Valence electrons Valence electrons – those in outer shell. METALS NON - METALS

8. Covalent bond A shared PAIR of electrons.  Formed between _____________________________.  Pure covalent bonds have ____________________ SHARING of the electrons.  _________________ Molecules; H-H, (H 2 ), O 2, F 2, Cl 2, Br 2, I 2, N 2,  Common Ions also covalently bonded KNOW formulae; eg sulphate ion SO 4 2- H x H HH In covalent substances the electrons are strongly held in the bonds and so the substance will NOT conduct electricity.

9. Covalent bond A shared PAIR of electrons.  Electrons from one atom are attracted strongly by the nucleus of another atom.  Formed between non metals. (Attract electrons strongly!)  Pure covalent bonds have EQUAL SHARING of the electrons.  In diatomic Molecules; H 2, O 2, F 2, Cl 2, Br 2, I 2, N 2, HH In covalent substances the electrons are strongly held in the bonds and so the substance will NOT conduct electricity. H H x

10. Covalent bond – Bohr Diagrams (G9) x x x x O O O x x x x x x x O x O O H H H H x x x x

11. 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 O “Dot Cross Diagrams” - Lewis & Couper Notation Lewis Diagrams Couper Notation Chemical Formulae ………… Name:Oxygen

12. 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 O=O “Dot Cross Diagrams” - Lewis & Couper Notation Lewis Diagrams Couper Notation Chemical Formulae O 2 Name:Oxygen

13. Covalent Molecules H 2, O 2, F 2, Cl 2, Br 2, N 2, H 2 O, NH 3, CH 4 CO 2, NH 4 +, CL

14. Covalent Molecules H 2, O 2, F 2, Cl 2, Br 2, N 2, H 2 O, NH 3, CH 4 CO 2, NH 4 +,

15. Covalent Molecules H 2, O 2, F 2, Cl 2, Br 2, N 2, H 2 O, NH 3, CH 4 CO 2, NH 4 +,

16. 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

17. Formed when there is a transfer of electrons. Formed between metals and non metals. Metals lose electrons and become positively charged ions - CATIONS. Non metals gain electrons and become negatively charged ions – ANIONS – called CHLORIDE. (“ide” = negative ion) Electrostatic attraction between oppositely charged ions bonds the ions together. Na... :Cl: -.. Na +. :Cl:.. Na. + : Cl: --> [Na] + [Cl] -. ELECTROSTATICATTRACTION Ionic Bonding Sodium atom Chlorine atom Sodium ion Chloride ion

18. Ionic Bonding – Bohr Diagrams The final compound is ALWAYS NEUTRAL The total charges of the cations and anions must balance out. Cl - Na + +1 Chlorine atom Chloride ion Sodium atom Sodium ion

19. 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.

20. 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 -

21. Formation of Ionic Bond 1. Write down all the steps that need to take place for this change to take place. 2. Try and place the steps in order. 3. Decide which steps would be endothermic and which would be exothermic. Na (s) + Cl 2(g) NaCl (s)

22. 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)

23. 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

24. MUST BE LEARNT BY HEART! ONETWOTHREE Hydrogen H + Beryllium Be 2+ Aluminium Al 3+ Lithium Li + Magnesium Mg 2+ Iron(III) Fe 3+ Sodium Na + Calcium Ca 2+ Potassium K + Barium Ba 2+ Silver Ag + Lead Pb 2+ Copper(I) Cu + Zinc Zn 2+ Ammonium NH 4 + Iron(II) Fe 2+ Oxonium H 3 O + Copper(II) Cu 2+ VALENCY TABLE 1

25. VALENCY TABLE 2 Negative Ions Fluoride F - Oxide O 2- Nitride N 3- Chloride Cl - Sulphide S 2- Phosphate PO 4 3- Bromide Br - Carbonate CO 3 2- Iodide I - Sulphate SO 4 2- Hydroxide OH - Nitrate NO 3 - Hydrogencarbonate HCO 3 - Hydrogensulphate HSO 4 - Permanganate MnO 4 - Ethanoate CH 3 COO - The trivial names for HCO 3 - and HSO 4 - are bicarbonate and bisulphate, respectively.

26. Bonding - Metallic Bonding - Exists between _________________. - Metal electrons are _____________ - therefore they become ______________________ (move from one atom to another). - This leaves _______________ - which become surrounded by a ‘sea’ of ______________________ electrons. - A force of _______________________ exists between the delocalized ___________________ and the positive ___________- which forms the ___________________ bond. All the _____________ of metals can be explained in terms of this bonding.

27. Bonding - Metallic Bonding - Exists between metal atoms. - Metal electrons are weakly held - therefore they become delocalized (move from one atom to another). - This leaves 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.

28. Explaining Metal Properties PropertyExplanation Malleable Ductile Conductors of electricity Shiny (Luster)

29. Explaining Metal Properties PropertyExplanation Malleable Weakly held electrons are able to move from one atom to another (delocalized). If subjected to significant force the atoms/ions are able to change positions and change the shape of the metal. Ductile Weakly held electrons are able to move from one atom to another (delocalized). If subjected to significant stretching force the atoms/ions are able to change positions and be drawn into wires. Conductors of electricity Weakly held electrons are able to move from one atom to another (delocalized). If subjected to a potential difference the electrons can drift across the metal and conduct an electric current Shiny (Luster) Weakly held electrons are able to move from one atom to another (delocalized). Rough spots on surface can be smoothed over so surface can be polished – also electrons can absorb and radiate light.

30. Bonding Summary Covalent Ionic Metallic H x H Cl - Na + Properties Properties Properties. H-H Eg

31. 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 )

32. Network Solids Strong covalent bonds Network solids have strong bonds between all atoms. The structure is extended in three dimensions. Extensive bonding and a GIANT STRUCTURE ensure the substance has high boiling and melting points are insoluble

33. Strong covalent bonds Diamond Graphite Delocalised electrons - Weaker van der waals forces between the layers Network Solids Diamond Graphite Properties: Strong/ hard/brittle – high m. & bpts. Electrical insulator (electrons held in bonds) in all phases – no ions. insoluble

34. Phases of Matter PLASMA + + + + + + There are FOUR states or phases of matter. 12341234

35. Molecular solids Ice Iodine molecules Strong covalent bonds Weaker intermolecular bonds Water molecule

36. Metals (Delocalised electrons) Strong electromagnetic attraction between ions and ‘sea’ of delocalised electrons. GIANT structure – extensive strong bonding in 3 dimensions and extended structure High melting points & boiling points

37. Ionic Solids – Giant Network Ionic bonds An extended lattice (Regular 3D arrangement) of positive and negative ions. Strong bonding throughout ensures the structure has VERY high melting points.. Na + Cl - Properties: Hard/strong/brittle – high melting & boiling points Electrical insulator (solid – ions can’t move) – conductor in solution or liquid phase – ions free to move.

41. 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 In metal the delocalised electrons are free to carry an electric current.

42. Dissolution (dissolving) of an Ionic Solid

43. Types>>NetworkMolecularIonicMetallic Particles Bonds Structure Properties Examples

44. Types>>NetworkMolecularIonicMetallic Particlesatomsmolecules Ions + and - + Ions and Electrons Bonds Covalent bonds Inter molecular bonds (weak) Ionic bonds Metallic bonds Structure (Simple or giant) GIANT MOLECULE Simple MOLECULE GIANT LATTICE PropertiesHard/Strong VERY HIGH melting points Non conducting (usually) Low mp & Bp Non conductors Weak/Brittle Soluble in Non – polar solvents HIGH mp & bp Hard strong crystals /brittle Soluble in polar solvents Conduct in solution or in liquid form Malleable, ductile, lustre (shiny), good conductors Examples C – diamond GraphiteWater/IceIodineSulphurchlorine Salts, sodium chloride, zinc chloride… Copper iron etc.

45. Microscopy How do we know about these structures??