1. Bonding Metallic Ionic Covalent The attraction between two oppositely charged entities

2. Metallic Formed between Metals Caused by The metals’ outer shell electrons being delocalised within the lattice of the resultant metal ions Actual Bond The attraction between the +ve ions and the –ve electrons Resultant structure: A lattice of +ve ions (usually a cube) in a ‘sea’ of mobile delocalised electrons

3. Properties of Metals High mp Strong attraction between the +ve ion and –ve electron. Conductivity Delocalised mobile electrons are able to move towards a +ve plate. Malleable The layers of ions are able to move over each other.

4. Structure of metals Giant Structures – HIGH MP / BP Strong electrostatic forces of attraction ( metallic bond) between every particle in the structure. How does the metallic bond get stronger? Increase the charge on the ion (stronger attraction between ion and electron, so stronger metallic bond) How Does MP/BP vary from Na  Mg ? How do we increase the conductivity of metals? Increase the number of free electrons How Does conductivity vary from Na  Mg ? Increases as Charge on ion increases from +1  +2 Increases as number of delocalised electrons per mole of lattice increases from 1  2

5. Ionic Formed between Metals and Non-metals Caused by The metals’ outer shell electrons being transferred to the outer shell of the non-metal so that both achieve a full outer shell. The metal atom becomes a +ve ion and the non-metal atom a –ve ion Actual Bond The attraction between the +ve ions and the –ve ions Resultant structure: A lattice of +ve ions (usually a cube) surrounded by –ve ions F Li F

6. 1 2 43 5 6 Represents the electrostatic attraction between opposite charged ions

7. Properties of ionic compounds High mp Strong attraction between the +ve ion and –ve ion. Conductivity in solids very poor Attraction between the +ve and –ve ions hold the ions in a fixed position in the lattice. Conductivity in aqueous solutions or liquids very high The ions become mobile so +ve ions move towards –ve plates and –ve ions move towards +ve plates.

8. Structure of Ionic compounds Giant Structures – HIGH MP / BP Strong electrostatic forces of attraction ( ionic bond) between every particle in the structure. How does the ionic bond get stronger? Increase the charge on the ions (so that the attraction between the ions increase, stronger ionic bond) How Does MP/BP vary from NaCl  Na 2 O  MgO? Increases as ions charge increase from +1/-1  +1/-2  +2/-2

9. Polarised Ionic Small Highly charged cation (Nucleus of cation is not shielded greatly) Large anion Outer shell of anion is shielded from its’ nucleus RESULT: Outer shell of anion is attracted towards the nucleus of the cation Distorts (polarises) the shape of the anion

10. Polarised ionic 9+3+ + - + - 53+

11. Formed between Non-Metals. Caused by The un-paired non-metals’ outer shell electrons are shared to form a electron pair between the two atoms. Actual Bond The attraction between the shared electron pair and the two nuclei. Resultant structure: Small discrete molecules or large giant structures Covalent

13. What is really happening to the electrons during covalent bonding? GCSE = Shells overlap A level Electrons are in (atomic) orbitals In bonding the (atomic) orbitals overlap QUESTION: WHAT DO THE ORBITALS LOOK LIKE? S = 3 p = The 3 p orbitals are all at right angles to each other

14. Molecular orbitals Found in a covalent bond The region of space in which the shared electron pair in a covalent bond is located Formed by the overlapping of the ATOMIC ORBITALS Two types IN ALL SINGLE BONDS THE ELECTRON PAIRS ARE LOCATED IN SIGMA MO IN ALL DOUBLE BONDS 1 ELECTRON PAIR IS LOCTAED IN A SIGMA MO AND THE OTHER ELECTRON PAIR IS LOCATED IN A PI MO s and s Sigma (σ) p and p end on Sigma (σ) p and p side on Pi (π) Sigma MOPi MO Sigma MO

15. Properties of covalent compounds Small discrete Molecules Low mp Weak attraction between the small molecules Poor Conductivity No charged particles free to move SIMPLE COVALENT (MOLECULAR) STRUCTURES

16. Properties of covalent compounds Giant structures High mp Strong covalent bond between every atom in structure Low Conductivity No charged particles free to move EXCEPTION Graphite – Delocalised electrons within each layer enables conduction.

17. Strange Compounds GCSE idea: After sharing electrons, atoms posses a full outer shell Dot cross diagrams for NaCl, CaO, CO 2, HCl, SO 2, SO 3, BH 3 Note the problems with last three!

18. NaCl CaO CO 2, HCl SO 2 SO 3 NOTE S has 10 electrons in outer shell NOTE S has 12 electrons in outer shell BH 3 NOTE B has 6 electrons in outer shell

19. What causes Bonding to happen? Electronegativity The ability of an atom to attract electrons towards itself in a covalent bond AFFECTED BY? a.No of shells (and hence the shielding of the nuclei) b.No of protons in the nucleus

20. Electronegativity and Periodic Table Across a Period No of shells the same Shielding the same No of Protons increases Nuclear pull on electrons increases Electronegativity increases Down a Group No of shells increases Shielding increases Nuclear pull on electrons decreases Electronegativity decreases MOST ELECTRONEGATIVE ATOM = F LEAST ELECTRONEGATIVE ATOM = Cs ELECTRONEGATIVITY OF C < ELECTRONEGATIVITY OF HALOGENS ELECTRONEGATIVITY OF C = ELECTRONEGATIVITY OF HYDROGEN

21. Consider a Covalent bond X-Y Elec X = Elec Y Elec X > Elec Y Elec X >>> Elec Y ElecX >>>>>>>>>>Elec Y Hence Bond between a metal (low elec) and a non- metal (high elec) usually IONIC Bond between 2 non-metals (similar elec) usually Covalent Polarised covalent or ionic bonds ensue if somewhere in between IONIC BOND= No overlap of shells COVALENT Bond = Overlap of shells

22. Boron Compounds Boron is in group 3 Expect ionic compounds (lose 3 electrons to achieve a full outer shell) Electronic configuration of B +3 = 1s 2 Polarisation so great: Covalent bonds ensue (even with BF 3 ) NB: B-F bond electron pair nearer to F atom once covalent

23. Aluminium compounds Al +3 Electronic configuration = 1s 2 2s 2 2p 6 Most compounds form a degree of polarised ionic (except for AlF 3 ) All aluminium tri-halides (except for AlF 3 ) are covalent molecules

24. Covalent Bonding 1.Shared electron pair(s) of electrons between 2 nuclei 2.Pair usually results from 1 unpaired electron from one atom, pairing up with an unpaired electron from another atom. Doesn’t have to be the case 3.Two types of electron pairs are seen in molecules. Square = Bonding Electron Pair Circle= Non-Bonding Electron Pair or LONE pair Types of electron pairs in a molecule

25. Dative Covalent Bonding 2 H 2 + O 2  2 H 2 O H + + OH -  H 2 O One of the bonds in the water formed in the lower example, must involve a shared electron pair in which BOTH electrons originated from the same atom. Dative Covalent (or co-ordinate) bonding is where the shared electron pair originates from the same atom. Dative Covalent bonds behave like normal covalent bonds once formed

26. Requirements for Dative Covalent Bonding 1.An atom or an atom in a molecule or ion which has a lone pair of electrons available for donation. (often –ve charged) CALLED A NUCLEOPHILE 2.An atom or an atom in a molecule or ion which has an empty orbital which can accept an electron pair. (often +ve charged) CALLED AN ELECTROPHILE

27. H + and OH - example H + ion (electronic configuration = 1s 0 ) Has an empty s orbital to accept a lone pair OH - ion Has a lone pair which it can donate into the empty orbital During the reaction, this happens to form a O-H dative covalent bond O H H Dative covalent bond

28. Boron compounds All covalent compounds All have 6 electrons in outer shell of B All have an empty orbital All can accept an electron pair NN B B

29. Shapes of Molecules Valence Shell Electron Pair Repulsion theory (VSEPR theory) The shape of a molecule is dependent on the angle between the bonds in the molecule Electron Pairs will repel each other to the maximum extent Lone pairs repel to a greater extent than bonding pairs The bond angle (and hence shape) is therefore dependent on the number of bonding and lone pairs surrounds the central atom in the molecule

30. Working out shapes Draw dot cross diagrams (or displayed formulae) Work out the number of bonding and lone pairs around the central atom Treat multiple bonds as 1 bonding pair H 2 BeI 2 (covalent) AlCl 3 (covalent) CF 4 H 2 S PH 3 SiH 4 SO 3 MoleculeNumber of electron Pairs Max Distance repelled Number of lone pairs Actual Bond angle Shape H2H2 11800 Linear BeI 2 21800 Linear AlCl 3 31200 Trigonal Planar CF 4 4109.50 Tetrahedral H2SH2S4109.52104.5Bent PH 3 4109.51107Pyramidal SiH 4 4109.50 Tetrahedral SO 3 31200 Trigonal Planar SF 6 6900 Octahedral

31. BeI 2 CO 2 AlCl 3 BH 3 CF 4 CH 4 SO 3 SF 6 NF 3 PH 3 H2SH2OH2SH2O

32. Complex molecules C 2 Cl 4 CH 3 CH 2 OH Work out the bond angle (and hence shape around each central atom)

33. Intermolecular Forces What makes a substance a solid at RT? Why is F 2 a gas, while I 2 is a solid? How can you get liquid F 2 ? Why does H 2 O boil at 100 o C while H 2 S boils at - 60.3 o C? Intermolecular forces: THE FORCE OF ATTRACTION BETWEEN MOLECULES Still a force of attraction between 2 oppositely charged entities

34. VAN DER WAALS FORCES Consider two F 2 molecules What force of attraction could exist between them to enable F 2 to liquefy?

35. VAN DER WAALS FORCES Electrons in a bond are not static Closer to one atom  Temporary Dipole Causes dipole to form in neighbouring molecules  Induced Dipoles Attraction between these two dipoles = VDW

36. Permanent Dipole- Dipole Molecules containing atoms of different electronegativity eg HCl, CH 3 Cl, Permanent dipoles are present Interactions occur between the δ+ve atom in one molecule and the δ-ve atom in the other molecule

37. Molecules without permanent dipoles when expected Symmetrical molecules O=C=O Dipoles cancel out

38. Hydrogen Bonding Especially large dipoles in molecules Especially large dipole-dipole attraction Called H-Bond Formed between molecules containing OH, NH or HF bonds

39. Drawing Hydrogen bonds Need at least 2 molecules Need to draw on dipoles on O and H atoms Need to draw lone pairs on O Draw straight line (often dotted) between lone pair on O to H on DIFFERENT Molecule

41. Water and DNA WATER High MP than expected High surface tension Low density of ice DNA Helical structure Strands held together H-Bonds need to be broken H-Bonds on surface pull molecules on surface in H-Bonds hold Water molecules apart in Solid Water

42. Boiling Points Non-Polar compounds LOW as VDW weak BP increases as no of electrons in molecules increases Temporary dipoles / induced dipoles increase Force of attraction between molecules increase Polar compounds HIGHER AS permanent dipole dipole attractions stronger than VDW Need to look for a bond in the molecule containing two atoms of different electronegativities. Compounds containing OH, NH or HF bonds Much higher as Hydrogen bonds much stronger than permanent dipole dipole

43. BUT Intermolecular forces only found in simple molecular structures and are MUCH WEAKER than the covalent, ionic or metallic bonds found in GIANT STRUCTURES BP 0 C NaCl1250 MgO2390 Cu2600 Carbon4500 CH 4 -79 CH 3 Cl-25 H 2 O100 Giant Structures: BREAK STRONG CHEMICAL BONDS TO MELT Simple Molecular Structures: BREAK Weak intermolecular forces to melt In solid Methane Particle = CH 4 In Gaseous Methane Particle = CH 4 No Chemical bond is broken on melting / boiling