1. Electrons bonding and structure

2. Module 2 aims Electron structure Chemical bonding Molecular shapes Intermolecular forces Bonding physical properties Redox

3. Electron structure

4. ORBITALS a region in space where one is likely to find an electron. An orbital is... a region in space where one is likely to find an electron. Orbitals can hold up to two electrons as long as they have opposite spin; this is known as PAULI’S EXCLUSION PRINCIPAL. Orbitals have different shapes... ORBITAL SHAPE OCCURRENCE ssphericalone in every principal level pdumb-bellthree in levels from 2 upwards dvariousfive in levels from 3 upwards fvariousseven in levels from 4 upwards An orbital is a 3-dimensional statistical shape showing where one is most likely to find an electron. Because, according to Heisenberg, you cannot say exactly where an electron is you are only able to say where it might be found. DO NOT CONFUSE AN ORBITAL WITH AN ORBIT

5. Orbitals are not filled in numerical order because the principal energy levels get closer together as you get further from the nucleus. This results in overlap of sub levels. The first example occurs when the 4s orbital is filled before the 3d orbitals. INCREASING ENERGY / DISTANCE FROM NUCLEUS 1 1s 2 2s 2p 4s 3 3s 3p 3d 4 4p 4d 4f PRINCIPAL ENERGY LEVELS SUB LEVELS 1 1s 2 2s 2p 3d 3 3s 3p 4s 4 4p 4d 4f PRINCIPAL ENERGY LEVELS SUB LEVELS ORDER OF FILLING ORBITALS THE FILLING ORDER 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 7s 7p HOW TO REMEMBER...

6. 1 1s 2 2s 2p 4s 3 3s 3p 3d 4 4p 4d 4f INCREASING ENERGY / DISTANCE FROM NUCLEUS This states that… “ELECTRONS ENTER THE LOWEST AVAILABLE ENERGY LEVEL” THE ‘AUFBAU’ PRINCIPAL The following sequence will show the ‘building up’ of the electronic structures of the first 36 elements in the periodic table. Electrons are shown as half headed arrows and can spin in one of two directions or s orbitals p orbitals d orbitals

7. 1 1s 2 2s 2p 4s 3 3s 3p 3d 4 4p 4d 4f INCREASING ENERGY / DISTANCE FROM NUCLEUS HELIUM 1s 2 THE ELECTRONIC CONFIGURATIONS Every orbital can contain 2 electrons, provided the electrons are spinning in opposite directions. This is based on... PAULI’S EXCLUSION PRINCIPLE The two electrons in a helium atom can both go in the 1s orbital. ‘Aufbau’ Principle ‘Aufbau’ Principle

8. 1 1s 2 2s 2p 4s 3 3s 3p 3d 4 4p 4d 4f INCREASING ENERGY / DISTANCE FROM NUCLEUS THE ELECTRONIC CONFIGURATIONS OXYGEN With all three orbitals half- filled, the eighth electron in an oxygen atom must now pair up with one of the electrons already there. 1s 2 2s 2 2p 4 ‘Aufbau’ Principle ‘Aufbau’ Principle

9. 1 1s 2 2s 2p 4s 3 3s 3p 3d 4 4p 4d 4f INCREASING ENERGY / DISTANCE FROM NUCLEUS THE ELECTRONIC CONFIGURATIONS FLUORINE The electrons continue to pair up with those in the half-filled orbitals. 1s 2 2s 2 2p 5

10. 1 1s 2 2s 2p 4s 3 3s 3p 3d 4 4p 4d 4f INCREASING ENERGY / DISTANCE FROM NUCLEUS THE ELECTRONIC CONFIGURATIONS NEON The electrons continue to pair up with those in the half-filled orbitals. The 2p orbitals are now completely filled and so is the second principal energy level. In the older system of describing electronic configurations, this would have been written as 2,8. 1s 2 2s 2 2p 6

11. 1s 1 1s 2 1s 2 2s 1 1s 2 2s 2 1s 2 2s 2 2p 1 1s 2 2s 2 2p 2 1s 2 2s 2 2p 3 1s 2 2s 2 2p 4 1s 2 2s 2 2p 5 1s 2 2s 2 2p 6 1s 2 2s 2 2p 6 3s 1 1s 2 2s 2 2p 6 3s 2 1s 2 2s 2 2p 6 3s 2 3p 1 1s 2 2s 2 2p 6 3s 2 3p 2 1s 2 2s 2 2p 6 3s 2 3p 3 1s 2 2s 2 2p 6 3s 2 3p 4 1s 2 2s 2 2p 6 3s 2 3p 5 1s 2 2s 2 2p 6 3s 2 3p 6 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 1 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 2 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 3 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 5 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 5 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 6 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 7 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 8 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 10 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 H He Li Be B C N O F Ne Na Mg Al Si P S Cl Ar K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn ELECTRONIC CONFIGURATIONS OF ELEMENTS 1-30

12. ELECTRONIC CONFIGURATION OF IONS Positive ions (cations) are formed by removing electrons from atoms Negative ions (anions) are formed by adding electrons to atoms Electrons are removed first from the highest occupied orbitals (EXC. transition metals) SODIUMNa1s 2 2s 2 2p 6 3s 1 1 electron removed from the 3s orbital Na + 1s 2 2s 2 2p 6 CHLORINECl1s 2 2s 2 2p 6 3s 2 3p 5 1 electron added to the 3p orbital Cl¯1s 2 2s 2 2p 6 3s 2 3p 6 FIRST ROW TRANSITION METALS Despite being of lower energy and being filled first, electrons in the 4s orbital are removed before any electrons in the 3d orbitals. TITANIUMTi1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 2 Ti + 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 2 Ti 2+ 1s 2 2s 2 2p 6 3s 2 3p 6 3d 2 Ti 3+ 1s 2 2s 2 2p 6 3s 2 3p 6 3d 1 Ti 4+ 1s 2 2s 2 2p 6 3s 2 3p 6

13. Electronic structure summary Electrons occupy energy levels around the nucleus of the atom, where each shell has a principal quantum number. For principal quantum number, n = 1, the number of electrons is 2; for n = 2, the number is 8; then 18; then 32 electrons for n = 4. Main energy levels are sub-divided into sub-shells and these consist of orbitals called s, p and d-orbitals. Elements have an electronic configuration that can be shown in s, p or d notation, for example, sodium is 1s 2, 2s 2, 2p 6, 3s 1.

14. BONDING The physical properties of a substance depend on its structure and type of bonding present. Bonding determines the type of structure. TYPES OF BOND CHEMICAL ionic (or electrovalent) strong bonds covalent dative covalent (or co-ordinate) metallic PHYSICAL van der Waals‘ forces- weakest weak bonds dipole-dipole interaction hydrogen bonds - strongest

15. THE IONIC BOND Ionic bonds tend to be formed between elements whose atoms need to “lose” electrons to gain the nearest noble gas electronic configuration (n.g.e.c.) and those which need to gain electrons. The electrons are transferred from one atom to the other. Sodium Chloride Na ——> Na+ + e¯ and Cl + e¯ ——> Cl¯ 1s 2 2s 2 2p 6 3s 1 1s 2 2s 2 2p 6 1s 2 2s 2 2p 6 3s 2 3p 5 1s 2 2s 2 2p 6 3s 2 3p 6 or 2,8,1 2,8 2,8,7 2,8,8 An electron is transferred from the 3s orbital of sodium to the 3p orbital of chlorine; both species end up with the electronic configuration of the nearest noble gas the resulting ions are held together in a crystal lattice by electrostatic attraction.

16. Physical properties of ionic compounds Melting point very highA large amount of energy must be put in to overcome the strong electrostatic attractions and separate the ions. Strength Very brittleAny dislocation leads to the layers moving and similar ions being adjacent. The repulsion splits the crystal. Electricaldon’t conduct when solid - ions held strongly in the lattice conduct when molten or in aqueous solution - the ions become mobile and conduction takes place. SolubilityInsoluble in non-polar solvents but soluble in water Water is a polar solvent and stabilises the separated ions. Much energy is needed to overcome the electrostatic attraction and separate the ions stability attained by being surrounded by polar water molecules compensates for this

17. IONIC BONDING BRITTLE IONIC LATTICES ++ + + ++ ++ - - - - -- - - + + ++ IF YOU MOVE A LAYER OF IONS, YOU GET IONS OF THE SAME CHARGE NEXT TO EACH OTHER. THE LAYERS REPEL EACH OTHER AND THE CRYSTAL BREAKS UP.

18. IONIC COMPOUNDS - ELECTRICAL PROPERTIES SOLID IONIC COMPOUNDS DO NOT CONDUCT ELECTRICITY Na + Cl - Na + Cl - Na + Cl - Na + Cl - Na + Cl - Na + Cl - IONS ARE HELD STRONGLY TOGETHER + IONS CAN’T MOVE TO THE CATHODE - IONS CAN’T MOVE TO THE ANODE MOLTEN IONIC COMPOUNDS DO CONDUCT ELECTRICITY Na + Cl - Na + Cl - Na + Cl - Na + Cl - IONS HAVE MORE FREEDOM IN A LIQUID SO CAN MOVE TO THE ELECTRODES SOLUTIONS OF IONIC COMPOUNDS IN WATER DO CONDUCT ELECTRICITY DISSOLVING AN IONIC COMPOUND IN WATER BREAKS UP THE STRUCTURE SO IONS ARE FREE TO MOVE TO THE ELECTRODES

19. Definitionconsists of a shared pair of electrons with one electron being supplied by each atom either side of the bond. compare this with dative covalent bonding atoms are held together because their nuclei which have an overall positive charge are attracted to the shared electrons Formationbetween atoms of the same elementN 2, O 2, diamond, graphite between atoms of different elementsCO 2, SO 2 on the RHS of the table; when one of the elements is in theCCl 4, SiCl 4 middle of the table; with head-of-the-group elements BeCl 2 with high ionisation energies; COVALENT BONDING + +

20. atoms share electrons to get the nearest noble gas electronic configuration some don’t achieve an “octet” as they haven’t got enough electrons eg Al in AlCl 3 others share only some - if they share all they will exceed their “octet” egNH 3 and H 2 O atoms of elements in the 3rd period onwards can exceed their “octet” if they wish as they are not restricted to eight electrons in their “outer shell” egPCl 5 and SF 6 COVALENT BONDING

21. Orbital theory Covalent bonds are formed when orbitals, each containing one electron, overlap. This forms a region in space where an electron pair can be found; new molecular orbitals are formed. SIMPLE MOLECULES The greater the overlap the stronger the bond. orbital containing 1 electron overlap of orbitals provides a region in space which can contain a pair of electrons

22. HYDROGEN H Another hydrogen atom also needs one electron to complete its outer shell Hydrogen atom needs one electron to complete its outer shell atoms share a pair of electrons to form a single covalent bond A hydrogen MOLECULE is formed H H H WAYS TO REPRESENT THE MOLECULE PRESSING THE SPACE BAR WILL ACTIVATE EACH STEP OF THE ANIMATION

23. HYDROGEN CHLORIDE Cl H Hydrogen atom also needs one electron to complete its outer shell Chlorine atom needs one electron to complete its outer shell atoms share a pair of electrons to form a single covalent bond H Cl WAYS TO REPRESENT THE MOLECULE PRESSING THE SPACE BAR WILL ACTIVATE EACH STEP OF THE ANIMATION

24. METHANE C Each hydrogen atom needs 1 electron to complete its outer shell A carbon atom needs 4 electrons to complete its outer shell Carbon shares all 4 of its electrons to form 4 single covalent bonds H H H H H C H H H H H WAYS TO REPRESENT THE MOLECULE PRESSING THE SPACE BAR WILL ACTIVATE EACH STEP OF THE ANIMATION

25. WATER O Each hydrogen atom needs one electron to complete its outer shell Oxygen atom needs 2 electrons to complete its outer shell Oxygen can only share 2 of its 6 electrons otherwise it will exceed the maximum of 8 2 LONE PAIRS REMAIN H H H O H H WAYS TO REPRESENT THE MOLECULE PRESSING THE SPACE BAR WILL ACTIVATE EACH STEP OF THE ANIMATION

26. AMMONIA N H H H N H HH H N H H H each atom needs one electron to complete its outer shell atom needs three electrons to complete its outer shell Nitrogen can only share 3 of its 5 electrons otherwise it will exceed the maximum of 8 A LONE PAIR REMAINS

27. WATER O H H O H H each atom needs one electron to complete its outer shell atom needs two electrons to complete its outer shell Oxygen can only share 2 of its 6 electrons otherwise it will exceed the maximum of 8 TWO LONE PAIRS REMAIN H O H H

28. OXYGEN O each atom needs two electrons to complete its outer shell each oxygen shares 2 of its electrons to form a DOUBLE COVALENT BOND OO O O

29. BondingAtoms are joined together within the molecule by covalent bonds. ElectricalDon’t conduct electricity as they have no mobile ions or electrons Solubility Tend to be more soluble in organic solvents than in water; some are hydrolysed Boiling pointLow - intermolecular forces (van der Waals’ forces) are weak; they increase as molecules get a larger surface area e.g. CH 4 -161°CC 2 H 6 - 88°C C 3 H 8 -42°C as the intermolecular forces are weak, little energy is required to to separate molecules from each other so boiling points are low some boiling points are higher than expected for a given mass because you can get additional forces of attraction SIMPLE COVALENT MOLECULES

30. A dative covalent bond differs from covalent bond only in its formation Both electrons of the shared pair are provided by one species (donor) and it shares the electrons with the acceptor Donor species will have lone pairs in their outer shells Acceptor species will be short of their “octet” or maximum. Lewis basea lone pair donor Lewis acida lone pair acceptor DATIVE COVALENT (CO-ORDINATE) BONDING Ammonium ion, NH 4 + The lone pair on N is used to share with the hydrogen ion which needs two electrons to fill its outer shell. The N now has a +ive charge as - it is now sharing rather than owning two electrons.

31. Boron trifluoride-ammonia NH 3 BF 3 Boron has an incomplete shell in BF 3 and can accept a share of a pair of electrons donated by ammonia. The B becomes -ive as it is now shares a pair of electrons (i.e. it is up one electron) it didn’t have before.

32. METALLIC BONDING Involves a lattice of positive ions surrounded by delocalised electrons Metal atoms achieve stability by “off-loading” electrons to attain the electronic structure of the nearest noble gas. These electrons join up to form a mobile cloud which prevents the newly-formed positive ions from flying apart due to repulsion between similar charges. Atoms arrange in regular close packed 3-dimensional crystal lattices. The outer shell electrons of each atom leave to join a mobile “cloud” or “sea” of electrons which can roam throughout the metal. The electron cloud binds the newly-formed positive ions together.

33. METALLIC BOND STRENGTH Depends on the number of outer electrons donated to the cloud and the size of the metal atom/ion. The strength of the metallic bonding in sodium is relatively weak because each atom donates one electron to the cloud. The metallic bonding in potassium is weaker than in sodium because the resulting ion is larger and the electron cloud has a bigger volume to cover so is less effective at holding the ions together. The metallic bonding in magnesium is stronger than in sodium because each atom has donated two electrons to the cloud. The greater the electron density holds the ions together more strongly. Na Mg K

34. METALLIC PROPERTIES MOBILE ELECTRON CLOUD ALLOWS THE CONDUCTION OF ELECTRICITY For a substance to conduct electricity it must have mobile ions or electrons. Because the ELECTRON CLOUD IS MOBILE, electrons are free to move throughout its structure. Electrons attracted to the positive end are replaced by those entering from the negative end. Metals are excellent conductors of electricity

35. MALLEABLE CAN BE HAMMERED INTO SHEETS DUCTILE CAN BE DRAWN INTO RODS AND WIRES As the metal is beaten into another shape the delocalised electron cloud continues to bind the “ions” together. Some metals, such as gold, can be hammered into sheets thin enough to be translucent. METALLIC PROPERTIES Metals can have their shapes changed relatively easily

36. HIGH MELTING POINTS Melting point is a measure of how easy it is to separate individual particles. In metals it is a measure of how strong the electron cloud holds the + ions. The ease of separation of ions depends on the... ELECTRON DENSITY OF THE CLOUD IONIC / ATOMIC SIZE PERIODSNa (2,8,1) < Mg (2,8,2) < Al (2,8,3) m.pt 98°C 650°C 659°C b.pt 890°C 1110°C 2470°C METALLIC PROPERTIES Na + Al 3+ Mg 2+ MELTING POINT INCREASES ACROSS THE PERIOD THE ELECTRON CLOUD DENSITY INCREASES DUE TO THE GREATER NUMBER OF ELECTRONS DONATED PER ATOM. AS A RESULT THE IONS ARE HELD MORE STRONGLY.

37. HIGH MELTING POINTS Melting point is a measure of how easy it is to separate individual particles. In metals it is a measure of how strong the electron cloud holds the + ions. The ease of separation of ions depends on the... ELECTRON DENSITY OF THE CLOUD IONIC / ATOMIC SIZE GROUPS Li (2,1) < Na (2,8,1) < K (2,8,8,1) m.pt 181°C 98°C 63°C b.pt 1313°C 890°C 774°C METALLIC PROPERTIES MELTING POINT INCREASES DOWN A GROUP IONIC RADIUS INCREASES DOWN THE GROUP. AS THE IONS GET BIGGER THE ELECTRON CLOUD BECOMES LESS EFFECTIVE HOLDING THEM TOGETHER SO THEY ARE EASIER TO SEPARATE. Na + K+K+ Li +

38. Chemical bonding summary Ionic bonding takes place when positive ions and negative ions are attracted in a giant ionic structure. Covalent bonding is the sharing of electron pair(s) between nuclei of atoms. The covalent bond and ionic bond are both very strong chemical bonds. A dative covalent bond is one formed in which both electrons are donated from the same atom.

39. REGULAR SHAPES Molecules, or ions, possessing ONLY BOND PAIRS of electrons fit into a set of standard shapes. All the bond pair-bond pair repulsions are equal. All you need to do is to count up the number of bond pairs and chose one of the following examples... C 2LINEAR 180ºBeCl 2 3TRIGONAL PLANAR 120ºAlCl 3 4TETRAHEDRAL 109.5ºCH 4 5TRIGONAL BIPYRAMIDAL 90º & 120º PCl 5 6OCTAHEDRAL 90ºSF 6 BOND PAIRS SHAPE ANGLE(S)EXAMPLE A covalent bond will repel another covalent bond

41. IRREGULAR SHAPES If a molecule, or ion, has lone pairs on the central atom, the shapes are slightly distorted away from the regular shapes. This is because of the extra repulsion caused by the lone pairs. BOND PAIR - BOND PAIR < LONE PAIR - BOND PAIR < LONE PAIR - LONE PAIR O O O As a result of the extra repulsion, bond angles tend to be slightly less as the bonds are squeezed together.

42. AMMONIA ANGLE... 107° SHAPE... PYRAMIDAL H N N H H H BOND PAIRS3 LONE PAIRS1 TOTAL PAIRS4 H H N H H H N H 107° H H N H The shape is based on a tetrahedron but not all the repulsions are the same LP-BP REPULSIONS > BP-BP REPULSIONS The N-H bonds are pushed closer together Lone pairs are not included in the shape

43. SHAPES OF IONS N BOND PAIRS 3PYRAMIDAL LONE PAIRS 1H-N-H 107° BOND PAIRS 4TETRAHEDRAL LONE PAIRS 0 H-N-H 109.5° N H H H N+N+ H H H H N+N+ BOND PAIRS 2ANGULAR LONE PAIRS 2H-N-H 104.5° N H H N NH 4 + NH 2 - NH 3 REVIEWREVIEW

44. MOLECULES WITH DOUBLE BONDS C O C OO Carbon - needs four electrons to complete its shell Oxygen - needs two electron to complete its shell The atoms share two electrons each to form two double bonds DOUBLE BOND PAIRS2 LONE PAIRS0 BOND ANGLE... SHAPE... 180° LINEAR OOC 180° Double bonds behave exactly as single bonds for repulsion purposes so the shape will be the same as a molecule with two single bonds and no lone pairs. The shape of a compound with a double bond is calculated in the same way. A double bond repels other bonds as if it was single e.g. carbon dioxide

45. OTHER EXAMPLES BrF 5 BOND PAIRS 5 LONE PAIRS 1 ‘UMBRELLA’ ANGLES 90° <90° F F F F Br F F F F F F BrF 3 BOND PAIRS 3 LONE PAIRS 2 ’T’ SHAPED ANGLE <90° F F Br F F F F SO 4 2- O S O-O- O-O- O O S O-O- O-O- O BOND PAIRS 4 LONE PAIRS 0 TETRAHEDRAL ANGLE 109.5°

46. ANSWERS ON NEXT PAGE TEST QUESTIONS For each of the following ions/molecules, state the number of bond pairs state the number of lone pairs state the bond angle(s) state, or draw, the shape SiCl 4 PCl 6 - H2SH2S SiCl 6 2- PCl 4 + BF 3

47. TEST QUESTIONS 3 bp0 lp120ºtrigonal planarboron pairs up all 3 electrons in its outer shell 4 bp0 lp109.5ºtetrahedralsilicon pairs up all 4 electrons in its outer shell 4 bp0 lp109.5ºtetrahedralas ion is +, remove an electron in the outer shell then pair up 6 bp0 lp90ºoctahedralas the ion is -, add one electron to the 5 in the outer shell then pair up 6 bp0 lp90ºoctahedralas the ion is 2-, add two electrons to the outer shell then pair up 2 bp2 lp92ºangularsulphur pairs up 2 of its 6 electrons in its outer shell - 2 lone pairs are left BF 3 SiCl 4 PCl 6 - H2SH2S SiCl 6 2- PCl 4 + ANSWER For each of the following ions/molecules, state the number of bond pairs state the number of lone pairs state the bond angle(s) state, or draw, the shape

48. Shapes Number of bonded pairs of electrons Number of lone pairs of electrons ShapeBond angle 20Linear180 30120 40Tetrahedral 5090 & 120 60Octahedral 31107 22Angular

49. Shapes Number of bonded pairs of electrons Number of lone pairs of electrons ShapeBond angle 30Linear180 30Trigonal planar120 40Tetrahedral109.5 50Trigonal- bipyramidal 90 & 120 60Octahedral90 31Pyramidal107 22Angular104

50. Molecular shapes summary The shape of a molecule is determined by the repulsion between bonded electrons and non-bonded electrons (lone pairs). Lone electron pairs repel more than bonded pairs of electrons and give rise to distorted shapes. By deducing the number of bonded electron pairs and lone pairs of electrons, the shape of a molecule may be predicted. BF 3 is trigonal planar; CH 4 and NH 4 + are tetrahedral; SF 6 is octahedral; H 2 O is non-linear (V-shaped/bent); CO 2 is linear and ammonia, NH 3, as pyramidal

51. ‘The ability of an atom to attract the electron pair in a covalent bond to itself’ Pauling Scalea scale for measuring electronegativity values increase across periods values decrease down groups fluorine has the highest value H 2.1 Li Be B C N O F 1.0 1.5 2.0 2.5 3.0 3.5 4.0 Na Mg Al Si P S Cl 0.9 1.2 1.5 1.8 2.1 2.5 3.0 K Br 0.8 2.8 ELECTRONEGATIVITY INCREASE

52. Electronegativity No electronegativity difference between two atoms leads to a pure non-polar covalent bond. A small electronegativity difference leads to a polar covalent bond. A large electronegativity difference leads to an ionic bond.

53. Spectrum of bonding 100% covalent100% ionic 0.03.2Difference in electronegativity As a rough guide: A difference greater than 2.0 will lead to essentially ionic bonding. A difference that is less than 1.0 will lead to bonding that is mainly covalent

54. Occurrencenot all molecules containing polar bonds are polar overall if bond dipoles ‘cancel each other’ the molecule isn’t polar if there is a ‘net dipole’ the molecule will be polar HYDROGEN CHLORIDE TETRACHLOROMETHANE WATER POLAR MOLECULES NET DIPOLE - POLAR NON-POLAR NET DIPOLE - POLAR

55. Factors affecting polarising ability Charge density of the positive ion Increases as the positive ion gets smaller and the charge density gets larger. As a negative ion gets bigger it becomes easier to polarise.

56. Type of dipole Permanent Dipole – dipole forces InstantaneousInduced Van der Waal forces

57. Although the bonding within molecules is strong, between molecules it is weak. Molecules and monatomic gases are subject to weak attractive forces. Instantaneous dipole-induced dipole forces Electrons move quickly in orbitals, so their position is constantly changing; at any given time they could be Anywhere in an atom. The possibility exists that one side has More electrons than the other. This will give rise to a dipole... The dipole on one atom induces dipoles on others Atoms are now attracted to each other by a weak forces The greater the number of electrons, the stronger the attraction and the greater the energy needed to separate the particles. NOBLE GASESALKANES Electrons B pt. Electrons B pt. He 2 -269°C CH 4 10 -161°C Ne10 -246°C C 2 H 6 18 - 88°C Ar18 -186°C C 3 H 8 26 - 42°C Kr36 -152°C VAN DER WAALS’ FORCES INSTANTANEOUS DIPOLE-INDUCED DIPOLE FORCES

58. Occurrenceoccurs between molecules containing polar bonds acts in addition to the basic van der Waals’ forces the extra attraction between dipoles means that more energy must be put in to separate molecules get higher boiling points than expected for a given mass DIPOLE-DIPOLE INTERACTION Mr °C CH 4 16-161 SiH 4 32-117 GeH 4 77-90 SnH 4 123-50 NH 3 17-33 PH 3 34-90 AsH 3 78-55 SbH 3 125-17 Mr°C H 2 O18+100 H 2 S34-61 H 2 Se81-40 H 2 Te130-2 HF20+20 HCl36.5-85 HBr81-69 HI128-35 Boiling points of hydrides

59. an extension of dipole-dipole interaction gives rise to even higher boiling points bonds between H and the three most electronegative elements, F, O and N are extremely polar because of the small sizes of H, F, N and O the partial charges are concentrated in a small volume thus leading to a high charge density makes the intermolecular attractions greater and leads to even higher boiling points HYDROGEN BONDING Hydrogen bonds are formed by electronegative elements with a non-bonding pair of electrons available to bond with a H attached to an electronegative element.

60. HYDROGEN BONDING - ICE each water molecule is hydrogen-bonded to 4 others in a tetrahedral formation ice has a “diamond-like” structure volume is larger than the liquid making it when ice melts, the structure collapses slightly and the molecules come closer; they then move a little further apart as they get more energy as they warm up this is why… a)water has a maximum density at 4°C b)ice floats. hydrogen bonding lone pair

61. HYDROGEN BONDING - HF Hydrogen fluoride has a much higher boiling point than one would expect for a molecule with a relative molecular mass of 20 Fluorine has the highest electronegativity of all and is a small atom so the bonding with hydrogen is extremely polar F H F H H F H F  +  ¯  +  ¯  +  ¯  +  ¯ hydrogen bonding

62. Intermolecular forces An intermolecular force exists between molecules and may include hydrogen bonding, dipole-dipole or van der Waals’ forces. Electronegativity is the ability of an atom in a covalent bond to attract a bonded pair of electrons towards itself. Hydrogen bonding arises in molecules in which a hydrogen atom is bonded to either an N or O atom. Water molecules, and other substances consisting of hydrogen bonding, have anomalous properties as a result.

63. Bonding and physical properties Metals consist of a close-packed arrangement of positive ions, through which delocalised electrons move. Metals are very good electrical conductors as a result of having mobile electrons. Giant structures have high melting and boiling points due to strong chemical bonds acting throughout the structure. Giant ionic structures conduct electricity when molten, and when dissolved in water due to mobile ions, not electrons.

65. Working out oxidation states Enter zero for any uncombined element (e.g. Na or Cl 2 ) Put in the oxidation state(s) of the elements you are sure of. (see above table) Multiply this figure by the number of atoms of the element that are present Take this number change the sign and assign it to the element unknown oxidation state. Divide the number by the number of atoms of that element in the compound This final figure is the oxidation state of the element.

66. OXIDATION STATES Q. Name the following...PbO 2 lead(IV) oxide SnCl 2 tin(II) chloride SbCl 3 antimony(III) chloride TiCl 4 titanium(IV) chloride BrF 5 bromine(V) fluoride manganese(IV) oxide shows that Mn is in the +4 oxidation state in MnO 2 sulphur(VI) oxide for SO 3 S is in the +6 oxidation state dichromate(VI) for Cr 2 O 7 2- Cr is in the +6 oxidation state phosphorus(V) chloride for PCl 5 P is in the +5 oxidation state phosphorus(III) chloride for PCl 3 P is in the +3 oxidation state THE ROLE OF OXIDATION STATE IN NAMING SPECIES To avoid ambiguity, the oxidation state is often included in the name of a species

67. Redox An oxidation number indicates the formal charge of a chemically combined particle in a compound. The oxidation number of metals usually equals the group number (as a positive value) and minus (8 – group number) for non-metals. An element has been oxidised if the oxidation number increases, and reduced if the oxidation number decreases. When they react, metals are normally oxidised (they lose electrons), whereas non-metals gain electrons and are reduced.