1 / 67

Electrons bonding and structure

Electrons bonding and structure. Module 2 aims. Electron structure Chemical bonding Molecular shapes Intermolecular forces Bonding physical properties Redox. Electron structure. ORBITALS. An orbital is... a region in space where one is likely to find an electron.

stasia
Télécharger la présentation

Electrons bonding and structure

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  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 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 s spherical one in every principal level p dumb-bell three in levels from 2 upwards d various five in levels from 3 upwards f various seven 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. PRINCIPAL ENERGY LEVELS SUB LEVELS PRINCIPAL ENERGY LEVELS SUB LEVELS 4f 4f HOW TO REMEMBER ... 4d 4d 4 4 4p 4p 3d 4s THE FILLING ORDER 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 7s 7p 4s 3p 3 3d 3s 3p 3 3s 2p INCREASING ENERGY / DISTANCE FROM NUCLEUS 2 2p 2s 2 2s 1 1 1s 1s ORDER OF FILLING ORBITALS 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.

  6. 4f 4d 4 3d 4p 3 3p 3s 2p 2 2s 1 1s THE ‘AUFBAU’ PRINCIPAL This states that… “ELECTRONS ENTER THE LOWEST AVAILABLE ENERGY LEVEL” 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 4s INCREASING ENERGY / DISTANCE FROM NUCLEUS

  7. 4f 4d 4 3d 4p 3 3p 3s 2p 2 2s 1 1s THE ELECTRONIC CONFIGURATIONS HELIUM 1s2 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. 4s INCREASING ENERGY / DISTANCE FROM NUCLEUS ‘Aufbau’ Principle

  8. 4f 4d 4 3d 4p 3 3p 3s 2p 2 2s 1 1s THE ELECTRONIC CONFIGURATIONS OXYGEN 1s2 2s2 2p4 With all three orbitals half-filled, the eighth electron in an oxygen atom must now pair up with one of the electrons already there. 4s INCREASING ENERGY / DISTANCE FROM NUCLEUS ‘Aufbau’ Principle

  9. 4f 4d 4 3d 4p 3 3p 3s 2p 2 2s 1 1s THE ELECTRONIC CONFIGURATIONS FLUORINE 1s2 2s2 2p5 The electrons continue to pair up with those in the half-filled orbitals. 4s INCREASING ENERGY / DISTANCE FROM NUCLEUS

  10. 4f 4d 4 3d 4p 3 3p 3s 2p 2 2s 1 1s THE ELECTRONIC CONFIGURATIONS NEON 1s2 2s2 2p6 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. 4s INCREASING ENERGY / DISTANCE FROM NUCLEUS

  11. 1s1 1s2 1s2 2s1 1s2 2s2 1s2 2s2 2p1 1s2 2s2 2p2 1s2 2s2 2p3 1s2 2s2 2p4 1s2 2s2 2p5 1s2 2s2 2p6 1s2 2s2 2p6 3s1 1s2 2s2 2p6 3s2 1s2 2s2 2p6 3s2 3p1 1s2 2s2 2p6 3s2 3p2 1s2 2s2 2p6 3s2 3p3 1s2 2s2 2p6 3s2 3p4 1s2 2s2 2p6 3s2 3p5 1s2 2s2 2p6 3s2 3p6 1s2 2s2 2p6 3s2 3p6 4s1 1s2 2s2 2p6 3s2 3p6 4s2 1s2 2s2 2p6 3s2 3p6 4s2 3d1 1s2 2s2 2p6 3s2 3p6 4s2 3d2 1s2 2s2 2p6 3s2 3p6 4s2 3d3 1s2 2s2 2p6 3s2 3p6 4s1 3d5 1s2 2s2 2p6 3s2 3p6 4s2 3d5 1s2 2s2 2p6 3s2 3p6 4s2 3d6 1s2 2s2 2p6 3s2 3p6 4s2 3d7 1s2 2s2 2p6 3s2 3p6 4s2 3d8 1s2 2s2 2p6 3s2 3p6 4s1 3d10 1s2 2s2 2p6 3s2 3p6 4s2 3d10 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) SODIUM Na 1s2 2s2 2p6 3s1 1 electron removed from the 3s orbital Na+ 1s2 2s2 2p6 CHLORINE Cl 1s2 2s2 2p6 3s2 3p5 1 electron added to the 3p orbital Cl¯1s2 2s2 2p6 3s2 3p6 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. TITANIUMTi 1s2 2s2 2p6 3s2 3p6 4s2 3d2 Ti+ 1s2 2s2 2p6 3s2 3p6 4s1 3d2 Ti2+ 1s2 2s2 2p6 3s2 3p6 3d2 Ti3+ 1s2 2s2 2p6 3s2 3p6 3d1 Ti4+ 1s2 2s2 2p6 3s2 3p6

  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 1s2, 2s2, 2p6, 3s1.

  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 CHEMICALionic (or electrovalent) strong bonds covalent dative covalent (or co-ordinate) metallic PHYSICALvan 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¯ 1s2 2s2 2p6 3s1 1s2 2s2 2p6 1s2 2s2 2p6 3s2 3p5 1s2 2s2 2p6 3s2 3p6 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 theresulting ions are held together in a crystal lattice by electrostatic attraction.

  16. Physical properties of ionic compounds Melting point very high A large amount of energy must be put in to overcome the strong electrostatic attractions and separate the ions. Strength Very brittle Any dislocation leads to the layers moving and similar ions being adjacent. The repulsion splits the crystal. Electrical don’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. Solubility Insoluble 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. Na+ Na+ Na+ Na+ Na+ Na+ Na+ Na+ Na+ Na+ Cl- Cl- Cl- Cl- Cl- Cl- Cl- Cl- Cl- Cl- IONIC COMPOUNDS - ELECTRICAL PROPERTIES SOLID IONIC COMPOUNDS DO NOT CONDUCT ELECTRICITY 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 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. COVALENT BONDING Definition consists 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 Formation between atoms of the same element N2, O2, diamond, graphite between atoms of different elements CO2, SO2 on the RHS of the table; when one of the elements is in the CCl4, SiCl4 middle of the table; with head-of-the-group elements BeCl2 with high ionisation energies; + +

  20. COVALENT BONDING • 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 AlCl3 • others share only some - if they share all they will exceed their “octet” eg NH3 and H2O • 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” eg PCl5 and SF6

  21. SIMPLE MOLECULES 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. orbital containing 1 electron orbital containing 1 electron overlap of orbitals provides a region in space which can contain a pair of electrons The greater the overlap the stronger the bond.

  22. H H H H H HYDROGEN H 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 Another hydrogen atom also needs one electron to complete its outer shell WAYS TO REPRESENT THE MOLECULE PRESSING THE SPACE BAR WILL ACTIVATE EACH STEP OF THE ANIMATION

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

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

  25. H H H O H O H H WATER WAYS TO REPRESENT THE MOLECULE Each hydrogen atom needs one electron to complete its outer shell O 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 PRESSING THE SPACE BAR WILL ACTIVATE EACH STEP OF THE ANIMATION

  26. H N H H AMMONIA H H N H H N H H atom needs three electrons to complete its outer shell each atom needs one electron 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 H N H H

  27. H O H WATER H O H O H H atom needs two electrons to complete its outer shell each atom needs one electron 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

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

  29. SIMPLE COVALENT MOLECULES Bonding Atoms are joined together within the molecule by covalent bonds. Electrical Don’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 point Low - intermolecular forces (van der Waals’ forces) are weak; they increase as molecules get a larger surface area e.g. CH4 -161°C C2H6 - 88°C C3H8 -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

  30. DATIVE COVALENT (CO-ORDINATE) BONDING 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 base a lone pair donor Lewis acid a lone pair acceptor Ammonium ion, NH4+ 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 NH3BF3 Boron has an incomplete shell in BF3 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 K Mg

  34. METALLIC PROPERTIES Metals are excellent conductors 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. MOBILE ELECTRON CLOUD ALLOWS THE CONDUCTION OF ELECTRICITY

  35. METALLIC PROPERTIES Metals can have their shapes changed relatively easily MALLEABLECAN BE HAMMERED INTO SHEETS DUCTILECAN 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.

  36. METALLIC PROPERTIES 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 PERIODS Na (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 Na+ Mg2+ Al3+ 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. METALLIC PROPERTIES 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 Li+ Na+ K+ 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.

  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. C 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... A covalent bond will repel another covalent bond BOND BOND PAIRS SHAPE ANGLE(S) EXAMPLE 2 LINEAR 180º BeCl2 3 TRIGONAL PLANAR 120º AlCl3 4 TETRAHEDRAL 109.5º CH4 5 TRIGONAL BIPYRAMIDAL 90º & 120º PCl5 6 OCTAHEDRAL 90º SF6

  40. O O O 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 As a result of the extra repulsion, bond angles tend to be slightly less as the bonds are squeezed together.

  41. H BOND PAIRS 3 LONE PAIRS 1 TOTAL PAIRS 4 N H H N H N N H H 107° H H H H AMMONIA • 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 N H ANGLE... 107° SHAPE... PYRAMIDAL H H

  42. N H H H H H H H H H N N+ N+ N N SHAPES OF IONS REVIEW BOND PAIRS 3 PYRAMIDAL LONE PAIRS 1 H-N-H 107° NH3 NH4+ BOND PAIRS 4 TETRAHEDRAL LONE PAIRS 0 H-N-H 109.5° BOND PAIRS 2 ANGULAR LONE PAIRS 2 H-N-H 104.5° NH2-

  43. O C 180° O C O BOND ANGLE... SHAPE... 180° LINEAR MOLECULES WITH DOUBLE BONDS 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 O O C 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 PAIRS 2 LONE PAIRS 0 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.

  44. F BrF3 F BOND PAIRS 3 LONE PAIRS 2 ’T’ SHAPED ANGLE <90° Br F Br F F Br F F F F F F O O SO42- BOND PAIRS 4 LONE PAIRS 0 TETRAHEDRAL ANGLE 109.5° O- S O S O- O O- O- OTHER EXAMPLES BrF5 BOND PAIRS 5 LONE PAIRS 1 ‘UMBRELLA’ ANGLES 90° <90° F F Br F F F

  45. 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 BF3 SiCl4 PCl4+ PCl6- SiCl62- H2S ANSWERS ON NEXT PAGE

  46. TEST QUESTIONS 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 BF3 3 bp 0 lp 120º trigonal planar boron pairs up all 3 electrons in its outer shell 4 bp 0 lp 109.5º tetrahedral silicon pairs up all 4 electrons in its outer shell 4 bp 0 lp 109.5º tetrahedral as ion is +, remove an electron in the outer shell then pair up 6 bp 0 lp 90º octahedral as the ion is - , add one electron to the 5 in the outer shell then pair up 6 bp 0 lp 90º octahedral as the ion is 2-, add two electrons to the outer shell then pair up 2 bp 2 lp 92º angular sulphur pairs up 2 of its 6 electrons in its outer shell - 2 lone pairs are left SiCl4 PCl4+ PCl6- SiCl62- H2S

  47. Shapes

  48. Shapes

  49. 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 distortedshapes. • By deducing the number of bonded electron pairs and lone pairs of electrons, the shape of a molecule may be predicted. • BF3 is trigonal planar; CH4 and NH4+ are tetrahedral; SF6 is octahedral; H2O is non-linear (V-shaped/bent); CO2 is linear and ammonia, NH3, as pyramidal

More Related