chemical Bonding, covalent and ioinic bonding

this lecture is about chemical binding, formation and properties

1. Bonding

2. Bonding A molecule/compound is formed when elements bond together With the aim in most cases to achieve a full outer shell of electrons Types of bonding: Covalent Ionic Metallic

3. Ionic bonding Ionic bonding is the electrostatic attraction between positive and negative ions The transfer of electrons from one atoms to another forms ions Atoms lose or gain electrons in order to have a full outer shell of electrons Strong electrostatic attraction holds the positive and negative ions together – ionic bonding The number of electrons gained/loss can be predicted using the atoms position in the periodic table

4. Group 2 2 + ions Group 1 1+ ions Group 6 2- ions Group 7 1- ions

5. Group of atoms with an overall charge – compound ion Compound ions Sulfate, SO 4 2- Hydroxide, - OH Ammonium, + NH 4 Carbonate, CO 3 2 - Nitrate, NO 3 -

6. Ionic compounds Ionic compounds are made up of positive and negative components Overall the charge of the compound is zero The charges on the ions can be used to determine the formula of the compound to give an overall neutral compound Potassium chloride – K + (1+) and Cl - (1-) ions form KCl, as the charges balance to give 0 charge. Sodium sulfate – Na + (1+) and SO 4 2- (2-) ions form Na 2 SO 4 . Two sodium ions are needed to balance out the charge of the of the 2- sulfate ion.

7. Covalent bonding Covalent bond is the strong electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atom In a single covalent bond two atoms share one pair of electrons in order to get a full out shell of electrons Br Br Methane, CH 4 Water, H 2 O Bromine, Br 2

8. Multiple covalent bonds Some molecules share more than one pair of electrons – double and triple bonds can be formed when two and three pairs of electrons are shared respectively Carbon dioxide, CO 2 Ethene , C 2 H 4 Nitrogen, N 2

9. Dative covalent Dative covalent bond is a bond that contains a shared pair of electrons with both electrons supplied by one atom. Once formed, there is no different between a ‘dative’ and ‘normal’ bond Ammonium, N H 4 Aluminium chloride, Al 2 Cl 6

10. Bond length The bond length is dependent on the strength of the bond The stronger the bond the shorter the bond length The average bond enthalpy is a measure of how strong a covalent bond is Bond enthalpy is measured in kJ mol -1 Hydrogen – halide bond, stronger the bond shorter the bond length Bond Bond enthalpy (kJ mol -1 ) Bond length ( Å ) H – F 562 0.937 H – Cl 431 1.267 H – Br 366 1.417 H – I 299 1.631

11. Metallic bonding Metallic bonding is the strong electrostatic attraction between metal ions and the delocalised electrons The outermost electrons are delocalised – free to move. Delocalised electrons are attracted to positive metal ions. Lattice of p ositive metal ions surrounded by a sea of electrons.

12. Shapes The shapes of molecule and ions vary depending on the pairs of electrons in the outer shell of the atom Pairs of electrons arrange themselves arrange themselves as far apart as possible Minimise repulsion Known as Valence-Shell Electron-Pair Repulsion Theory (VSEPR) Lone pair – lone pair repulsion > lone pair – bond pair repulsion > bond pair – bond pair repulsion  

13. Shapes – 2 electron pairs Linear 180 ° Beryllium chloride, BeCl 2

14. Shapes – 3 electron pairs Trigonal Planar Boron trifluoride , BF 3 120 °

15. Shapes – 4 electron pairs Tetrahedral Methane, CH 4 109.5 °

16. Shapes – 4 electron pairs (1 lone pair) Trigonal pyramidal Ammonia, N H 3 106.7 °

17. Shapes – 4 electron pairs (2 lone pairs) Bent Water, H 2 O 104.5 °

18. Shapes – 5 electron pairs Trigonal bipyramidal Phosphorous pentafluoride , PF 5 120 ° 90 °

19. Shapes – 5 electron pairs (1 lone pair) Seesaw Sulfur pentafluoride , SF 5 102 ° 87 °

20. Shapes – 5 electron pairs (2 lone pairs) T-shaped Chlorine t rifluoride , ClF 3 88 °

21. Shapes – 6 electron pairs Octahedral Sulfur hex afluoride, SF 6 90 °

22. Shapes – 6 electron pairs (2 lone pairs) Square planar Xenon tetrafluoride, XeF 4 90 °

23. Electronegativity Electronegativity is the ability of an atom to attract the electron pair in a covalent bond Strongly electronegative atoms include: F, O, Cl, N – have a electronegativity of >3 compare to hydrogen which is 2.2 Electronegativity increases left to right across a Period and up a Grou p Differences in electronegativity's between atoms in a covalent bond – polarised bond The atom with a greater electronegativity will pull the pair of bonding electrons towards it Electronegativity: 2.2 3.0 ( Pauling electronegativity values)

24. Polarity The greater the difference in electronegativity, the more polar the bond Covalent bonds between the same atoms are NOT polar as they have the same electronegativity Atoms with similar electronegativity's such as C (2.6) – H (2.2) are essentially non-polar

25. Polar Molecules A whole molecule can be polar when there is an uneven distribution of charge Some molecules that contain polar bonds are not polar overall – if the shape and symmetry of the molecule cause the dipole to cancel Water - Permanent dipole Partially negative oxygen at one end of the molecule, and partially positive hydrogens at the other end. Carbon dioxide – NO permanent dipole The polar bonds are symmetrical, therefore the charges cancel out, and there is no permanent dipole.

26. Forces between molecules Molecules interact with each other Intermolecular forces are very weak compared to bonds Physical properties of a compound are affected by the intermolecular forces Types of intermolecular force: Permanent dipole - dipole Instantaneous dipole – induced dipole Hydrogen bonding

27. Permanent dipole - dipole Molecules with permanent dipoles interact with each other Weak electrostatic attraction between partial positive and negative charges on neighbouring molecule Molecules align themselves so the partial positive and negative regions are next to each other Three chloromethane molecules – permanent dipole Partially negative chlorine is attracted to the partially positive carbon of the neighbouring molecule

28. Instantaneous dipole - Induced dipole Instantaneous dipole – induced dipole/Van der Waals/London forces Lead to all molecules being attracted to each other They are the weakest intermolecular forces Electrons are always moving, at any point the electrons may be at one side of the charge cloud – temporary dipole This dipole can induce a dipole in a neighbouring atom The two dipoles are attracted to each other As the electrons are always moving the dipoles are formed and destroyed all the time The overall effect is the attraction between molecules

29. Instantaneous dipole - Induced dipole At any moment the electrons can be more towards one side. The temporary dipole induces a dipole in a neighbouring molecule. A domino effect occurs, where each temporary dipole induces another in a neighbouring molecule.

30. Strength of Van der Waals Forces The size and shape of a molecule effects the strength of van der Waals forces. The larger the molecule the stronger the van der Waals forces, due to the larger electron cloud Long straight molecules have stronger van der Waals forces as they can lie closer together and therefore interact more Branched molecules have weaker van der Waals forces as they cant get close together. The stronger the van der Waals forces the greater the boiling point of a compound The LONGER the molecule the HIGHER the boiling point The MORE BRANCHED the molecule the LOWER the boiling point

31. Hydrogen bonding Hydrogen bonding is the strongest intermolecular force Hydrogen bonding occurs when hydrogen is bonded to a very electronegative atom (O, F, N) The electronegative atom draws the electrons away from the hydrogen – polarised bond Hydrogen has a high charge density Hydrogens form a weak bond with lone pairs on electronegative atom δ - δ + δ - δ - δ - δ - δ + δ + δ + δ + δ + δ + δ +

32. H - Bonding Properties Substances with hydrogen bonds have higher melting and boiling points than similar substances without hydrogen bonds. Extra energy is required to break the hydrogen bonds, therefore higher temperatures are needed, for the compounds to melt and boil. Hydrogen bonding also explains why ice is less dense than water. When water cools to form ice, more hydrogen bonds are formed, molecules arrange themselves into an ordered lattice The molecules a.re further apart in the lattice than in liquid water – ice is less dense. H 2 O H 2 S H 2 Se H 2 Te P H 3 NH 3 AsH 3 SbH 3 HF HCl HBr HI

33. Structure There are four types of crystal structure: Ionic Metallic Giant covalent (Macromolecular) Molecular

34. Ionic lattice Atoms arranged in an orderly repeating pattern Ionic compounds contain positive and negative ions The ions in the crystal are arranged in a ‘ giant ’, 3-dimensional structure called an ionic lattice. The lattice is held together by electrostatic attraction between the cations and anions.

35. Properties of ionic compounds T end to dissolve in water – Water is polar so can surround ions , and pull ion away from the lattice. C onduct electricity when molten or dissolved – When ions are molten/dissolved they are free to move and can carry charge. Cannot conduct when solid – ions are fixed so are unable to move. Have high melting points – due to the strong electrostatic forces.

36. Metallic structure Lattice of positive metal ions surrounded by a sea of electrons. Metal atoms are arranged in layers. The atoms are as close together as possible – high density Mg 2+ Mg 2+ Mg 2+ Mg 2+ Mg 2+ Mg 2+ Mg 2+ e - e - e - e - e - e - e - e - e - Lattice of Mg 2+ ions Sea of delocalised electrons

37. Properties of Metals The electrons are free to move around the metal lattice, they can easily move through the metal when a potential difference is applied – i.e. an electric current can flow through the metal . Metals are good thermal conductors, kinetic energy can be transferred between electrons. Metals have high melting points due to the strong electrostatic attraction between positive ions and delocalised electrons. Metals are malleable as the layers of metal ions can slide over one another.

38. Giant Covalent Structures Diamond Diamond is a form of carbon This has a very strong three dimensional structure Every carbon is joined to four others by covalent bonds. This giant molecule can contain billions of atoms. The very regular way that the ions are packed leads to the distinctive regular crystal shape with flat surfaces at fixed angles .

39. Properties of Diamond The strong covalent bonds mean that diamond has a very high melting point : 3800°C. The strong bonds and rigid structure makes diamond hard . As the atoms are not arranged in layers, they cannot slide over each other, so diamonds are not malleable or ductile . Won’t dissolve in any solvent There are no free electrons, so diamond does not conduct electricity

40. Giant Covalent Structures Graphite Graphite (the “lead” inside pencils) is another allotrope of carbon. Each carbon atom is joined with strong covalent bonds to three others, forming sheets of atoms arranged in hexagon patterns . The fourth outer electron of each carbon is delocalised. There are only weak van der Waals forces between the layers.

41. Properties of Graphite The strong covalent bonds between the atoms give graphite a very high melting point: 3600°C The bonds between the layers are longer and weaker than the other bonds - graphite less dense than diamond. The weak bonds between the layers mean that graphite can be split easily in the direction parallel to the layers – can be used as a dry lubricant and in pencils. While three electrons from each carbon are used in covalent bonds, the fourth electron is delocalized over all the atoms in its layer - graphite will conduct electricity and heat. The conductivity parallel to the layers is much greater than in a perpendicular direction.

42. Molecular structures Most covalently bonded substances exist as simple small or medium sized molecules, rather than as giant structures. When simple molecules are heated moderately, the atoms do not get enough kinetic energy to break the strong covalent bonds, and so most molecules do not break apart (decompose) on heating. The forces between molecules – van der Waals or dipole-dipole are much weaker than the covalent bonds. Just a little heat gives enough kinetic energy to separate each molecule from neighbouring molecules. When a molecular substance melts, only the weak intermolecular bonds are broken.

43. Properties of Simple Molecules Low melting and boiling points – only weak intermolecular forces holding the molecules together Do not conduct electricity or heat, as there are no charged particles to carry the charge Molecules with stronger intermolecular forces – hydrogen bonds, have higher melting and boiling points