Bonding

1. Bonding Covalent Bonding

2. Types of Chemical Bond • When atoms combine to achieve more stable structures, three types of bonding are possible • Ionic Bond – results when metallic atoms combine with non-metallic atoms to form and ionic lattice • Metallic Bond – results when metallic atoms combine to form a metallic lattice • Covalent Bond – results when non-metallic atoms combine to form either molecules or covalent lattices.

3. Covalent Bonds • Non-metal atoms are held together by a covalent bond. • Covalent bonds are formed when atoms, rather than gaining or losing electrons, share them. There is no transfer of electrons. • Watch a covalent bond form

4. Covalent Molecular Substances • Covalent molecular substances are held together by covalent bonds and are all around us. • Most non-metals elements and compounds of non-metals have; • Low melting and boiling temperatures (gases or liquids at room temp and pressure) • Poor conductors of electricity • Therefore structure • The forces of attraction between the particles must be very weak • No charged particles are free to move through the substance

5. Covalent Molecular Substances • The basic units of covalent molecular substances are groups of atoms called molecules. All molecules within a structure are identical • Adjacent atoms within a molecule share electrons in order to achieve a full outer shell • Electrical attraction between the nuclei of adjacent atoms and the shared electrons causes the atoms in a molecule to be held together. This force of attraction is called covalent bonding. • The overall charge on each molecule is zero and so adjacent molecules are held together by weak intermolecular forces

6. The Molecule • A molecule is a group of non-metallic atoms held together by covalent bonds. The atoms are combined in a fixed ratio and are electrically neutral. • Identical elements held together by covalent bonds – can be found in various arrangements (monatomic, diatomic)

7. The Molecule • Monatomic – He • Diatomic – H2 Hydrogen, O2 Oxygen, N2 Nitrogen, F2 Fluorine, • Other molecular elements - S8 Sulphur, P4 Phosphorus • Covalent Molecular Compounds – H2O water, CO2 Carbon Dioxide, CH4 Methane

8. Covalent Bonding • Sharing electrons to form full outer shell for stability • Most elements in the second and third periods of the periodic table need eight outer shell electrons for stability • Non-metallic atoms generally have high electronegativities(they are able to attract electrons easily)

9. Lewis Dot Diagrams • When writing out molecules… • Bonding pairs can be represented by the electrons themselves or by a single line • If there is more than one bonding pair, two lines are used • Up to four lines can be used • EACH LINE REPRESENTS A BONDING PAIR OF ELECTRONS IN A MOLECULE • 1 line=2 electrons • Lines represent bonding pairs…you must still draw out lone pairs

10. 3 basic rules • Sum up valence electrons of all atoms in the molecule • Use a pair of electrons to form a bond between atoms…use a line to represent bond • Arrange remaining electrons to satisfy the duet rule for hydrogen and the octet rule for other atoms (Verify that each atom has a complete set of electrons )

11. Elements is group 7A have how many valence electrons? • How many electrons would they share in a covalent bond with an atom in group 7A? • How many covalent bonds do are formed between bonding atoms in 7A? • Elements in group 6A have how many valence electrons? • How many electrons would they have to share in a covalent bond with another atom in group 6A? • How many covalent bonds do are formed between bonding atoms in 6A?

12. Elements in group 5A have how many valence electrons? • How many electrons would they have to share in a covalent bond with another atom in group 5A? • How many covalent bonds do are formed between bonding atoms in 5A? • Elements in group 4A have how many valence electrons? • How many electrons would they have to share in a covalent bond to be stable? • How many covalent bonds are formed between bonding atoms in group 4A?

13. Review • Complete revision questions page 113 (1)

14. Structural Formula • Structural Formulae • Substitute a dash for each shared electron pair • Double and triple covalent bonds are represented by double and triple dashes respectively

15. Cl Cl It is called aSINGLE BOND circle the electrons for each atom that completes their octets

16. O O two bonding pairs, making a doublebond

17. O O O O = For convenience, the double bond can be shown as two dashes.

18. Cl Cl Single bonds are abbreviated with a dash circle the electrons for each atom that completes their octets

19. Review • Complete the revision questions page 115 (2 – 4)

20. Electron dot diagrams for Molecular Compounds • Carbon Dioxide – Carbon is the central atom in a carbon dioxide molecule. It forms a double bond with each of the oxygen atoms • Draw the dot diagram and structural formula

21. Electron dot diagrams for Molecular Compounds • Water – In each H2O molecule, an oxygen atom shares electrons with two hydrogen atoms so that they all achieve complete outer shells. • The lone pairs of unbonded electrons in the oxygen atom affect the shape of the molecule and are shown in the structural formula by the two ‘unbonded’ dashes. • Draw the dot diagram and structural formula

22. Electron dot diagrams for Molecular Compounds • Methane – formula CH4. Carbon forms a single bond with each of the four hydrogen atoms. All atoms achieve complete outer shells. • Draw the dot diagram and structural formula

23. Electron dot diagrams for Molecular Compounds • Ammonia – formula NH3. Nitrogen forms a single bond with each of the three hydrogen atoms, in addition to having one lone pair. All atoms achieve complete outer shells. • Draw the dot diagram and structural formula

24. Review • Complete the revision questions page116 (5 – 7)

25. Molecular Models • Ball and stick • Space filling • Computer generated • Shape diagram • Electron dot diagrams and structural formulae fail to represent the three-dimensional shapes of molecules. All discrete molecules have a definite three-dimensional shape. • Note the tetrahedral shape of methane

26. Valence Shell electron Pair Repulsion Theory • VSEPR theory – a simple basis for predicting molecular geometry. • Outer shell electrons • The electron pairs repel each other and take up positions as far from one another as possible

29. Names of Covalent Compounds Prefixes are used • In the names of covalent compounds. • Because two non-metals can form two or more different compounds. • Examples of compounds of N and O: • NO nitrogen oxide • NO2 nitrogen dioxide • N2O dinitrogen oxide • N2O4dinitrogentetroxide • N2O5dinitrogenpentoxide

30. Naming Covalent Compounds • STEP 1 Name the first non-metal by its element name. • STEP 2 Name the second non-metal with an ide ending • STEP 3 Add prefixes to indicate the number (from subscripts) of atoms of each nonmetal. Mono is usually omitted.

31. Naming Covalent Compounds • Work through the Sample Problems page 119

32. Review • Complete the revision questions page 120 (9 – 12)

33. Electronegativity + – 0 0 H Cl H H

34. Electronegativity • The attraction for electrons • Different atoms have different electron attracting abilities. • The relative attraction that an atom has for shared electrons in a covalent bond is known as electronegativity • Linus Pauling developed a scale in which the most electronegative atom, Fluorine, is assigned a value of 4.0. Fluorine attracts electrons almost twice as well as hydrogen – electronegativity 2.1 • No values are assigned to the noble gases.

35. Pauling Scale of Electronegativities • The following electronegativity trends may be seen in the periodic table • Electronegativities increase from left to right within a period • Electronegativities decrease from top to bottom within a group • Metals generally have lower electronegativities than non-metals

36. H H H H H H Na Na Na Na Cl Cl Cl Cl O O O The basic units: ionic vs. covalent • Ionic compounds form repeating units. • Covalent compounds form distinct molecules. • Consider adding to NaCl(s) vs. H2O(s): • NaCl: atoms of Cl and Na can add individually forming a compound with million of atoms. • H2O: O and H cannot add individually, instead molecules of H2O form the basic unit.

37. Intramolecular forces occur between atoms Intermolecular forces occur between molecules Holding it together Q: Consider a glass of water. Why do molecules of water stay together? A: there must be attractive forces. Intramolecular forces are much stronger • We do not consider intermolecular forces in ionic bonding because there are no molecules. • We will see that the type of intramolecular bond determines the type of intermolecular force.

38. + + 0 – 0 – H2 HCl LiCl H H Cl H [Cl]– [Li]+ I’m not stealing, I’m sharing unequally • We described ionic bonds as stealing electrons • In fact, all bonds share – equally or unequally. • Note how bonding electrons spend their time: covalent (non-polar) polarcovalent ionic • Point: the bonding electrons are shared in each compound, but are not always shared equally. • The greek symbol  indicates “partial charge”.

39. Electronegativity • Recall that electronegativity is “a number that describes the relative ability of an atom, when bonded, to attract electrons”. • The periodic table has electronegativity values. • We can determine the nature of a bond based on EN (electronegativity difference). • EN = higher EN – lower EN NBr3: EN = 3.0 – 2.8 = 0.2 (for all 3 bonds). • Basically: a EN below 0.5 = covalent, 0.5 - 1.7 = polar covalent, above 1.7 = ionic • Determine the EN and bond type for these: HCl, CrO, Br2, H2O, CH4, KCl

40. Electronegativity Answers HCl: 3.0 – 2.1 = 0.9 polar covalent CrO: 3.5 – 1.6 = 1.9 ionic Br2: 2.8 – 2.8 = 0 covalent H2O: 3.5 – 2.1 = 1.4 polar covalent CH4: 2.5 – 2.1 = 0.4 covalent KCl: 3.0 – 0.8 = 2.2 ionic

41. + + + – + – – – + – + – Electronegativity & physical properties CaCl2 would have a lower melting/boiling point: CaCl2 = 3.0 – 1.0 = 2.0 CaF2 = 4.0 – 1.0 = 3.0 Note: other factors such as atomic size within molecules also affects melting and boiling points. EN is an important factor but not the only factor. It is most useful when comparing atoms and molecules of similar size. • Electronegativity can help to explain properties of compounds like those in the lab. • Lets look at HCl: partial charges keep molecules together. LiBr would have a lower melting/boiling point: KCl = 3.0 – 0.8 = 2.2 LiBr = 2.8 – 1.0 = 1.8 • The situation is similar in NaCl, but the attraction is even greater (EN = 2.1 vs. 0.9 for HCl). H2S would have a lower melting/boiling point: H2O= 3.5 – 2.1 = 1.4 H2S = 2.5 – 2.1 = 0.4 • Whichwouldhaveahighermelting/boilingpoint? • NaCl because of its greater EN. • For each, pick the one with the lower boiling point a) CaCl2, CaF2 b) KCl, LiBr c) H2O, H2S

42. CaCl2 would have a lower melting/boiling point: CaCl2 = 3.0 – 1.0 = 2.0 CaF2 = 4.0 – 1.0 = 3.0 Note: other factors such as atomic size within molecules also affects melting and boiling points. EN is an important factor but not the only factor. It is most useful when comparing atoms and molecules of similar size. LiBr would have a lower melting/boiling point: KCl = 3.0 – 0.8 = 2.2 LiBr = 2.8 – 1.0 = 1.8 H2S would have a lower melting/boiling point: H2O= 3.5 – 2.1 = 1.4 H2S = 2.5 – 2.1 = 0.4

44. Non-polar and polar covalent bonds • Non-polar – bonding electron pair are shared equally and is uniformly distributed between the nuclei of two bonded atoms.

45. Methane The molecule methane has four Carbon-Hydrogen single covalent bonds. These covalent bonds are called non-polar covalent bonds because • the electrons shared by the adjacent atoms in the bonds are shared equally • The consequence of this equal sharing of electrons is that there is no charge separation (dipole moment • Since there is no charge separation in the covalent bonds this molecule cannot enter into a charge interaction with water and will therefore be hydrophobic.

46. Pentane This molecule is another example of a compound that is composed of Hydrogen and Carbon and is therefore in the class of compounds known as hydrocarbons. The covalent bonds between adjacent atoms are non-polar covalent bonds and this compound is therefore hydrophobic.

47. Non-polar and polar covalent bonds • Polar– bonding electrons are unequally shared and are therefore unsymmetrically distributed between the nuclei of two bonded atoms. Such bonds occur between atoms of different electronegativities. • The shared pair of electrons move closer to the more electonegative atom. • This means that the atom that has greater control of the electron pair becomes slightly negatively charged, while the atom that lost some control of the electron pair becomes slightly positively charged. • Polar covalent bonds have a charge separation or a bond dipole.

48. Water • Water molecule having polar covalent bonds between the Oxygen atom and the Hydrogen atoms.

49. What are polar covalent Bonds? • Polar covalent bonds are a particular type of covalent bond. • In a polar covalent bond, the electrons shared by the atoms spend a greater amount of time, on the average, closer to the Oxygen nucleus than the Hydrogen nucleus. This is because of the geometry of the molecule and the great electronegativity difference between the Hydrogen atom and the Oxygen atom. • The result of this pattern of unequal electron association is a charge separation in the molecule, where one part of the molecule, the Oxygen, has a partial negative charge and the Hydrogens have a partial positive charge. • You should note this molecule is not an ionbecause there is no excess of proton or electrons, but there is a simple charge separation in this electrically neutral molecule. • Water is not the only molecule that can have polor covalent bonds. Examples of other molecules that have polar covalent bonds are Peptide bonds and amines . • The biological consequence of polar covalent bonds is that these kinds of bonds can lead to the formation of a weak bond called a hydrogen bond.

50. What are polar covalent Bonds? • A pure covalent bond is not purely covalent, since there is not an equal sharing of electrons. • It has some characteristics of ionic bonding, although the transfer of electrons from one atom to the other is not complete. • If the difference in electronegativity between two cations is 2 or greater on the Pauling scale, an ionic bond does form