1. Ionic or Covalent Bond? Ionic compounds form from metals (low electronegativity – electropositive elements) and nonmetals (high electronegativity – electronegative elements). Usually the difference in electronegativity in ionic compounds is greater than 1.8. Molecular compounds with covalent bonds are formed from nonmetals (and metalloids) with low electronegativity differences. SL Chemistry Topic 4 Bonding

2. Ionic or Covalent Bonds?

3. Ionic Bonding The electrostatic attraction between oppositely charged particles that are formed by the transfer of electrons. The electrons are donated or accepted to achieve the noble gas configuration.

4. Ionic Bonding Ionic compounds are held together by strong electrostatic attractions in a strong crystal lattice structure. These strong attractions cause ionic compounds to have high melting points since it takes a large quantity of energy to overcome these forces of attraction.

5. Sodium chloride as an example of an ionic compound Even though chlorine is already a molecule with a noble gas configuration, as sodium burns in it, the energy released when the ionic lattice is formed is sufficient to break the covalent bond.

6. Formulas of ions and ionic compounds Metals lose electrons & form cations Nonmetals gain electrons & form anions Polyatomic ions – more than one element joined by covalent bonds and the charge is often spread (delocalized) over the whole ion.

7. Covalent Bonds The sharing of one or more pairs of electrons so that each atom in a molecule achieves an inert gas configuration. Hydrogen, H 2, is the simplest molecular compound. The two electrons are shared and attracted equally, directly between the nuclei or the atoms.

8. Single Covalent bonds In a single covalent bond, 2 electrons are being shared equally be two different nuclei. The orbitals of the two atoms overlap.

9. Diatomic molecules Many nonmetals exist as diatomic molecules where one or more pair or electrons Is shared. They are: hydrogen, nitrogen, oxygen, flourine, chlorine, bromine, and iodine – HNOFClBrI Ch. 19 p. 545, #13 & 15 due ____________

10. Lewis Structures Lewis structures or electron dot structures show all of the valence electrons Sometimes the bonding is shown for a molecule without the valence electrons, this is not a true Lewis structure. Electrons can be shown as dashes, dots, or small x’s.

11. Electron Dot Structures One method of showing valence electrons for all types of elements is using the electron dot structure. It shows only valence electrons and they are filled by following Hund’s rule. X 1 2 3 4 5 6 7 8

14. Simple molecules with single covalent bonds Carbon forms 4 single covalent bonds with hydrogen or halogens, since they require only one more electron as in methane and carbon tetrachloride –3-D shape is important for each of these molecules Nitrogen forms three single covalent bonds leaving one non- bonded pair or lone pair of electrons as in ammonia

15. Simple molecules with single covalent bonds Oxygen forms two single covalent bonds with 2 unshared pairs as in water Halogens such as fluorine form a single covalent bond leaving 3 unshared pairs of electrons as in hydrogen fluoride or hydrogen chloride

16. How do you know what shape to draw? Typically the first atom listed is the central atom or the base chain as in hydrocarbons. Most everything else is bonded off of the central atom.

17. Steps in drawing a molecule 1. Count and add up the number of valence electrons for each molecule. 2. Test different drawings until each atom has a full octet, except for hydrogen. As the number of pairs of electrons decreases, the number of double or triple bonds would increase. If you give each atom 8 e - and only single bonds and you are one pair short of the valence number you need one double bond. If you are two pairs of e - short, you need a triple bond or two double bonds. Remember that once you are past the e- conf. of phosphorus, you can have an expanded octet.

18. Examples of Lewis structures with single covalent bonds BF 3 BeF 2 H 2 S CH 3 Cl PCl 5 SF 6 XeO 4 ClO 4 -

19. Multiple Covalent Bonds At times more than one pair of electrons are shared so that the noble gas configuration can be achieved. These are called double and triple bonds. Diatomic oxygen and nitrogen have multiple bonds, as does carbon dioxide and the hydrocarbons ethene and ethyne.

20. Multiple Covalent Bonds Co-ordinate covalent bonds (dative) are formed when one of two atoms donates both electrons for a covalent bond. It is like normal covalent bonds in all other ways. SO 2 SO 3

21. Draw some Lewis structure containing multiple bonds CO CO 2 CN - N 2 NO + CO 3 2-

22. Bond Length and Bond Strength Single bonds are the longest and weakest of the three types of bonds. Double bonds are between single and triple. Triple bonds are the strongest, but shortest of the three. Bond strength varies with size of the nucleus of the different atoms and also the interactions between the other atoms than make up the molecules. P. 19, IB Diploma book gives some examples as does page 214 in Masterton.

23. Bond Polarity (not the molecule just each bond) A bond between atoms in diatomic molecules is nonpolar since the attraction for the electrons is equal between the two atoms’ nuclei. When there are two different atoms involved a polar bond is formed and the polarity is determined by the electronegativity of the atoms. –The atom with higher electronegativity will be richer in electrons and be slightly negative. –Polar ends of a molecule are indicated with and Ch. 7 p. 192 ? 1, 3, 5, 7, 11, 21 due ____________

24. Shapes of simple molecules and ions – VSEPR Theory Valence shell electron pair repulsion –Pairs of electrons arrange themselves around the central atom so that they are as far apart from each other as possible. –There will be greater repulsion between non- bonded pairs of electrons than bonded pairs. (They will take up more space. –Double and triple bonds count as one pair of electrons in VSEPR theory. –By lose definition the electrons are called negative charge centers.

25. Shapes and bond angles 2 negative charge centers – linear with a bond angle of 180 o BeF 2 CO 2 C 2 H 2 HCN

26. Shapes and bond angles 3 negative charge centers - two possible shapes –Trigonal planar – 3 bonding pairs or electrons with a bond angle of 120 o –BF 3 –C 2 H 4 –CO 3 2-

27. Shapes and bond angles 3 negative charge centers - 2 nd shape –Bent or V-shaped made of 2 bonding pairs and one non-bonded pair of 120 o (since there are only three areas of charge, the non-bonded pair doesn’t make a difference in bond angle) –SO 2 –NO 2 -

28. Shapes and bond angles 4 negative charge centers – tetrahedral – 4 bonding pairs with a bond angle of 109.5 o –CH 4 –CCl 4 –NH 4 + –BF4 -

29. Shapes and bond angles 4 negative charge centers – triangular pyramid – 3 bonding pairs and one non-bonding pair of electrons with a bond angle of 107 o Ammonia, NH 3

30. Shapes and bond angles 4 negative charge centers – bent or V- shaped – 2 bonding pairs and two non-bonding pair of electrons with a bond angle of 105 o (104.5 o ) H 2 O

31. Shapes and bond angles

32. 5 negative charge centers - trigonal bipyramidal – 90 o, 120 o and 180 o PCl 5 There are several other shapes, but IB doesn’t emphasize. Distorted tetrahedral (see-saw) - 4 bonding pairs with 1 non- bonding T-shaped – 3 bonding pairs with 2 non-bonding pair Linear – 2 bonding pairs with 3 non-bonding pairs of electrons

33. Shapes and bond angles 6 negative charge centers – 90 o and 180 o – octahedral – square planar – 4 bonding pairs with 2 non-bonding as far apart as possible above and below the plane of the molecule

35. Molecular polarity Molecular polarity is dependent on both the individual bond polarity and then also the shape of the molecule Symmetrical molecules with like atoms all the was around tend to be nonpolar even with polar bonds in the molecules. Non-symmetrical molecules, especially with non- bonding pairs of electrons tend to be polar. CO 2 H 2 O CCl 4 CH 3 Cl Ch. 7 #27, 29, 33, 35, 45, 49, 51, 53 due _________

36. Intermolecular Forces Intermolecular forces are weak attractions between molecules that are 10% and less the strength of covalent bonds. They are dependent on the elements that make up the molecules their arrangement and the shape of the final molecule.

37. Intermolecular forces – van der Waal forces Also called dispersion and London forces In any molecule at any one moment in time the electrons may not be evenly dispersed which creates an instantaneous dipole. This can then induce temporary dipoles in adjacent molecules, resulting in a weak attraction between the molecules.

38. Intermolecular forces – van der Waal forces It is the only force of attraction between nonpolar molecules and is extremely weak. As mass increases, van der Waal forces increase.

39. Intermolecular forces – Dipole:dipole forces Continual electrostatic attractions between polar molecules that is stronger than van der Waal forces. –Butane, M r = 58, b.p. -0.5 o C –Propanone, M r = 58, p.t. 56.2 o C –Same molar mass with different intermolecular forces and therefore boiling point –Examples – halogens and hydrocarbons

40. Intermolecular forces – Dipole:dipole forces

41. Dipole-Dipole forces Molecules line up so that the positive and negative ends attract.

42. Induced dipole between a polar molecule and a nonpolar molecule Ch. 10 #35, 36, 38, 40, 42

43. Intermolecular Forces – hydrogen bonding The strongest of the intermolecular forces that is not an actual bond, but a very strong intermolecular force between molecules. Occurs because of the polar bond between hydrogen and either nitrogen, oxygen or fluorine.

44. Intermolecular Forces – hydrogen bonding The electron pair is drawn away from the hydrogen by the highly electronegative element so that all that remains of the hydrogen is the proton of the nucleus since there are no inner electrons The proton then attracts to a nonbonding pair of electrons in an adjacent nitrogen, oxygen or fluorine.

45. Intermolecular Forces – hydrogen bonding & boiling point Water’s boiling point is much higher than HF and NH 3 because of the ratio of attractions between nonbonding pairs and hydrogen protons. –H 2 O has two nonbonding pairs and two nearly bare protons, which results in a greater number of attractions –HF has three nonbonding pairs, but only one nearly bare proton –NH 3 has one nonbonding pair of electrons with three nearly bare protons. Ch. 9 p. 255 #1, 3, 5, 7, 9, 11, 13, 15, 19, 21, 23 due _____________

46. Intermolecular Forces – hydrogen bonding & boiling point

47. Water has a very high boiling point because of hydrogen bonding. Without hydrogen bonding it would be a gas at room temperature. The hydrogen bonds also result in an open framework for ice molecules – a six sided crystal. Water is has its highest density at 4 o C

48. Metallic Bonding Metals consist of a sea of delocalized electrons that surround the cations, the nuclei and nonvalence electrons A metallic bond is the attraction tht two neighboring positive ions have for the delocalized electrons between them.

49. Metallic bonding Metals are malleable, which means they can be bent and reshaped, as well as ductile which is the ability to shape into wires. Metals have these properties because the layers of positive ion shift position and form new bonds as other bonds are broken

50. Metallic bonding Alloys have new/different properties because it disrupts the lattice structure of the metal causing it to be less malleable and less ductile in addition to a higher melting point

51. Metallic Bonding

52. Physical properties as related to bonding type – melting and boiling points The type of bonds between atoms and also the intermolecular forces between molecules determines the melting and boiling points of a substance. Melting a solid disrupts the attractive forces but since they particles are very close to each other, the attractive forces are still fairly high. Boiling results in a complete disruption of attractive forces as a liquid moves to a gas. Impurities in a substance weaken the structure and therefore result in lower melting points.

53. Physical properties as related to bonding type – melting and boiling points The highest melting and boiling points are found in macromolecular structures such as diamond and silicon dioxide/sand. The lattice made of covalent bonds makes for very strong attractions.

54. Physical properties as related to bonding type – melting and boiling points Ionic compounds – have high electrostatic attractions between ions therefore they have high melting points. It also keeps them from being malleable and ductile. Metals also have high melting points due to attractions between ions.

55. Physical properties as related to bonding type – melting and boiling points In simple molecular compounds the intermolecular forces are disrupted between the molecules, but not the covalent bonds. In water, you disrupt H-bonds between the water molecules, but not the covalent bonds between the hydrogens and oxygen. Hydrogen bonds are about 1/10 the strength of a covalent bond and van der Waals’ forces are only about 1/100 the strength of a covalent bond.

56. Physical properties as related to bonding type – melting and boiling points The weaker the intermolecular force, the greater the volatility of a substance. Propane Ethanal ethanol M r 44 44 46 m.p. -42.2 o C 20.8 o C 78.5 o C Polarity non-polar polar polar Intermol. Vdw dipole:dipole H- bond force

57. Physical properties as related to bonding type – solubility “Like dissolves like” Polar substances tend to dissolve polar substances whereas nonpolar substances dissolve other nonpolar substances. Oil and water don’t mix. Organic molecules that have a polar end and a nonpolar end tend to be soluble if they are small. The larger the nonpolar hydrocarbon end in a homologous series, the less soluble in water the compound becomes. Ethanol is water soluble but at 1-hexanol solubility is only slightly soluble in water.

58. Physical properties as related to bonding type – conductivity Conductivity occurs with two types of particles, electrons that are free to move or ions that are free to move. Macromolecular compounds and simple molecules that have electrons held in a fixed position do not conduct electricity.

59. Physical properties as related to bonding type – conductivity Metals and graphite contain delocalized electrons, electrons that aren’t held in a fixed position, so they can move toward a positive gradient. Ions in ionic compounds that are in a molten state can also conduct electricity in the process of electrolysis, but the substances are chemically decomposed in the process. The ions flow toward the oppositely charged electrode. IB likes to ask about this and you need to answer very specifically about electrons or ions.

60. Physical properties as related to bonding type – conductivity