1. Section 6.5 – Molecular Compounds Binary molecular compounds are composed of two nonmetallic atoms. Because atoms can combine in different ratios (for example CO and CO 2 ) we use prefixes to help distinguish between compounds. CO is carbon monoxide CO 2 is carbon dioxide CCl 4 is carbon tetrachloride Note the –ide ending (similar to how an anion works, but these aren’t ionic compounds) PrefixNumber mono-1 di-2 tri-3 tetra-4 penta-5 hexa-6 hepta-7 octa-8 nona-9 deca-10 1

2. Section 6.5 – Molecular Compound Naming To convert a name to a formula, write the correct symbols for the two elements, then add appropriate subscripts. If there is just one of the first atom, you don’t need to write the mono-, it is assumed. But for the second atom, if there is one, use mono- Ex: tetraiodine nonoxide is I 4 O 9 sulfur trioxide SO 3 phosphorus pentafluoridePF 5 Ex: N 2 O dinitrogen monoxide PCl 3 phosphorus trichloride SF 6 sulfur hexafluoride H 2 O dihydrogen monoxide 2

3. Molecular Bonding - Acids Here is a list of some of the most common acids which have covalent bonds and their names, which don’t always follow the standard naming convention. HCl Hydrochloric acid H 2 SO 4 Sulfuric acid HNO 3 Nitric acid CH 3 COOH Acetic acid (also written HC 2 H 3 O 2 ) H 3 PO 4 Phosphoric acid H 2 CO 3 Carbonic acid These are the most common ones and good to memorize 3

4. Single Covalent Bonds Hydrogen is the simplest model of a covalent bond H· + ·H H:H Hydrogen Hydrogen Hydrogen atom atom molecule H : H Each Hydrogen has one electron and they share them to form a single covalent bond. The single covalent bond can be represented by the pair of electrons or as a dash as shown below H:H or H-H Each dash represents a pair of shared electrons. 4

5. Conventions for naming The chemical formulas of ionic compounds describe formula units (Example: NaCl is a formula unit) The chemical formulas of covalent compounds describe molecules. (Example H 2 O is a molecule) Ionic compounds do not have molecular formulas because they are not composed of molecules. What does that mean? Example: Ionic copper(II) oxide is composed of equal numbers of Cu 2+ and O 2- ions in a crystal lattice. The formula unit shows the lowest whole-number ratio of Cu 2+ to O 2-, which is 1:1, CuO. 5

6. Ionic versus Molecular Compounds 6

7. Covalent Molecules Combinations of atoms of the nonmetallic elements in groups 4A, 5A, 6A and 7A of the periodic table are likely to form covalent bonds. octet rule for covalent bonding: Chemist Gilbert Lewis summarized this tendency in his formulation of the octet rule for covalent bonding: Sharing of electrons occurs if the atoms involved acquire the electron configuration of noble gases. 7

8. Covalent Bonding – Diatomic Gas Fluorine 8

9. In a water molecule, two hydrogen atoms form one single covalent bond each with one oxygen atom. Note how the O atom ends up with eight electrons around it. Covalent molecules will form if each atom will end up with 8 electrons around it (except H). Each dot is one electron. Each line is two electrons. 9

10. In an ammonia molecule, NH 3, three H atoms form single covalent bonds with one N atom. Note how the N has eight electrons around it. When you consider the N and the three H’s, you can see the 2s and 2p orbitals are now full with eight electrons. 10

11. A methane molecule has four carbon-hydrogen bonds. In each bond, C and H share the 1s e- from the hydrogen and a 2s or 2p e- from the carbon. Normally C would start with 1s 2 2s 2 2p 2 configuration, but by promoting one 2s e- to 2p, resulting in 1s 2 2s 1 2p 3 it can create a stable octet with the four H atoms. 11

12. Covalent Bonding Different bonding models for methane, CH 4. Models are NOT reality. Each has its own strengths and limitations. 12

13. Review: The Octet Rule and Covalent Compounds 13

14. A single bond The Diatomic Fluorine Molecule F F 1s 2s 2p seven Each has seven valence electrons FF 14

15. Some double bonds: Carbon Dioxide A carbon dioxide molecule has two C=O bonds Note how C and the O’s each have 8 electrons now 15

16. An exception to the rule: The Diatomic Oxygen Molecule O O 1s 2s 2p six Oxygen has six valence electrons. You would think O 2 would form a double bond by the looks of it, but experiments show its nonstandard and has two unpaired e- O O 16

17. A triple bond: The Diatomic Nitrogen Molecule N N 1s 2s 2p five Each has five valence electrons N N 17

18. I Bring Clay For Our New House 18

19.  Lewis structures show how valence electrons are arranged among atoms in a molecule.  Lewis structures reflect the central idea that stability of a compound relates to noble gas electron configuration (atoms will react if they can arrange themselves to have 8 electrons around them).  Shared electron pairs are covalent bonds and can be represented by two dots (:) or by a single line ( - ) Lewis Dot Structures 19

20. The HONC Rule  HH  Hydrogen (and Halogens) form one covalent bond  O  Oxygen (and sulfur) form two covalent bonds  One double bond, or two single bonds  N  Nitrogen (and phosphorus) form three covalent bonds  One triple bond, or three single bonds, or one double bond and a single bond  C  Carbon (and silicon) form four covalent bonds.  Two double bonds, or four single bonds, or a triple and a single, or a double and two singles 20

21. C H H H Cl........ Completing a Lewis Structure: CH 3 Cl  Add up available valence electrons:  C = 4, H = (3)(1), Cl = 7 Total = 14  Join peripheral atoms to the central atom with electron pairs.  Complete octets on atoms other than hydrogen with the remaining electrons  Draw carbon as the central atom (it wants the most bonds, 4)...... Check: Final structure should have 14 e- 21

22. Coordinate Covalent Bonds A covalent bond in which one atom contributes both bonding electrons is called a coordinate covalent bond. This is signified by showing coordinate covalent bonds as arrows that point from the atom donating the pair of electrons to the atom receiving the bond. Many polyatomic cations and anions contain both covalent and coordinate bonds. NH 4 + is an example. 22

23. Carbon Monoxide – coordinate covalent bonding In a coordinate covalent compound, one atom contributes both electrons of a bonding pair. In carbon monoxide, which atom contributes two electrons in one of the carbon-oxygen bonds? :C O: Triple bond – one of them is a coordinate bond 23

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25. exceptions to octet rule 25

26. How to represent ions (covalent bonds within the ion) 26

27. Electron dot structure of the sulfite ion (SO 3 2- ) each O has 6 each S has 6 plus 2 extra = 26 e- Sulfur has to create one coordinate covalent bond to make this work. 27

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29. Bond Dissociation Energies Large amounts of heat are given off when hydrogen atoms combine to make H 2, which implies that the product is more stable than the reactants. If you try to break H 2 apart, it will require a large amount of energy to do it. The same thing is true for Carbon. A typical C-C single covalent bond has a bond dissociation energy of 347 kJ. Since Carbon forms such strong C-C bonds, that explains why it’s compounds are so stable. See table of bond dissociation energies on the next page. Note which one is weakest. 29

30. Bond Length and Bond Energy What trend do you notice? 30

31. Resonance  Occurs when more than one valid Lewis structure can be written for a particular molecule, such as ozone, below.  These are resonance structures. The actual structure is an average or a blend of the resonance structures. 31

32. Resonance in Benzene, C 6 H 6 Each of these junctions represents where a Carbon is 32

33. Exceptions to the Octet Rule NO 2 (also known as smog in LA) has one unpaired electron, so it is an exception to the octet rule. 33

34. Oxygen: an exception to the octet rule O O O The measured distance between oxygen atoms indicates that O 2 does have some double bond character. This suggests that oxygen is a hybrid of the two structures shown on this page. 34

35. Other exceptions to the octet rule F – B – F F BF 3 is deficient by 2 e-. Sometimes Phosphorus or Sulfur expand the octet to include 10 or 12 electrons. –Examples are PCl 5 and SF 6 P S 35

36. Section 16.3: Polar bonds and molecules Covalent bonds involve sharing electrons between two atoms. Sometimes the sharing is equal and the electron resides halfway in between the atoms, as in a diatomic gas like N 2, Cl 2, etc. This is called a nonpolar covalent bond. This sharing isn’t always equal, because one atom may pull harder than the other atom, and then the electrons will not be in the middle. If the bonding electrons are shared unequally, this is called a polar covalent bond (or just a polar bond). 36

37. Polar Bonds The greater the electronegativity value, the greater the ability of an atom to attract electrons to itself. A high electronegativity atom is not “stealing” electrons as in the ionic case, but it is moving them in its direction. Consider HCl. Hydrogen has an electronegativity of 2.1 and Chlorine has an electronegativity of 3.0. These values are quite different, so the covalent bond in HCl is polar. The shared electrons are pulled in the direction of Cl, because it is more electronegative. This can be represented as follows:   H – Cl H - Cl 37

38. Bond Polarity Electronegativity Differences and Bond Types Electronegativity Difference Range Most probable type of bond Example 0.0 – 0.4Nonpolar covalentH – H (2.1-2.1=0.0) 0.4 – 1.0Moderately polar covalent  H – Cl (3.0-2.1=0.9) 1.0 – 2.0Very polar covalent  H – F (4.0-2.1=1.9) ≥ 2.0IonicNa + Cl - (3.0-0.9=2.1) 38

39. Sample Problem 16-4 Find which type of bond (nonpolar covalent, moderately polar covalent, very polar covalent or ionic) will form between each of the following? A)N (3.0) and H (2.1) Δ = 0.9, moderately polar covalent B) F (4.0) and F (4.0) Δ = 0.0, nonpolar covalent C) Ca (1.0) and O (3.5) Δ = 2.5, ionic D) Al (1.5) and Cl (3.0) Δ = 1.5, very polar covalent 39

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41. Polar Molecules The presence of a polar bond in a molecule makes the entire molecule polar. That means one end of the molecule is slightly negative and the other end slightly positive. A molecule that has two poles is called a dipolar molecule or a dipole. If this kind of molecule is placed in an electric field, the molecules orient themselves with respect to the positive and negative plates creating the field. 41

42. Polar Molecules Some molecules have dipoles, but their polarities line up in such a way that they cancel out. Carbon dioxide is one such example O = C = O 42

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44. Intermolecular Attractions- van der Waals forces The weakest intermolecular attractions are van der Waals forces. These consist of two possible types, London dispersion forces and dipole interactions. London dispersion forces, (weakest of all intermolecular interactions) are caused by the motion of electrons. The strength of dispersion forces increases as the number of electrons increases. For halogens, which have more e- in their outer shell, the major attraction between them is dispersion forces. These forces are weaker for F and Cl (gases at STP). They are stronger for Bromine, a liquid at STP, and even stronger for Iodine, a solid at STP. 44

45. Dipole interaction forces The second type of van der Waals force is the dipole interaction, when polar molecules are attracted to one another. The positive region of one molecule is attracted to the negative region of another. HCl molecules 45

46. Hydrogen Bonds A hydrogen bond is an attractive force where a hydrogen which is covalently bonded to a very electronegative atom (meaning the H has a slight δ + charge on it) is also weakly bonded to an unshared electron pair of another atom (pair has δ – charge to it). This happens because when H bonds to O, F or N, the very polar bond leaves the H very electron deficient, with essentially an exposed nucleus (a proton) with no electrons. The H nucleus is then attracted to a negatively charged unshared electron pair on another atom. The resulting hydrogen bond is only about 5% of the strength of a regular covalent bond, but it is still the strongest of the intermolecular forces. –This is what causes water to be a liquid at room temperature 46

47. Hydrogen bonds in water Hydrogen has valence e- that are not shielded from the nucleus by another layer of electrons. Water has this type of interaction because the hydrogens have a slightly + charge and the oxygen has a slightly – charge. This relatively strong interaction is called a hydrogen bond. 47

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50. Hydrogen Bonding videos Hydrogen bonding http://www.youtube.com/watch?v=lkl5cbfqFRM Water drops on a penny http://www.youtube.com/watch?v=tv4Jrc06yLA&feature=fvwrel 50

51. Intermolecular Forces Summary Weakest – London dispersion forces Middle – dipole-dipole interactions Strongest – hydrogen bonds But all three are still much weaker than a covalent bond (max 5% of the strength of an average covalent bond) 51

52. Intermolecular Attractions and Molecular Properties The physical properties of a compound depend on the type of bonding it has – ionic or covalent. Here are some comparisons of physical properties Characteristics of Ionic and Covalent Compounds CharacteristicIonic CompoundCovalent Compound Representative unitFormula unitMolecule Bond formationTransfer of one or more electrons between atoms Sharing of electron pairs between atoms Type of elementsMetallic & nonmetallicNonmetallic Physical stateSolidSolid, liquid and gas Melting pointHigh (usually > 300 C)Low (usually < 300 C) Solubility in waterUsually highHigh to low Electrical conductivity of aqueous solution Good conductorPoor conductor or nonconducting 52