8.2 The Nature of Covalent Bonding

1. 8.2 The Nature of Covalent Bonding

2. How Does the Octet Rule Apply? • When atoms bond covalently, they do so to acquire a more stable electron configuration – usually this means to acquire a full octet of valence electrons. • Electron sharing (covalent bonding) involves sharing pairs of electrons. • When two atoms share one pair of electrons we call it a single covalent bond.

3. Why a “single” covalent bond? • We call it a “single” covalent bond because the atoms are sharing a single pair of electrons. Consider F2. • This shared pair of electrons is frequently represented with a single dash between the atoms – indicating a single bond.

4. How are bonding pairs represented? • Structural formulas will most often use dashes to indicate bonding between atoms. Remember that each dash represents 2 electrons. • Try to draw the structural formula for ammonia, NH3. Begin with drawing electron dot diagrams of the atoms involved: Structural formula: N H H H

5. What about those other electrons? • Note the pair of electrons on nitrogen that are not part of a bond – these electrons are called an unshared pair of electrons. They are also sometimes called lone pairs or nonbonding pairs. • Some elements have several unshared pairs:

6. Using the Octet Rule • Consider the following elements. How many electrons would each of them need to share (or how many single bonds would each need to form) in order to achieve the same electron configuration as neon? C N O F Ne

7. Draw the structural formulas for: • OF2 • SCl2 • N2H4

8. Draw the structural formulas for: • CCl4 • CHCl3 • C2H6

9. Is that all there is? • Remember the diatomic elements – H2, N2, O2, F2, Cl2, Br2, I2? • Draw dot diagrams for 2 oxygen atoms: • How could these two atoms bond together and both attain their octet of electrons? O O

10. Multiple bonds • If the atoms formed 2 bonds between them, then each would have an octet. • When 2 atoms form multiple bonds between them, the bonds are referred to as double or triple covalent bonds (rather than being thought of as multiple single covalent bonds). • In the oxygen example, the structure appears to work out if the oxygens double bond: O O

11. Not all molecules bond as predicted: • Your textbook points out that oxygen appears that it should work out this way, but that experimental evidence suggests that two of the electrons in the oxygen molecule remain unpaired instead of bonding. • Although these rules provide guidelines to help you figure out molecular bonding, this example demonstrates that sometimes things don’t turn out as expected – they “break the rules.” O O

12. Let’s try another one: • Now try predicting nitrogen’s structure, N2. (It doesn’t break the rules!) • Draw dot diagrams for 2 nitrogen atoms: • How could these two atoms bond together and both attain their octet of electrons? N N

13. Now try a larger molecule: • Predict the structure of a molecule of carbon dioxide, CO2. • What types of bonds are these? C O O

14. More Examples • Now try hydrogen cyanide, HCN. • What types of bonds are these? C N H

15. More Examples • Last of all, try formaldehyde, H2CO. (HINT: Atoms that can form more bonds, like carbon, are likely to be central atoms – atoms with lots of atoms bonded to them. Atoms that can only bond once are terminal atoms.) C O H H

16. Bonding for 2nd Period Elements • Consider the valence electrons of the elements in the 2nd period: • Predict how many bonds each atom must form to attain a noble-gas configuration: Li Be B C N O F Ne Li Be B C N O F Ne

17. Bonding for 2nd Period Elements Li Be B C N O F Ne • Can lithium form a covalent bond and reach stability? • Which elements can reach stability (an octet) by forming covalent bonds? • Can fluorine form an ionic bond? • Are the bonds in nitrogen molecules (N2) ionic or covalent?

18. Link to Biology • Earth’s atmosphere is almost 80% nitrogen (N2) but only about 20% oxygen. So why are there relatively few nitrogen-containing compounds as compared to oxygen-containing compounds? • The triple bond in nitrogen molecules (N≡N) is harder to break than a double bond, and considerably harder to break than a single bond (which is what oxygen has been experimentally shown to have ).

19. Link to Biology • In order for plants and animals to get access to nitrogen, the element must be converted to a compound – this is called nitrogen fixing. It occurs naturally when lightning provides sufficient energy for atmospheric nitrogen to react with oxygen, forming nitrogen oxides. These dissolve in rain and fall to the ground where they can be utilized by plants.

20. Link to Biology • There are also some nitrogen-fixing bacteria found in soil that are able to convert atmospheric nitrogen into usable compounds.

21. Are those all the bonds? • If you try drawing carbon monoxide (CO) the result makes it appear that carbon will break the octet rule: • However, if oxygen were to slide over an unshared pair and make it a bonding pair (so oxygen contributes both of the electrons in the bond), then both molecules will achieve an octet of electrons. C +  C=O O C≡O

22. So what is it called? • When one atom contributes both of the electrons in a bond, the bond is called a coordinate covalent bond. • A coordinate covalent bond can be indicated in the structural formula by drawing an arrow that points from the donating atom to the receiving atom instead of just a line for the bond. Thus carbon monoxide is best represented by: C O =

23. Is it really that different? • Once formed, a coordinate covalent bond is just like any other covalent bond. The arrow is used just as a reminder of where the electrons came from. Thus, the three bonds between carbon and oxygen are still considered a triple bond. It was just formed in an unusual manner. • The ammonium and hydronium polyatomic ions are also examples involving coordinate covalent bonds.

24. How do you explain ammonium? • Consider the structures of ammonia and of a hydrogen ion: • Note that the hydrogen ion currently has no valence electrons, but would be stable with 2. Note also that the nitrogen atom in ammonia has an unshared pair of electrons. H+

25. How do you explain ammonium? • The empty s orbital of the hydrogen ion must overlap the filled orbital in ammonia. The electrons from nitrogen are attracted to both the nitrogen nucleus and hydrogen nucleus. A bond forms when electrons are simultaneously attracted to two nuclei – in this case, a coordinate covalent bond: H H+ H N+ H  H

26. How do you explain hydronium? • Consider the structures of water and of a hydrogen ion: • What do you believe the structure of the hydronium ion looks like? Try drawing it: H+

27. Another Example • Try to draw sulfur dioxide (SO2) so that the structure satisfies the bonding requirements of all three atoms. • The only possibility is that one oxygen is double bonded to sulfur, and the other oxygen is single bonded through a coordinate covalent bond where sulfur donates both electrons. This can be drawn as: or O S = O O = S O

28. Another Example • These two structures are called resonance structures. • It was once believed that the molecule would resonate (oscillate back and forth) between these two structures. • Once again, experimental evidence contradicts the prediction. Evidence indicates that the bonds in SO2 are equivalent, which shouldn’t be the case since one is a single bond and the other a double bond. How can we explain this? ↔ O S = O O = S O

29. Another Example • The bonds have been found to be the same length and strength, which can only be explained if you assume that the actual bonding in this molecule is the average of the two structures. In reality, the bonding is a hybrid (mixture) of the two extremes represented by these resonance forms. ↔ O S = O O = S O

30. Resonance Structures • Resonance structures occur when there are two or more valid structural formulas that have the same number of electron pairs for a molecule or ion. They are drawn with a double-headed arrow between the structures to indicate both structures can occur. • You should remember that they are just a way to envision the bonding in certain molecules. ↔ O S = O O = S O

31. Breaking the Rules • There are exceptions to the octet rule. • Some molecules have an odd number of total valence electrons – making it impossible for each atom to satisfy the octet rule. • Others have some atoms with fewer or more than 8 valence electrons. • So we see that the “octet rule” is more of a general guideline that can be helpful in many circumstances, but is not always going to be satisfied.

32. Fume-hood Demo • Put on goggles, apron and gloves. • Under a fume hood, place a small piece of copper metal in an evaporating dish, make sure the fume hood is on, and then pour concentrated nitric acid (HNO3) over it. • Students make observations of the reaction below:

33. The Reaction • The balanced chemical equation for this reaction is: • NO2 is one of the pollutants in car exhaust, giving smog its reddish-brown color, and is very reactive and poisonous. • Try to write a structural formula for NO2.

34. The Resonance Structures • The most stable structures that can be drawn are: (Note: it is also possible to draw structures that place the unpaired electron on oxygen.) • Any molecule that contains an unpaired electron is called a free radical and tends to be very reactive(which explains the highly reactive nature of oxygen, O2). ↔ O N = O O = N O

35. Dimerization • When free radicals interact with each other, they share their unpaired electrons and create a dimer – a larger molecule constructed of two of the radicals. O = N O O = N O  + O N = O O N = O

36. Breaking the rules – too few • Boron trifluoride only has 6 electrons around boron: • Remember that fluorine (4.0) is much more electronegative than boron (2.0), so the bonds are more ionic in nature. (Boron is essentially emptying out its valence level.) F – B – F F

37. Breaking the rules – too few • Boron trifluoride can achieve 8 electrons around boron if it forms a coordinate covalent bond with another atom that has an unshared pair, such as the nitrogen of ammonia: H H F F N H  F – B + N H F – B F F H H

38. Breaking the rules – too many • Some elements, most notably phosphorus and sulfur, will form bonds to more atoms than expected and end up with more than an octet. PCl3 follows the octet rule, PCl5 does not. Draw them below:

39. Breaking the rules – too many • Sulfur hexafluoride is a very dense gas. What is its structural formula? • How many electrons surround the sulfur atom?

40. References: • http://www.green-planet-solar-energy.com/images/lewis-dot-structure-fluorin.gif • http://3.bp.blogspot.com/-pnWFaI3vfzM/TWZlR_l5JFI/AAAAAAAAABA/Z-7-i-nqkW4/s1600/vvvv.GIF • http://www.eou.edu/chemweb/molmodel/images/hclf.gif • http://www.eou.edu/chemweb/molmodel/images/h2ofc.gif • http://www.eou.edu/chemweb/molmodel/images/methstr.gif • http://1.bp.blogspot.com/-OlnUXIHKtDg/Teh5ekJUTjI/AAAAAAAAAD8/UIgoC2JK2PI/s1600/770px-Water-2D-flat.png • http://www.etsu.edu/physics/ignace/physics/lightning.jpg • http://www.bbc.co.uk/scotland/learning/bitesize/standard/chemistry/images/nitrogen_fixation.gif • http://imagine2050.newcomm.org/wp-content/uploads/2010/09/safetyfirst.PNG