Covalent bonds and Lewis Structure

All you need to know about covalent bonds and lewis structure.

1. COVALENT BONDS & LEWIS STRUCTURE Prepared by: Mrs. Eden C. Sanchez

2. Learning Objectives Illustrate the formation of covalent bonds in terms of electron sharing Apply the octet rule in forming covalent compounds Define electronegativity Describe the electronegativity trends in the periodic table 2

3. Learning Objectives Draw Lewis structure of covalent compounds Identify lone pairs and bond pairs; Draw the resonance structures of covalent compounds Determine the polarity of a bond based on the electronegativities of the bonding atoms 3

4. Learning Objectives Determine whether a bond is ionic, polar covalent, or covalent based on the differences in electronegativities of the bonding atoms 4

5. Keywords Lewis structure Covalent bond Lone pair Bond pair Single bond Double bond 5

6. Keywords Triple bond Nonpolar covalent bond Polar covalent bond Electronegativity Percent ionic character Resonance 6

7. Keywords Incomplete octet n. Expanded octet 7

8. FORMATION OF THE COVALENT BOND Gilbert Lewis suggested that the chemical bond is formed by sharing of electrons in atoms. Example: The two electrons are shared equally between the two atoms forming a covalent bond. 8

9. FORMATION OF THE COVALENT BOND The electrons are attracted to the nuclei of both atoms keeping the atoms together to form a molecule . formation of the covalent bond for the F2 molecule 9

10. FORMATION OF THE COVALENT BOND The representation of the covalent compound is called the Lewis structure . In the Lewis structure , shared electrons that form a bond is represented by a line or a pair of dots ; lone pairs are represented by dots above the atom . 10

11. FORMATION OF THE COVALENT BOND From the Lewis structure of F 2 , how many electrons are around each fluorine atom in F 2 ? How many bond pairs are there in the F 2 molecule? 11

12. FORMATION OF THE COVALENT BOND How many lone pairs are there in the F 2 molecule? I llustrate the formation of the covalent bond in Cl 2 . How many bond pairs are there? How many lone pairs ? 12

13. FORMATION OF THE COVALENT BOND Illustrate the formation of the covalent bond in HCl . 13

14. Exercises: Draw the Lewis structure for H 2 O, CH 4 (methane), and for NH 3 . Which of the three molecules has the largest number of bond pairs (covalent bonds)? Draw the Lewis structure for carbon dioxide, CO 2 . 14

15. The examples of CO 2 and N 2 show that there are different types of covalent bonds that are formed. Single bonds are formed when two atoms are held together by one pair of electrons. Multiple bonds can be formed also 15

16. A double bond is from the sharing of two pairs of electrons such as in the case of O and C in CO 2 . A triple bond exists in N 2 where the two N atoms are held by three pairs of electrons. 16

17. Types of Covalent Bond Experimental evidence has shown that electrons are not equally shared between H and F; the electrons spend more time near F rather than H. Therefore the electron density is shifted more towards F rather than H. 17

18. This leaves the F end of the molecule partially negative, δ - , and the H end of the molecule partially positive, δ+, such bond is referred to as a polar covalent bond. 18

19. The polar covalent bond is somewhere between a purely covalent (nonpolar) bond and an ionic bond (where there is almost complete transfer of electrons). 19

20. Electronegativity property that distinguishes the polarity of bonds the tendency of an atom in a chemical bond to attract electrons toward itself is a theoretical concept and devised as a relative scale. That is, it can be estimated relative to, or in comparison to, other elements in chemical bonds. 20

21. Electronegativity Linus Pauling developed a relative scale of electronegativities which is widely used in General Chemistry In contrast, ionization energies and electron affinities are physically measurable properties of elements. 21

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23. The difference in the electronegativity values ( EN ) of two bonded atoms determines the percent ionic character of the bond. If the bond is between two identical elements, like F—F , then the bond is purely covalent with 0 percent ionic character. The difference in electronegativity is 0. 23

24. For the molecule H— Cl , the difference in electronegativity is 0.9 showing that the bond is a polar covalent bond . A 50% ionic character corresponds to EN=1.7. While there is no bond that is 100% ionic , an electronegativity difference of 2.0 or greater is usually classified to be predominantly ionic. 24

25. When EN ≥ 2.0, the bond is predominantly ionic. 25

26. Exercises: I. Classify the following bonds as ionic , polar covalent , or covalent . The C-C bond in H 3 CCH 3 The K-I bond in KI The C-F bond in CF 4 The N-H bond in NH 3 26

27. Exercises: II. Arrange the following bonds according to increasing bond polarity: Cs to F, Cl to Cl , Br to Cl , Si to C . 27

28. Guidelines in Writing the Lewis Structure of Covalent Molecules Draw a skeletal structure of the molecule putting bonded atoms next to each other. In general , the least electronegative atom occupies the central position. H and F usually occupy terminal (end) positions . 28

29. Guidelines in Writing the Lewis Structure of Covalent Molecules Count the total number of valence electrons from all the atoms in the structure. Add electrons corresponding to the charge for negative ions; subtract electrons corresponding to the charge for positive ions. 29

30. Guidelines in Writing the Lewis Structure of Covalent Molecules Distribute the valence electrons to the non-central atoms such that these atoms fulfill the octet rule. Remaining electrons are assigned to the central atom. Remember that bonds are equivalent to 2 electrons. If the valence electrons are not enough, multiple bonds may be formed. 30

31. Exercises Write the Lewis structure for NCl 3 . Write the Lewis structure of OCS. C is the central atom . Write the Lewis structure of CN –. 31

32. Exercises Write the Lewis structure of the following molecules: Ethylene , C 2 H 4 Acetylene , C 2 H 2 Carbon tetrachloride, CCl 4 COBr 2 (for the skeletal structure, C is bonded to O and Br atoms) 32

33. Lewis Structure & Resonance Write the Lewis structure for the ozone molecule, O 3 . To resolve this discrepancy, we represent the ozone molecule using the two structures presented as: 33

34. Lewis Structure & Resonance T he above structures is called a resonance structure . The double sided arrow shows that the structures are resonance structures. A resonance structure is one of two or more Lewis structures for a molecule that cannot be represented accurately by only one Lewis structure. 34

35. Exercise: Draw the resonance structures for the carbonate ion , CO 3 2- . 35

36. EXCEPTIONS TO THE OCTET RULE: The octet rule works best for second-period elements. Hence there are many exceptions. They fall into three categories: Incomplete octet Odd number of electrons Expanded Octet 36

37. EXCEPTIONS TO THE OCTET RULE: Incomplete octet An example of a molecule with incomplete octet is BeH 2 , beryllium hydride . H – Be – H There are only 4 electrons around Be and not 8. Boron and aluminum also form molecules with incomplete octets. 37

38. EXCEPTIONS TO THE OCTET RULE: Draw the Lewis structure of aluminum triiodide , AlI 3 , showing the incomplete octet . 38

39. EXCEPTIONS TO THE OCTET RULE: Molecules with Odd Number of Electrons Examples are nitric oxide, NO, and dinitrogen dioxide, N 2 O. 39

40. EXCEPTIONS TO THE OCTET RULE: The odd numbered molecules are sometimes referred to as radicals . They are generally highly reactive . 40

41. EXCEPTIONS TO THE OCTET RULE: Expanded Octets Atoms belonging to the second period cannot have more than eight valence electrons around the central atom because they only have the 2s and 2p subshells. This is different for atoms of elements in the 3rd period and beyond. 41

42. EXCEPTIONS TO THE OCTET RULE: These elements have 3d orbitals that can participate in the bonding. Hence they can have more than eight valence electrons around the central atom. An example is SF 6 , sulfur hexafluoride and phosphorus pentafluoride , PF 5 . 42

43. EXCEPTIONS TO THE OCTET RULE: 43

44. EXCEPTIONS TO THE OCTET RULE: Another example is BrF 5 44

45. NAMING COVALENT COMPOUNDS: (A REVIEW) For binary compounds, state the name of the first element. The name of the second element ends in –ide . HF Hydrogen fluoride HI Hydrogen iodide SiC Silicon carbide 45

46. NAMING COVALENT COMPOUNDS: (A REVIEW) Prefixes are used to denote the number of atoms in the formula. The prefix “mono” usually omitted for the first element in the formula . CO carbon monoxide CO 2 carbon dioxide NO 2 nitrogen dioxide 46

47. NAMING COVALENT COMPOUNDS: (A REVIEW) N 2 O 4 dinitrogen tetroxide CCl 4 carbon tetrachloride SF 6 sulfur hexafluoride 47

48. Seatwork Look for at least 2 examples of covalent compounds that can be found in nature or used in everyday life. Include the following information: uses of the covalent compound c hemical formula and chemical name of the covalent compound s tructure of the compound 48

49. 49 Thank You! Any questions ?

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