2. Chemical Bonding

3. Learning Objectives To understand what covalent bonding is. To predict when covalent bonding will occur. To use chemical formulas to build a visual model.

4. Covalent bonding takes place between non- metals atoms only. Atoms attain a full outer shell of electrons - ‘noble gas’ structure - by sharing electrons. Atoms bonded in this way form molecules

5. Chlorine Cl 2

6. Chlorine has atomic number 17 It has 17 protons in its nucleusIt also has 17 electrons 2 in the inner shell8 in the next shell7 in the outer shell CHLORINE We need 2 chlorine atoms These atoms can then share electrons to complete their outer shells A chlorine molecule is formed Also written Cl-Cl or Cl 2

7. Methane CH 4

8. Carbon has atomic number 6 It therefore has 6 electrons 2 in the inner shell 4 in the outer shell Hydrogen has just one electron And we need 4 hydrogens Now both carbon and all 4 hydrogens have complete outer shells 4 covalent bonds have been made Covalent bonds can be shown like this. Or like this

9. Oxygen O 2

10. Oxygen has atomic number 8 Therefore it has 8 electrons 2 in the inner shell 6 in the outer shell We need a second oxygen To complete the outer shell, they each need to share 2 electrons, so a total 4 electrons are shared The covalent bonds can be shown like this. Or like this O=O

11. Things to notice about covalent bonding Carbon (group 4) always forms 4 bonds Nitrogen (group 5) often makes 3 bonds Oxygen (group 6) often makes double bonds All of these arrangements allow the atoms to have a full outer shell of 8 (or 2 for Hydrogen). Double bonds count as two bonds, and satisfy the need to share two electrons for BOTH the atoms eg.

12. The most stable electronic pattern is what the ‘noble gases’ have: a full outer shell Atoms combine with other atoms to achieve this. SHELL NUMBER ONE  SHELL NUMBER TWO  SHELL NUMBER THREE 

13. The most stable electronic pattern is what the ‘noble gases’ have: a full outer shell Atoms combine with other atoms to achieve this. SHELL NUMBER TWO  SHELL NUMBER THREE 

14. Group 4 needs to share 4 electrons to get to a full shell of 8 (so makes 4 bonds). Group 5 needs 3 electrons (makes 3 bonds). Group 6 needs 2 electrons (makes 2 bonds). Group 7 only needs 1 electron. HYDROGEN ONLY NEEDS 1 ELECTRON (ONE BOND) TO BE LIKE HELIUM 

15. NameFormu la Structural Formula Lewis Diagram MethaneCH 4 Carbon DioxideCO 2 Water Ammonia Hydrogen cyanide Formaldehy de

16. NameFormu la Structural Formula Dot & Cross Diagram MethaneCH 4 Carbon DioxideCO 2 WaterH2OH2O AmmoniaNH 3 Hydrogen cyanide HCN Formaldehy de H 2 CO

17. Covalent Bonding with Lewis Dot

18. Covalent Bond When nonmetallic elements react with other nonmetallic elements, they share electrons in order to obtain eight valence electrons. Each fluorine atom has seven valence electrons. They each require one more electron to satisfy the Octet Rule.

19. The left fluorine atom now has a total of eight electrons and the right fluorine atom now has a total of eight electrons around it. When nonmetallic elements react with other nonmetallic elements, they share electrons in order to obtain eight valence electrons.

20. The two electrons that form the covalent bond are often Represented by a single line. The F 2 molecule can be represented using a line and dots to show the bonding pair and the six lone pairs, respectively. This is called a Lewis dot structure.

21. Multiple Covalent Bond Some atoms have to share more than one electron in order to satisfy the Octet Rule.

22. Each oxygen atom has six valence electrons. They each require two more electrons to satisfy the Octet Rule.

23. The left oxygen atom now has a total of eight electrons around it. The right oxygen atom now has a total of eight electrons around it.

24. The four electrons shared by the oxygen atoms form a double bond. The double bond is represented by two single lines. Each line in the Lewis dot structure represents two electrons

25. The element hydrogen is an exception to the Octet Rule. It only needs two electrons, rather than eight, to be stable. The hydrogen atom has one valence electron. It requires one more electron to be stable. The fluorine atom has seven valence electrons. It requires one more to satisfy the Octet rule.

26. The hydrogen atom now has a total of two electrons around it and is stable. The fluorine atom now has a total of eight electrons around it and is stable.

27. The Lewis dot structure of the HF molecule shows a line and 6 dots to represent the bonding pair and the 3 lone pairs of electrons, respectively.

28. Rules for writing Lewis Dot structures Rule 1 Add together the number of valence electrons for each atom in the molecule. For example, CF 4 Carbon has four valence electrons and each fluorine has seven valence electrons = 4 + 4(7) = 32

29. Rule 2 Write out the elements of the molecule so that the least electronegative elements is in the center surrounded by the other elements. For example, CF 4

30. Rule 3 Place a covalent bond between the central atom and the outside atoms. Remember each covalent bond contains two electrons.

31. The four covalent bonds use eight of the 32 valence electrons in CF 4 Rule 4 There are 24 valence electrons remaining. Add electrons to the outer atoms as lose pairs to satisfy the Octet Rule. This uses 24 electrons. There Are no electrons left, so this is The Lewis dot structure for CF 4

32. Rule 5 for example, NH 3 First apply Rules 1-4 to the molecule Rule 1: Count the valence electrons Rule 2: Place the least electronegative element at the centre, except for H which is always an outer atom Rule 3: Add covalent bonds between the centre atom and the outer atoms Rule 4: Add lone pairs to the outer atoms Rule 5: Add lone pairs to the centre atom

33. Rule 1 Nitrogen has 5 valence electrons and each hydrogen has 1 valence electron The total number of valence electrons = 5 + 3 (1) = 8 Rule 2 Hydrogen is always an outer atom and is never at the centre of a molecule

34. Rule 3 Add the bonding electrons. This uses 6 of the 8 valence electrons. Rule 4 The 2 remaining valence electrons are not added to the outer atoms, because each H has its maximum of 2 valence electrons.

35. Rule 3 Add the bonding electrons. This uses 6 of the 8 valence electrons. Rule 4 The 2 remaining valence electrons are not added to the outer atoms, because each H has its maximum of 2 valence electrons.

36. Rule 5 Place the remaining 2 Valence electrons on the central nitrogen atom Rule 6 Check all atoms in the molecule to ensure that each has 8 electrons(2 for hydrogen). If an atom has fewer than 8 electrons, create double or triple bonds. (Note: Double bonds only exist between C,N,O and S atoms) This is the Lewis structure For NH 3

37. Apply rule 6 to the following; CH 4, CF 4, Hydrogen : 1 bond = 2 electrons (stable) Carbon : 4 bonds = 8 electrons (stable) Fluorine : 1 bond + 3 lone pairs = 2 + 3 (2) = 8 electrons (stable) Carbon : 4 bonds = 8 electrons (stable)

38. Example; CH 2 O Apply Rules 1-5 to the molecule Rule 1: Count the valency electrons Rule 2: Place the least electronegative element at the centre, except for H, which is always an outer atom Rule 3: Add covalent bonds between the centre and the outer atoms Rule 4: Add lone pairs to the outer atoms Rule 5: Add lone pairs to the centre atom

39. Rule 1 Carbon has 4 valence electrons, each hydrogen has 1 valence electron, and oxygen has 6 valence electrons. Total number of valence electrons : 4 + 2(1) + 6 = 12 Rule 2 Carbon is at the centre of the molecule because it is less electronegative than oxygen. Hydrogen is always an outer atom and is never at the centre of the molecule.

40. Rule 3 Add the bonding electrons. This uses 6 of the 12 valence electrons Rule 4 Add the remaining 6 lectrons to the outer atom. Hydrogen does not need any more electrons, but Oxygen needs 6 to complete its octet.

41. Rule 6 Oxygen shares one of its lone pairs with C and O and give the desired 8 electron total Rule 5 There are no valence electrons left to add to the centre This is the Lewis dot Structure for CH 2 O

42. Exceptions to the Octet Rule The Octet Rule applies to Groups IVA through VIIA in the second row of the Periodic Table, but there are exceptions to the rule among some other elements. The following two cases are an example Example BF 3 Rule 1 Boron has 3 valence electrons and each Fluorine has 7 valence electrons Total number of electrons = 3 + 3 (7) = 24

43. Rule 2 Boron is at the centre of the molecule because it is less electronegative than fluorine Rule 3 Add the bonding electrons. This uses 6 of the 24 valence electrons

44. Rule 4 Add the remaining electrons to the outer atoms. Each Fluorine has the required 8 electrons Rule 5 This uses the remaining electrons leaving none to add to the Boron central atom

45. Rule 6 Check the number of electrons around each atom. Each Fluorine atom has 8 electrons, but the Boron Atom has only 6. This is an exception to the Octet Rule. A B=F bond is not an option, because double bonds exist only between C,N,O, and S atoms This is the Lewis dot structure BF 3

46. M ETALLIC B ONDING 45 However, metals behave differently. Metallic bonding is similar to both covalent and ionic bonding The valence (outermost) electrons are loosely held by the metal ions, so much so that they move away from the atom to form a positively charged ION. The electrons are free to move from one positively charged ION to the next (i.e. They are DELOCALISED) and are shared (just like in covalent bonding among the various metallic positively charged ions The number of electrons = the number of protons. The metal is therefore electrically NEUTRAL Source: www.daviddarling.info/images/metallic_bond.jpg

47. C OMPARE AND CONTRAST TYPES OF BONDING 46 Metallic and ionic bonding involve electrostatic attractions between positive and negatively charged particles. Metallic bonding shares electrons among the ions in a similar manner to how electrons are shared in covalent bonding. Covalent bonding shares electrons rather than having electrostatic charges. Ionic bonding will form compounds whereas covalent bonding can form a compound or element and metallic bonding is strictly found in elements SimilaritiesDifferences