2. Chapter 8 Pages 238-273 Covalent Bonding

3. Learning Goals I can recognize the difference between an ionic, covalent and metallic compound. I can explain the difference between ionic, covalent and metallic bonding. I can correctly name ionic, metallic and covalent compounds.

4. Chemical Bonding Chemical Bond The force that holds 2 or more atoms together. Bonds form by the attraction between the positive nucleus of one atom and the negative electrons of another atom. Only the valence electrons are involved in chemical bonds. Chemical bonding results in a more stable situation for each atom in the bond.

5. Ionic, Metallic & Covalent Compounds 3 types of compounds – Ionic: NaCl, FeCl 3, CuCl 2 – Metallic: Na, Fe, Cu, brass (alloy of Cu and Zn) – Covalent:H 2 0, Dextrose- C 6 H 12 O 6 ; methane-CH 4 ; NH 3

6. Covalent Bonding The chemical bond that results from sharing electrons is a covalent bond. A molecule is formed when two or more atoms bond. We will use Lewis structures to represent covalent bonding. Example: H 2 O Basic bonding video

7. Ionic, Metallic and Covalent Compounds Are the following compounds ionic, metallic or covalent? – CaBr 2 – Ca – Bronze- alloy of Cu, Zn and Sn – HBr – LiF – CO 2 – LiOH – Zn – Zn(NO 3 ) 2

8. Ionic vs. Metallic vs. Covalent Bonding Explain the difference in bonding between the 3 different types of compounds. Ionic bonding- Metallic bonding- Covalent bonding-

9. Ionic vs. Metallic vs. Covalent Bonding Ionic bonding- – Valence electrons are transferred forming a cation (loses the e - ) and an anion (gains the e - ). Generally formed between a metal and a nonmetal. Metallic bonding- – Valence electrons are delocalized and are attracted to all of the metal cations. No electrons are lost or gained. Formed within all elemental metals and alloys. Covalent bonding- – Valence electrons are shared between 2 atoms forming a molecule. Generally formed between 2 nonmetals.

10. Naming Covalent Compounds First element in the formula is always name first. The second element in the formula is named using its root name and adding the suffix –ide. Prefixes are used to indicate the number of atoms of each element that are present in the compound. NEVER use “mono” for the first element

11. Naming Covalent Compounds

12. CO 2 NH 3 N 2 P 2 O 5 H 2 H 2 O NF 3 CCl 4

13. Covalent Bonding There are 7 diatomic molecules: H 2, N 2, O 2, F 2, Cl 2, Br 2, I 2 The mnemonic device used to remember this is “Have No Fear Of Ice Cold Beer Diatomic molecules form when 2 atoms of the same element share electrons to form a covalent bond. These 2-atom molecules are more stable than the individual atoms. By sharing electrons each atom achieves the stable octet of (8) valence electrons.

14. Covalent Bonding The figure shows two hydrogen atoms forming a hydrogen molecule with a single covalent bond, resulting in a more stable H 2 molecule.

15. Single Covalent Bonds Sigma bonds are single covalent bonds. Sigma bonds Sigma bonds occur when the pair of shared electrons is in an area centered between the two atoms.

16. Covalent Bonding Double bonds form when two pairs of electrons are shared between two atoms. Triple bonds form when three pairs of electrons are shared between two atoms.

17. Multiple Covalent Bonds A multiple covalent bond consists of one sigma bond and at least one pi bond. The pi bond is formed when parallel orbitals overlap and share electrons.pi bond Double bonds contain 1 sigma bond and 1 pi bond. Triple bonds contain 1 sigma bond and 2 pi bonds.

18. Lewis Structures In a Lewis structure dots or a line are used to symbolize a single covalent bond.Lewis structure The halogens—the group 17 elements—have 7 valence electrons and form single covalent bonds with atoms of other non-metals.

19. Lewis Structures, cont’d. Atoms in group 16 can share two electrons and form two covalent bonds. Water is formed from one oxygen with two hydrogen atoms covalently bonded to it.

20. Lewis Structures, cont’d. Atoms in group 15 form three single covalent bonds, such as in ammonia.

21. Lewis Structures, cont’d. Atoms of group 14 elements form four single covalent bonds, such as in methane.

22. Lewis Structures, Cont’d Group 1 metals can only form 1 bond, and thus disobey the octet rule. Group 2 metals can only form 2 bonds and thus disobey the octet rule. Group 3 metals can only form 3 bonds and thus disobey the octet rule.

23. Drawing Lewis Structures – Predict the location of certain atoms. – Carbon will always be in the middle if it is present – Hydrogen will always be on the outside if it is present – Otherwise, the most electropositive atom goes in the middle. – Count the number of valence electrons on each atom and the total number and draw them in place. – Draw a single bond between each atom – Check to see if each has the octet or however many it needs. – If you have 7-7, add a double bond, if you have 6-6, add a triple bond.

24. Resonance Structures Resonance is a condition that occurs when more than one valid Lewis structure can be written for a molecule or ion. Resonance This figure shows three correct ways to draw the structure for (NO 3 ) -1.

25. Exceptions to the Octet Rule Some molecules do not obey the octet rule. A small group of molecules might have an odd number of valence electrons. NO 2 has five valence electrons from nitrogen and 12 from oxygen and cannot form an exact number of electron pairs.

26. Exceptions to the Octet Rule A third group of compounds has central atoms with more than eight valence electrons, called an expanded octet. Elements in period 3 or higher have a d-orbital and can form more than four covalent bonds.

27. Covalent Bonding We can determine whether a covalent bond is polar or nonpolar based upon the ELECTRONEGATIVITY difference between the bonded atoms and the shape the bonded molecule has. What is electronegativity? – The ability of an atom to attract electrons to itself in a chemical bond

28. Electronegativity

29. Electonegativity & Bond Character A chemical bond is never 100% ionic or 100% covalent. The character of a chemical bond (ionic vs. covalent) depends on the ELECTRONEGATIVITY DIFFERENCE between the bonded atoms.

30. Electonegativity & Bond Character Example: H 2 O Electronegativity of O = 3.44 Electronegativity of H = 2.20 EN Difference 3.44- 2.20 = 1.24 According to table 8.7 this makes H 2 O a polar covalent compound.

31. Electonegativity & Bond Character Example: NaCl Electronegativity of Cl = 3.16 Electronegativity of Na = 0.93 EN Difference 3.16 – 0.93 = 2.23 According to table 8.7 this makes NaCl an ionic compound.

32. Electonegativity & Bond Character Example: CH 4 Electronegativity of C = 2.55 Electronegativity of H = 2.20 EN Difference 2.55- 2.20 = 0.35 According to table 8.7 this makes CH 4 a mostly nonpolar covalent compound.

33. Your Turn Use the electronegativity values and table 8.7 to determine if the following compounds are either (in textbook): – mostly ionic – mostly polar covalent or – mostly nonpolar covalent – Then provide the NAME of each compound using our naming rules.

34. Your Turn KBr CaO CO N 2 H 2 S MgCl 2 NO 2 LiF HBr CS 2 Apply the concept of “like dissolves like”. Only compounds “like” water will dissolve in water (the ionic and polar compounds). Put a star next to the above compounds which you predict will dissolve in water