Introduction to Bonding

1. Introduction to Bonding Topic #13 Essential Question: What are all the differences and similarities between covalent bonds and ionic bonds?

2. What is a bond? • A bond can be thought of as a force that holds groups of two or more atoms together and makes them function as a unit • Example : water O Bonds require energy to break and release energy when made H H

3. Types of bonds • Ionic bonds - typically formed between metals and nonmetals • Covalent bonds - typically formed between nonmetals • Metallic bonds - formed between metals

4. Ionic Bonds • Ionic bonding results from the transfer of electrons. Then, the opposite charges attract each other. • Ionic bonds are strong

5. Ionic Bonds • Na and Cl • Na is a metal and likes to lose one electron and form a +1 ion. • Cl is a nonmetal and likes to gain one electron and form a -1 ion. • the final ionic compounds is NaCl + Na+ Cl- NaCl The electrostatic interaction keeps them together!

6. Ionic Bonds • They do this to achieve an octet!

7. Covalent Bonds • Covalent Bonds • exist between nonmetals bonded together • form when atoms of nonmetals share electrons • electrons can be shared equally or unequally • The unequal sharing results in polar molecules

8. Metallic Bonds • Metallic bonds exist between metals • Occur when two metals, usually the same metal, are bonded together • “sea of electrons” • “delocalized electrons” • ttp://www.youtube.com/watch?v=XAnTCYZPJsE&feature=bf_next&list=PLBFE28832E577A62B

9. Regents Chemistry • Electronegativity

10. How can we tell really tell which type of bond we have? • Electronegativity – is the relative ability of an atom in a molecule to attract shared electrons to itself • This tells us what type of bond we have; • Covalent, polar covalent or ionic • Electronegativity values are determined by measuring the polarities of bonds between various elements to determine a specific value for each element

11. Electronegativity • Electronegativity values for each element are obtained by using the Periodic Table • In fact, there is a general trend in electronegativity we observe in the Periodic Table • Electronegativity values increase across and up the Periodic Table • See table on pg. 332

12. Electronegativity • We take the difference between the electronegativity values to determine exactly what type of bond exists, in essence the polarity of the bond See table 12.1

13. Determining Bond Polarity • If the difference between the electronegativity values is: • 0.0 – 0.5: covalent bond (equal sharing) • 0.6 – 1.6: polar covalent bond (unequal sharing) • 1.7 – up: ionic bond (transferring electrons)

14. Examples • Use your Reference Tables to determine the difference in electronegativity values and the type of bond for each of the following: • H-H • H-Cl • H-O • H-S • H-F • NaCl • O2 • KBr Worksheet

15. Regents Chemistry • Intro to valence electrons

16. Electrons in an atom • Electrons surround the nucleus of an atom in specific energy levels or shells • Each level can hold only a certain amount of electrons • It is an atoms ability to the lose, gain or share electrons from its outer shell that determine its reactivity

17. The outer shell • The outer shell in an atom contains the valence electrons • Valence electrons can be lost, gained or shared to have eight electrons in the outer shell • Each group on the table tells the number of valence electrons

18. Periodic Table • Groups 1, 2, 13, 14, 15, 16, 17, 18 have 1,2,3,4,5,6,7,8 valence electrons, respectively • We will not consider the transition metals • See periodic table

19. Sharing to reach the Octet Rule • The octet rule states that an atom cannot have more than 8 electrons in its outer shell • Valence electrons are lost, gained or shared with other atoms to attain 8 electrons in the outer shell • Eight valence electrons means a filled and happy shell - like the noble gases

20. Nonmetals share • Nonmetals share electrons to reach eight valence electrons • Single, double and triple bonds can be formed by sharing electrons

21. Metals + non-metals = lose/gain e- • metals and nonmetals interact by losing and gaining electrons to reach 8 electrons (filled outer shell) • The oxidation states on the periodic table represent this desire to move electrons • ex: K+ want to lose 1 electron to reach noble gas configuration of eight electrons

22. Lewis structures: your assignment • The reading and problems focus on drawing Lewis structures • Lewis structures are a means to represent bond formation between atoms • Covalent bonded compounds have different Lewis structures than Ionic bonded compounds

23. Example of a Lewis Structure C CH4 Covalent bonds H H H C H H

24. Regents Chemistry • Lewis Structures

25. Lewis Structures • The Lewis Structure is a representation of a molecule that shows how the valence electrons are arranged among the atoms in a molecule • We used dots around the elemental symbol to represent the valence electrons C

26. Single Lewis Structure - Practice • Draw four lone electrons first (if necessary) them pair them up • Draw Lewis Structures for the following atoms Br Al Na Be

27. Lewis Structures for Ionic Compounds • For Lewis Structures of ionic bonds the atoms are not joined but draw next to each other example: KBr - K+ [ Br ] Bromine gains an electron to achieve the noble gas configuration of Krypton Potassium loses an electron to achieve the noble gas configuration of Argon

28. Lewis Structures – Covalent Bonds • Hydrogen forms stable molecules when it shares two electrons • Two electrons fill Hydrogen’s valence shell • Helium does not form bonds because its valence shell is already filled; it is a noble gas • Second row non-metals Carbon through Fluorine from stable molecules when surrounded by eight electrons – the Octet Rule

29. Lewis Structures – Covalent Bonds • Valence electrons in covalent bonds can either be bonding pairs, if involved directly in the bond or lone pairs if not involved in the bond

30. Writing Lewis Structures - Rules • Obtain the total sum of the valence electrons from all of the atoms • Use one pair of electrons to form a bond between each pair of bound atoms. For convenience, a line (instead of a pair of dots) can be used to indicate each pair of bonding electrons • Arrange the electrons to satisfy the duet rule for hydrogen and the octet rule for second row non metals

31. Lewis Structures – Covalent Bonds • Examples Step 1) 8 total valence e- total Step 2) Draw one pair of electrons per bond 8-6 = 2 left Step 3) Arrange the remaining electrons according to octet rule PH3 H l H– P –H •• H H P H

32. Lewis Structures – Covalent Bond Practice Examples .. H:Br: ·· HBr CF4 Worksheet

33. Regents Chemistry • Ionic Lewis Structures • Multiple bonds in Lewis Structures • Polyatomic ion Lewis Structures and Resonance

34. Lewis Structures for Ionic Compounds • For Lewis Structures of ionic bonds the atoms are not joined but draw next to each other example: KBr - K+ [ Br ] Bromine gains an electron to achieve the noble gas configuration of Krypton Potassium loses an electron to achieve the noble gas configuration of Argon

35. Examples of Ionic Lewis Structures • Draw Lewis Structures for the following: NaCl LiBr KI

36. Multiple Bonds and Lewis Structures…review first • We have seen how to draw Lewis Structures for molecules with single bonds • For example • NH3 8 total valence e- 3 bonds x 2e- = 6 bonding 2 e- left over • Sum the total • valence e- • Subtract number • of bonding e- • Place remaining • valence e- H N H H

37. Multiple Bonds • Between atoms of the same element • Example • Oxygen • O O Also a Lewis Structure O = O Just O = O is called a structural model

38. Example of Multiple Bonds Nitrogen N N N N We now meet the octet rule!

39. Multiple Bonds • Between atoms of different elements • CO2 O C O O = C = O We must use double bonds to meet the octet rule!

40. Lewis Structures for Polyatomic Ions and Resonance Structures • Read pg. 344 (bottom) to 349 and answer questions a-g in example 12.4 (pg. 347) and a-i in the Self Check exercise 12.4 (pg. 348)