1. Unit 13 - Bonding Chapter 12 Chemical Bonding Pages 398-439

2. Name the two types of compounds Ionic Compounds –Metal bonded to nonmetal –Cation bonded to anion Molecular compounds –Nonmetal bonded to Nonmetal What is a bond?

3. Chemical Bond A chemical bond is a force that holds groups of two or more atoms together and makes them function as a unit. There are three types of bonds: –Metallic –Ionic –Covalent

4. #1 Metallic Bonding Metals bond with each other through metallic bonding Metals have ____ ionization energy –So they _____ electrons easily The valences electrons lost are free to move throughout the entire metal and the result is a “sea of electrons” Sea of electrons lose low

5. #2 Ionic Bonding An ionic bond is a chemical bond formed by the electrostatic attraction between a positive ion (cation) and a negative ion (anion). –Opposites Attract!!

6. Molecule or not? A molecule is two or more atoms bonded together through the sharing of electrons. Is an ionic compound a molecule? No, ions are held together by attractive forces not by sharing of electrons. Ionic compounds form solid crystal structures called a crystal lattice (with repeating patterns)

7. Crystal Lattice Sodium atom losses an electron to form a + 1 ion Chlorine atom gains an electron to form a -1 ion They approach each other and become attracted to one another. The lattice continues to form as long as there is ions available.

8. Ionic Compound Formation 1. Alkali metals form ionic bonds with halogens 2. Alkaline earth metals form ionic bonds with halogens 3. Alkali metals form ionic bonds with nitrogen and oxygen 4. Alkaline earth metals form ionic bonds with nitrogen and oxygen KCl K 2 O or K 3 N CaO or Ca 3 N 2 MgCl 2

9. How are ionic compounds formed? To form an ionic compound, a transfer of one or more electrons must occur. One way to represent the formation of an ionic bond is by using electron configurations. Draw the electron config and orbital diagram for a sodium atom and a chlorine atom

10. How does an Ionic Bond Form between Sodium and Chlorine Sodium Na1s 2 2s 2 2p 6 3s 1 1s 2s 2p 3s Chlorine Cl1s 2 2s 2 3s 2 3p 5 1s 2s 2p 3s3p 3p 6 Na + Cl -

11. Electron Dot symbol Lewis Dot symbol Rules The letter symbols represent the nucleus and the core electrons of an atom Each dot represents one valence electron Dots are placed one at a time around an imaginary square surrounding the nucleus and core electrons. For one through four valence electrons, each “side of the square” gets a single electron. When more than four valence electrons exist the electrons start to double up and two dots are place on each side when necessary.

12. Electron Dot Structure for Sodium Chloride [Na] + [ Cl ] -

13. Cesium and Phosphorus Cs 3 P simpler way of showing Lewis picture No electron configurations Cs P P -3 Cs + + +

14. # 3 Covalent Bonding A covalent bond is chemical bond in which atoms share electrons. –The atoms can share One pair of electrons (2) Two pairs of electrons (4) Three pairs of electrons (6) Review: what happens to atoms in an ionic bond? –Transfer of electrons

15. Two types of covalent bonds Covalent bonds exist as two varieties: –Polar covalent –Nonpolar covalent This will be discussed later in the Unit. Just be aware

16. The Octet Rule Atoms tend to form bonds to achieve a noble gas electron configuration. Atoms that achieve this configuration have eight valence electrons. Exception: Hydrogen (H) Yippie!! I have 8 v.e.

17. Example: Covalent Bonding of two fluorine atoms

18. The Duet Rule Hydrogen is the only exception to the octet rule. Since hydrogen’s single electron is in the first energy level and the first energy level contains only a single s orbital, obtaining enough electrons to have a full octet in its energy level is impossible. Hydrogen tends to form bonds to achieve the electron configuration of helium (its nearest noble gas) which has only two electrons.

19. Covalent Story Two hydrogen atoms and an oxygen atom are sitting on a couch. The oxygen atom has 6 valence electrons, and the hydrogen atoms each have 1 valence electron. They are all unhappy because none of them satisfy the octet rule (or in the case of the hydrogen atoms, they don't satisfy the duet rule).

20. The oxygen atom then comes up with a brilliant idea. He proposes that each of the hydrogen atoms share an electron with him.

21. Electron Dot Structures Also known as a Lewis Dot Structures is a simple way of showing covalent bonds. Example: Chlorine Write the electron dot figure for two atoms of chlorine Both chlorine atoms achieved the octet rule by sharing one pair of electrons. The shared electrons can be replaced by a line (represents the covalent bond). Cl

22. Covalent bonds can also be represented by using orbital diagrams: 1s2s 2p 3s 3p Cl Shared electrons

23. Double Bonds Some atoms share two pairs of electrons so they may achieve the valence electron configuration of a noble gas. The atoms that have a tendency to form double bonds are: Sulfur (S),Carbon (C),Oxygen (O), and Silicon (Si)

24. Example: Oxygen Write the electron dot symbol for two atoms of oxygen: The oxygen atoms form a double bond by sharing two pairs of electrons to satisfy the octet rule. The shared electrons can be replaced by lines (the lines represents the covalent bonds). OO

25. Triple Bonds Some atoms share three pairs of electrons so they may achieve the valence electron configuration of a noble gas. The atoms that have a tendency to form triple bonds are: Carbon (C) andNitrogen (N)

26. Example: Nitrogen Write the electron dot symbol for two atoms of Nitrogen: The nitrogen atoms form a triple bond by sharing three pairs of electrons to satisfy the octet rule. The shared electrons can be replaced by lines (lines represent the covalent bonds). N N

27. Bonding and Non-bonding electrons Bonding electrons are electrons that are involved in bonds between atoms, they are the electrons that are shared. Non-bonding electrons are electrons that are not involved in bonding. Pairs of non-bonding electrons are often referred to as lone pairs. NN 3 bonding pairs = 6 bonding electrons 2 lone pairs = 4 non- bonding electrons

28. Writing Lewis Structures There are two key pieces of information about a molecule you need to know in order to draw a Lewis Dot structure: 1. Number of bonds 2. Number of non-bonding electrons

29. Rules for Lewis Structures 1. Draw the skeleton of the molecule. –Example NF 3 –The element that only has one atom usually goes in the middle. Exception: hydrogen –Connect the elements with single bonds NF F F

30. Rules for Lewis Structures 2. Count the number of valence electrons that should be in the molecule. Element # of atoms valence electrons total Fluorine Nitrogen NF F F 3 1 x 7 x 5 = 21 = 5 26 v.e.

31. Rules for Lewis Structures 3. Subtract the electrons already in the bonds (bonded electrons) of the skeleton from the total number of valence electrons. This difference equals the number of additional electrons needed. NF F F 26 v.e.- = 6 bonded electrons 6 =20e -

32. Rules for Lewis Structures 4. Add electron pairs to the molecule to complete octets around each atom, starting with the side atoms and ending with the central atom, until the correct number of electrons have been added (work from outside to inside). NF F F 20 electrons

33. 3 possible cases A. All atoms, including the central atom, have a complete octet of electrons. No further change is necessary. Sit back and admire your structure B. The central atom has less than the complete octet . Convert one or more pairs of electrons from the side atoms into double bonds until the central atom has an octet. Sit back and admire your structure. C. The central atom has more than the complete octet . This is necessary in order for the molecule to have the correct number of electrons. Such a molecule has an expanded octet. No further change is necessary. Sit back and admire your structure.

34. Page. 15 CHI 3 1.Skeleton 2.Total valence electrons 3.Subtract bonding electrons 4.Add remaining electron pairs outside to inside to fulfill octet rule. CI I I H C 1 x 4 = 4 H 1 x 1 = 1 I 3 x 7 = 21 26 v.e - 8= 18 e -

35. Page. 15 HCN 1.Skeleton 2.Total valence electrons 3.Subtract bonding electrons 4.Add remaining electron pairs outside to inside to fulfill octet rule. CHN H 1 x 1 = 1 C 1 x 4 = 4 N 1 x 5 = 5 10 v.e - 4= 6 e - Carbon is missing 4 electrons. Carbon can share 4 more electrons (2 pairs = 2 additional bonds) with nitrogen creating a triple bond

36. Review: Ionic Electron Dot Structure KCl K Cl - K +

37. Electron Dot Structures for Polyatomics Polyatomic ions are simply covalently bonded molecules that have either gained or lost a single electron or multiple electrons.

38. Steps for electron dot structures for polyatomic ions 1. from the ion formula, determine the number of valence electrons contributed from atoms in the ion. Get a total 2. Add electrons or subtract electrons from the total in step 1 as appropriate. Recall: a) to become a negative ion, atoms gain electrons b) to become a positive ion, atoms loss electrons 3. Use the total number of electrons from step two and follow the steps for writing electron dot structures for covalent compounds. Avoid placing the extra electrons around the central atom whenever possible. 4. Remember to put the entire electron dot structure in brackets and indicate the charge.

39. Example 1 – Draw the electron dot configuration for chlorate ClO 3 -1 Cl 1 x 7 = 7 O 3 x 6 = 18 25 v.e + 1e electron from the charge = 26 e - ClO O O 26 e - - 6 = 20e -