1. Bond Types

2. Let’s Look At 3 Types of Bonds Ionic Polar Covalent (Molecular) Nonpolar Covalent (Molecular)

3. Remember Electronegativity? The tendency for an atom to attract electrons to itself in a bond –The higher the value, the better it is at attracting electrons. The difference in the electronegativity values determines what type of bond will be formed.

4. Electronegativity Electronegativity is a scale from 0.7 (Cs) to 4.0 (F). Electronegativity generally increases across a period and decreases down a group.

5. Electronegativity Values (these will be given to you) Why don’t the noble gases have a value? They don’t attract electrons!

6. Ionic Bonds If the electronegativity difference is greater than 1.7, one atom will pull the electron completely away from other atom. The electrons are NOT shared. Ionic BondAn Ionic Bond is formed as + and – attract. Electronegativity of Na is 0.9; Cl is 3.0. 3.0 – 0.9 = 2.1; difference >1.7, so…Ionic!

7. Polar Covalent Bonds Covalent bonds share electrons The shared pairs are pulled, similar to a tug-of-war, between the nuclei of the atoms sharing the electrons. If the electronegativity difference is between 0.3 – 1.7, one side of the bond becomes slightly more negative and the other side becomes slightly more positive. Polar Covalent BondThis is a Polar Covalent Bond.

8. Polar Covalent Bonds The electronegativity of H is 2.1; Cl is 3.0. 3.0 – 2.1 = 0.9; difference is b/w 0.3 – 1.7, so…Polar Covalent! ClH ∂+ ∂- Slightly

9. Polar Covalent Bonds The electronegativity of O is 3.5; H is 2.1. 3.5 – 2.1 = 1.4 Difference is b/w 0.3– 1.7, so… Polar Covalent!

10. Nonpolar Covalent Bonds Nonpolar CovalentWhen the atoms have equal pull, causing the electrons to be equally shared, the bond is Nonpolar Covalent. Neither side of the bond is even slightly positive or negative. The electronegativity difference is b/w 0.0 – 0.3. This is the type of bond that occurs between 2 atoms of the same element. (H 2, O 2, Cl 2, etc.)

11. Nonpolar Covalent Bonds The electronegativity of H is 2.1. 2.1 – 2.1 = 0 The difference is b/w 0.0 – 0.3, so…Nonpolar!

12. 3 Different Types of Bonds

13. Lewis Dot Structures Ionic & Metallic Bonding And…

14. Valence Electrons  Electrons in the outer energy level.  Determine chemical and physical properties of an element  The group number of the representative elements is the same as the number of valence electrons.  All of the elements within a given group will have the same number of valence electrons.

15. For example:  Be is in Group 2A.  There are 2 electrons in the outermost energy level.  Be has an e­ configuration of 1s² 2s²  How many valence electrons will F have? Valence Electrons 7

16. Lewis Dot Diagrams and the Octet Rule Noble gases have a FULL valence shell of 8 electrons (n s 2 n p 6 ). (Helium has a full valence shell with only 2 valence electrons.) Through bonding, other atoms “seek” a full shell of eight electrons. This is called the OCTET RULE. Noble gases are unreactive (inert) because they already have a full shell! Valence Electrons

17. Lewis Dot Diagrams and the Octet Rule A visual representation of where the bonding electrons are in an atom The VALENCE electrons are shown as dots around the symbol for the element. This is called an electron dot diagram or a Lewis dot structure. Electron Dot Diagrams

18. HOW TO MAKE A DOT DIAGRAM  Write the symbol for the element.  Decide how many valence electrons the element has.  Using dots, place one dot per electron on each side of an imaginary box around the symbol.  You must place one dot on each side of the box before doubling up. Otherwise, the order doesn’t matter. Otherwise, the order doesn’t matter. N 5 N

19. Electron Dot Diagrams Lithium has only 1 valence electron, so we only place one dot on our diagram. Lithium Li 1 valence electron!

20. Electron Dot Diagrams Beryllium has 2 valence electrons, so we place two dots on our diagram. Beryllium BeBe

22. Ionic Bonding Ionic Bonding  An atom is always trying to get a full outer energy level of eight electrons—Octet Rule.  Atoms can gain, lose, or share valence (outer) electrons to complete their outer shell.  When atoms get their full shells by completely giving or taking electrons from other atoms, an Ionic Bond is formed.

23. Ionic Bonding  Involves A metal and a nonmetal A metal and a nonmetal A positively charged ion ( the metal) A positively charged ion ( the metal) A negatively charged ion (the nonmetal) A negatively charged ion (the nonmetal)  An electrostatic attraction happens! (One atom loses an electron, the other gains it. They become oppositely charged  “bond” together!) (One atom loses an electron, the other gains it. They become oppositely charged  “bond” together!)

24. Ionic Bonding  Are these pairs likely to form ionic compounds?  Cl, Br  K, He  Na, Cl No—Both nonmetals that form negative ions. No—Helium is a noble gas that doesn’t bond with anything. Yes—Sodium is a metal that forms a positive ion, and chlorine is a nonmetal that forms a negative ion.

26. Ionic Bonding NaCl +- Two ions now stuck together!

27. NaCl

28. 4 Properties of Ionic Compounds 1. At room temp. most ionic compounds are a crystalline solid 2. Ionic compounds are brittle and shatter if hit Ions of like charge are forced near each other. 3. Because of the strong electrostatic attractions, crystalline solids are very stable and have high melting points! 4. When melted or dissolved, ionic compounds can conduct electricity!

29. All of the charged particles (ions) enable a flow of current. Saltwater Salt water Sugarwater Sugar water

30. Metallic Bonding The mobile valence electrons make them good conductors of heat and electricity! Metals are made of closely packed cations with mobile “de-localized” valence electrons. The attraction of the free-floating valence electrons for the positively charged metal ions forms the metallic bond.

31. Van der Waals Forces Attractions Between Molecules Van der Waals Forces Weaker than either the ionic or covalent bonds that form between atoms in a compound. Responsible for determining whether a compound is a liquid, gas, or solid 3 basic types from weakest to strongest –(London) Dispersion forces –Dipole interaction –Hydrogen bonding

32. London Dispersion Forces Weakest of all molecular attractions Caused by the motion of electrons producing a temporary polarity. Strength of dispersion forces generally increases as # of electrons in the molecule increases. All molecules have these weak attractions.

33. London Dispersion Forces

34. Dipole Interaction Occurs when polar molecules are attracted to one another

35. Hydrogen Bonds Occurs b/w molecules in which H is covalently bonded to either O, N, or F, which are very electronegative –Causes very polar molecules that are strongly attracted to each other –Still only has about 5% of the strength of a covalent bond

36. Hydrogen Bonds Very Important!! Reason ice is less dense than water Reason for the relatively high b.p. of water Responsible for the double helix of the DNA molecule O HH H H O + + 2 lone pairs

37. Molecules Nonpolar or Polar Molecules We now know how to determine if the bond b/w atom and atom in a compound is polar or nonpolar. But…what about the whole molecule?

38. Molecules Nonpolar or Polar Molecules Draw the Lewis Structure. If the central atom has any unshared pairs, the molecule is polar.

39. Molecules Nonpolar or Polar Molecules If there are no unshared pairs on the central atom, look at the atoms around the central atom. –If they are all the same, the molecule is nonpolar. –If any one of them is different, the molecule is polar. In a 2-atom molecule, if the bond between the 2 atoms is polar then the whole molecule is polar.

40. Molecules Nonpolar or Polar Molecules H2OH2O CO 2 HCN CH 3 Cl Polar Nonpolar Polar N2N2 Nonpolar HCl Polar

41.  Malleable – able to be pounded into sheets. Metals do not shatter as mobile electrons keep the positive metal ions from getting too close to each other.  Ductile – able to be drawn into wires