1. Chemical Bonding

2. Although we have talked about atoms and molecules individually, the world around us is almost entirely made of compounds and mixtures of compounds. We are going to take an in depth look at these compounds and the interactions of the atoms that hold them together and make up the compounds

3. Bonds Bonds are the force that holds groups of two or more atoms together and makes them function as a unit. Bond Energy is the energy required to break the bond between two atoms.

4. Ionic Bonding Generally occurs between a Metal and a Non- metal Cations lose electrons, Anions gain electrons Electrons are transferred Opposite charges on atoms attracts them to one another

5. Covalent Bonding Generally occurs between a nonmetal and a nonmetal Both atoms share electrons to achieve a lower energy state. Electrons are “shared”. More like a tug of war. Same charges, lower energy is responsible for bonds

6. Nonpolar Covalent Bonding Covalent bonds where electrons are shared equally between two atoms. Atoms must have the same values of electronegativity If a covalent bond is like a tug of war a nonpolar covalent bond would be a stalemate.

7. Polar Covalent Bonding During the “tug of war” in covalent bonding electrons aren’t always shared equally. Some atoms have a stronger attraction for electrons and pull them closer than other atoms. This unequal sharing of electrons causes one atom to have a small positive charge and one to have a small negative charge

8. Electronegativity Attraction of shared electrons to an atom. Determines the type of bond Can calculate a value of electronegativity based on relative values for each element.

10. Electronegativity For differences in electronegativity, generally: O = nonpolar covalent examples: Cl-Cl, C-S.1-1.5 = polar covalent examples: C-F, P-S 1.6-3.3 = Ionic examples: Na-O, K-I

11. Practice Identify the following as Ionic, polar covalent, or nonpolar covalent bonds: S-FMg-Cl Br-BrB-F B-NN-Cl P-IMn-S

12. Dipole Moments Polar covalent bonds that do not share electrons equally are said to have a dipole moment. The atom pulling the electrons the strongest or with the higher electronegativity will have a partial negative charge. The atom with the weaker pull on electrons will have a partial positive charge.

13. Lewis Dot Structures Representations of atoms or molecules which show the valence electrons around an atom or molecule Hydrogen follows a duet rule – two valence electrons give it the same electron configuration as helium Most other atoms follow a octet rule – eight valence electrons will give each atom the same number of valence electrons as a noble gas

14. Lewis Dot Structures - Ionic Metals lose electrons, nonmetals gain electrons Rules for LDS for ionic compounds: 1.Write each element symbol 2.Determine the number of valence electrons 3.Add the valence electrons to each atom. Clockwise- 12,3,6,9 one at a time. 4.Show the electron transfer from metal(s) to nonmetal(s)

15. Lewis Dot Structures-Practice Draw the Lewis Dot Structure for the following atoms: Ca F Se Al P Si

16. Lewis Dot Structures - Practice Draw the Lewis Dot Structures for the following ionic compounds: Na + Cl Mg + Br Al + O B + F

17. Lewis Dot Structures -Covalent Two nonmetals share electrons to achieve a lower energy Rules: 1-3 same as ionic 4. Circle electrons that will pair together 5. Rearrange the compound so shared electrons are aligned correctly(between atoms).

18. Lewis Dot Structures - Practice Draw the LDS for the following covalent compounds: CCl 4 NBr 3 Phosphorus + Iodine Silicon + Fluorine

19. Bond Strength Of the three covalent bonds: A triple bond is the strongest followed by a double bond and then a single bond which is the weakest of the three A triple bond has the highest bond energy, then a double bond, followed by a single bond

20. Bond Length Of the three covalent bonds: A triple bond is the shortest followed by a double bond and then a single bond which is the longest Why?

21. LDS- Covalent Compounds So far we have looked at simple covalent compounds and how they will share valence electrons Now we will look at more complex covalent compounds and how to determine the Lewis Dot Structure.

22. Try drawing the Lewis Dot Structure for SO 2 using the rules for covalent compounds that you have learned The actual structure of the compound is: Let’s take a look at how to draw LDS when we are given the formula and the compound is more complex

23. Rules for Complex Covalent LDS 1. Determine the total number of valence electrons in the compound 2. Begin by putting a single bond between each atom (Choose appropriate middle atom if necessary) 3. Fill in lone pair electrons to fulfill duet/octet rule 4. Add a double bond (or triple bond) if necessary to insure the duet/octet rules are fulfilled and the total number of valence electrons are correct.

24. Resonance Resonance occurs when several equally correct Lewis Dot structures can be assigned to compounds. Double arrows are used to show options for compounds with resonance structures. Resonance structures for SO 2 :

25. Practice: Draw the LDS for the following compounds with the new rules you have been given: HFN 2 NH 3 CH 4 NF 3 O 2 COPH 3

26. LDS for Ions For ions the rules for drawing Lewis Dot Structures are the same except the total number of valence electrons will either increase or decrease depending on the charge For a positive charge subtract a valence electron For a negative charge add a valance electron After Drawing the LDS brackets are added and the charge is added outside the brackets- top right

27. Example: NH 4 + Add up valence electrons for each atom: N-5 H-1 total = 9 Because of the +1 charge we assume a valence electron has been lost Our new total of valence electrons is 8 Draw the LDS using the same rules. Don’t forget the brackets and the charge

28. Practice Complete the LDS for the following ions. Show resonance structures if they exist NO + NO 3 - SO 4 -2 ClO 3 - PO 4 -3 SCN -

29. Lone Pair Electrons The unshared valence electrons represented in LDS are called lone pair electrons. Example: In CF 4 each fluorine has six lone pair electrons and carbon has zero for a total of twenty four.

30. VSEPR Stands for Valence Shell Electron Pair Repulsion This theory states that electrons pairs around an atom will spread out as far as possible This repulsion is due to the same charges on electrons

32. Molecule Polarity We have already discussed a polar covalent bond in terms of dipole moments caused by differences in electronegativity. We will now use this knowledge to determine whether a molecule is polar or non polar

33. Molecule Polarity In order for a molecule to be considered polar it needs to have a concentrated partial positive charge on one end and a concentrated partial negative charge on the other. Molecules that have polar bonds will not necessarily be polar molecules

34. Molecule Polarity Symmetrical molecules can have dipole moments cancel each other out causing them to be nonpolar. Examples CH 4, BF 3, CO 2

35. Molecule Polarity A molecule with a lone pair of electrons in place of a bond will always be polar (bent and trigonal pyramidal) Examples: NH 3, H 2 O, NO 2 -

36. Molecule Polarity Molecules that have a tetrahedral, trigonal planar, and linear geometry can be either polar or nonpolar. The central atom would have to have different atoms bonded to it to be a polar molecule. Example: HCN, HCO 2 -