Chpt 8 Chemical Bonding Forces holding atoms together = Chemical Bonds

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1 Chpt 8 Chemical Bonding Forces holding atoms together = Chemical Bonds Kinds of chemical bonds: 1. Ionic 2. Covalent 3. Metallic Useful guideline: Octet rule Atoms tend to gain, lose, or share e - to achieve the same number of e - as a noble gas. All noble gases (except He) have 8 e - in their outermost shell

2 Lewis symbols (e - dot symbols) for atoms " Use only outer shell (valence) e - s to write e - dot symbols Why? because only valence e - s are involved in bonding ex: I C Mg Cl - S 2- Al 3+

3 Charges of Ions, as we know, are related to e - configurations of atoms (Na +, Mg 2+, Al 3+, F -, O 2-, N 3-, etc), but: Another exception to the Octet Rule: Recall: e - configs of transition metal ions (& beyond) Examples: Ti Fe Ag Ga

4 Ionic Bonding: Metals with nonmetals; Electrostatic attraction between oppositely charged ions Results in solids with 3-dimensional lattices

5 Lattice enthalpy: Energy to separate ions in an ionic compound from each other Coulomb s law of attraction: greater charges stronger attraction smaller ions " stronger attraction NaCl(s) Na + (g) + Cl - (g) ΔH lattice = +788 kj/mol

6 Na(s) Cl 2 (g) Wieliczka Salt Mine, Krakow, Poland NaCl(s)

7 Larger ΔH lattice stronger attraction between ions in solid ex Which has the greater lattice enthalpy? a. NaCl or MgO b. NaCl or KBr

8 Ionic bonds Covalent bonds

9 Covalent Bonding =Sharing of valence e - s Between neutral nonmetal atoms Follows octet rule (usually) Types of Covalent Bonds: Polar NonPolar Single bonds Multiple bonds (double, triple)

10 Bond Polarity & Electronegativity Electronegativity (EN) = Ability of an atom in a molecule to attract and hold shared e - s to itself. Element with higher EN has greater attraction for the e - s in a bond Result: Unequal sharing of e - s = Polar Covalent Bond vs Equal sharing of e - s = Nonpolar Covalent Bond

11 Periodic trends in EN inc. EN F has greatest EN inc. EN ex Which has a greater electronegativity? a. O or C b. O or Cl c. P, S, As, or Se

12 The greater the difference in EN the more polar the bond H = 2.1 F = 4.0 Cl =3.0 Br = 2.8 I = 2.5 ex Which bond is most polar? S-S P-S As-S Se-S

13 Polar Molecules $A polar molecule has positive and negative sides Has a dipole moment ex. Water is a polar molecule: Overall dipole moment, µ, is a measure of polarity

14 So what??? Polar Molecules 1. Attract each other (therefore are soluble in each other) Sugar, C 6 H 12 O 6 (s) in coffee = Polar covalent compound dissolving in aqueous solution

15 Ethanol, C 2 H 5 OH(l) in water = Polar covalent compound dissolving in water

16 Polar Molecules 2. Attract ions (therefore dissolve ionic compounds) NaCl(s) in water = Ionic compound dissolving in water

17 Polar Molecules C 2 H 5 OH(aq) 1. Attract each other (therefore are soluble in each other) NaCl(aq) 2. Attract ions (therefore dissolve ionic compounds)

18 Polar Molecules 3. Align in electric fields (LCDs)

19 If dipole moment = 0, the molecule is nonpolar:

20 Drawing Lewis Structures for Molecules The goal: to determine the number and types of covalent bonds (polar, nonpolar, single, double, triple) The tools: # of valence e - octet rule formal charges The rules (see Handout & pages of text)

21 Drawing Lewis Structures (see also pages in text) Object: To determine the number and types of bonds in a molecule. Tools: 1. Number of Valence Electrons 2. Octet Rule 3. Formal Charges Example: Draw the Lewis Structure for NH 3 1. Add up the valence electrons from all atoms in the molecule. These now belong to the whole molecule to be distributed according to the guidelines below. N valence e - s 3 H 3 x = valence e - s Total = valence e - s 2. Assume the first atom in the formula is the central atom, unless told otherwise. Connect all other outer atoms to the central atom using single bonds. Complete the octets on outer atoms (if necessary). 3. Count the electrons in your structure. Does it equal the number you counted in Step 1? If the number of electrons in Step 1 is greater, put the additional electrons on the central atom (even if it means the central atom has more than an octet, but only if the valence shell of the central atom is n!3.) If the number in Step 1 is less, go to Step 4. Example: Draw the Lewis structures for POCl 3, NF 3 and NH If the central atom has less than an octet, try multiple bonds. Example: Draw the Lewis structures for SO 3 and CN -!! "#$%!&!'(!)!

22 Helpful Tips for Drawing Lewis Structures 1. Hydrogen has one valence e in the n = 1 shell, which contains only one s orbital (l = 0; m l = 0). As a result, H can have no more than 2 e - s. H can have a maximum of one shared pair of e - s (one single bond) and can never be a central or linker atom. 2. Fluorine, like hydrogen is never a central or linker atom (why?), so the best Lewis structure for a molecule containing F will have only one single bond to F. 3. Carbon has 4 valence e - s and so shares 4 pairs of e - s. Can form up to 4 single bonds. 4. Nitrogen has 5 valence e - s ( 3 unpaired e - s and 2 paired). Tends to share 3 pairs of e - s. Can share 4 pairs in molecular ions. 5. Some exceptions to the Octet Rule First 4 elements in the periodic table (H through Be) always have less than an octet. Element # 5, boron, can have less than an octet (most common) or can have an octet. Elements 5 through 10 in the periodic table (B through Ne) can never have more than an octet (why not??); yet all remaining elements can have more than an octet (why?).!! "#$%!&!'(!)!

23 Resonance Structures Some molecules are not adequately described by a single Lewis structure ex: Ozone, O 3 Benzene, C 6 H 6 ex Draw the three resonance structures for CO 3 2-

24 What if Not all resonance structures are the same? What is the best Lewis structure for a molecule with different resonance structures? Use Formal Charges: Charges assigned to atoms by dividing shared e - s equally

25 Formal Charges. Used to determine most likely structure for molecules where more than one possible correct Lewis structure exists. Formal Charge of an Atom = # of valence e - s - " # of unshared e - s # shared e- % s # $ & ' To use, you must calculate the formal charge on each atom in a structure. 1. The best structure will have a formal charge equal to, or closest to, zero on each atom in the molecule. Example: CO 2 2. If and only if the formal charge on each atom is not equal to zero, then the best structure will have a negative formal charge on the more electronegative element(s). Example: N 2 O!! "#$%!&!'(!&!

26 More Exceptions to the Octet Rule Odd e - molecules Atoms with less than an octet Central atom has more than an octet This is a very common exception ex PCl 5 Question: Where are these expanded octets found? When central atom has empty d orbitals in valence shell: P = [Ne]3s 2 3p 3 3d 0 [3rd shell has empty d orbitals] Answer: Only if valence shell, n, of central atom is 3

27 Bond Energies (D) Bond-Dissociation enthalpy = ΔH required to break a bond Trends 1. always positive quantities 2. ΔH (single) < ΔH (double) < ΔH (triple) bond

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29 Using bond energies to estimate ΔH rxn ΔH rxn = Σ(bonds broken) Σ(bonds made) endothermic exothermic ex Estimate ΔH for CO + H 2 O H 2 + CO 2 C O + H-O-H H- H + O=C=O Given: D(C O) = 1072 kj/mol D(O-H) = 463 kj/mol D(C=O) = 799 kj/mol D(H-H) = 436 kj/mol

30 Upon detonation, 2 moles of TNT forms 3 moles N 2 and 7 moles CO plus H 2 O and C.

31 Summary of Octet Rule Atoms often use just their outer s and p orbitals in bonding room for a max. of 8e - octet rule In drawing Lewis structures, octet rule is used to predict if molecules have single, double, or triple bonds. Type of bonds affects properties of molecules such as: Bond lengths Bond energies e - distributions Molecular shapes (Ch. 9)