Chapter 6: Chemical Bonding

Why do atoms bond? To achieve the most favorable (lowest energy) arrangement

Octet Rule Atoms will give, take, or share electrons in order to have 8 valence electrons Exception: Hydrogen wants only 2 & Helium has only two

Valence Electrons Ar Ar Na Mg Al N O Cl Electrons in the outermost energy level (s, p) Ex. Na – 1s22s22p63s1 – 1 valence e- Ex. Cl – 1s22s2p63s23p5 – 7 valence e- Shown in dot diagrams (1 dot per valence e-): Na Mg Al N O Cl Ar Ar or

What’s special about 8? The number of valence e-s in the noble gases Indicates stability (low energy)

Types of bonding Determined by the difference in electronegativity between atoms involved Covalent - sharing electrons Ionic – transfer of electrons Metallic - sea of electrons

Covalent Bonding Occurs when 2 atoms SHARE valence electrons so that both follow the octet rule No ions involved Nonmetal + Nonmetal

Example … F – 7 valence e-s, wants 8 F 3p 3s One unpaired e-

F F Example, cont. Another F also has 7 valence e-s 3p 3p 3s 3s One unpaired e-

Example, cont. By sharing electrons… F F

Example, cont. By sharing electrons… F F

Example, cont. By sharing electrons… F F

Example, cont. By sharing electrons… F F

Covalent bonding By sharing electrons… F F

F F Example, cont. …both end with full orbitals 3p 3p 3s 3s One unpaired e-

Example, cont. F F 8 Valence electrons

Example, cont. F F 8 Valence electrons

F F Lewis Structure Covalent bond – 2 e-s Lone pair – not in bond aka unshared pair F F Covalent bond – 2 e-s Molecule – group of atoms held together by covalent bonds

Structural Formula F F Covalent bond – 2 e-s

Covalent Compounds Bond length – distance between 2 covalently bonded atoms Bond energy – energy required to break a covalent bond ex.H2 = 436 kJ/mole, N2 = 163 kJ/mole

Polar Bonds Sometimes, the electrons in a covalent bond aren’t shared equally, causing partial charges Electronegativity difference determines polarity: 0-0.3 – nonpolar, 0.3-1.7 – polar, 1.7-3.3 - ionic H Cl Electronegativity = 2.10 Electronegativity = 3.16

Properties of Covalent Compounds Any state (solid, liquid, or gas) Nonconductors Low melting point and boiling point

Drawing Dot Structures Count total number of valence e-s Arrange atoms with bonds between them -Symmetry, C in center 3. Add lone pairs to fulfill octet rule -H never gets lone pairs 4. Count total number of e-s in structure 5. Add extra bonds if necessary -If there are too many e-s in structure

Example: CH4 C H Total valence = 4 + 1 + 1 + 1 + 1 =8 Arrange atoms: Add lone pair(s) Count e-s in structure = 8 *No extra bonds needed* H C

Example CH3I H I C Total valence = 4 + 1 + 1 + 1 + 7 = 14 Arrange atoms: Add lone pair(s) Count e-s in structure = 14 *No extra bonds needed* H I C

Example C2H4 H H H H C C C C H H H H Count : 4+4+1+1+1+1 = 12 Draw: 3. Add “dots” 4. Re-count: 14 5. Add extra bond(s): 14-12 = 2 extra e-s = add 1 bond Double bond (4 e-s) H H H H C C C C H H H H

Double/Triple bonds CONS like to multiple bond Determine number of bonds to add For each bond you add, you must subtract TWO lone pairs (one from each atom the bond touches) **Double bond- shorter/stronger than single ***Triple bond – shorter/stronger than double

Example N2 N N N N Count: 5+5 = 10 Draw: Add “dots” Re-count = 14 14-10 = 4  2 extra bonds needed N N N N

Example HCN H C N H C N Count: 1+4+5 = 10 Draw: Add “dots” Re-count = 14 14-10 = 4  2 extra bonds needed H C N H C N

Example CO2 O C O O C O Count: 4+6+6 = 16 Draw: Add “dots” Re-count = 20 20-16 = 4  2 extra bonds needed O C O O C O

Polyatomic Ions Group of atoms covalently bonded together that has an overall charge Add or subtract electrons from initial count

- Example NO2- O N O O N O Count: 5+6+6+1 = 18 Draw: Add “dots” Re-count = 20 20-18 = 4  1 extra bond needed - O N O O N O

PO4-

Resonance Occurs when a molecule has more than one plausible structure In nature, the actual structure is an average of all possible resonance structures - - O N O O N O

VSEPR Theory Valence shell electron pair repulsion theory Predicts the 3D shape of molecules Caused by repulsion of electrons in bonds and lone/unshared pairs 5 shapes (there are many more you will learn in college chemistry)

1. Linear -1 bond OR -2 bonds with no lone pair(s) on central atom O C O H F

2. Bent H O H 2 bonds with lone pair(s) on central atom

3. Trigonal Planar 3 bonds with no lone pair on central atom F B F F

4. Trigonal Pyramidal 3 bonds with a lone pair on central atom H N H H

5. Tetrahedral 4 bonds from central atom H H C H H

Molecular Models Ammonia – NH3 Carbon Sulfide – CS2 Methane – CH4 Propane – C3H8 Formaldehyde – H2CO Urea – H4N2CO Glucose – C6H12O6

Polar Molecules Just like individual bonds can be polar, so can molecules Shape of molecule helps determine if the molecule is polar If bonds within molecule are not polar, the molecule CANNOT be polar

Example – H2O Bonds are polar; molecule is polar

Example – CO2 Bonds are polar; molecule is NOT polar

Example – CCl4 Bonds are polar; molecule is NOT polar

Example – NH3 Bonds are polar; molecule is polar

Ionic Bonding - Cl Na+ Cl Na NaCl Metal bonded to a nonmetal One atom TRANSFERS electron(s) to another so both follow the octet rule - Cl Na+ Cl Na Cation 1s22s22p6 Wants to lose 1e- 1s22s22p63s1 Wants to gain 1e- 1s22s22p63s23p5 Anion 1s22s22p63s23p6 NaCl formula unit

Other examples… Al Cl - Al3+ 3 Cl Cl AlCl3 Cl

Try these Na and O Al and S

Lattice Energy Energy released when one mole of an ionic compound is formed

Properties of Ionic Compounds High melting point and boiling point Brittle Dissolve in water Conduct electricity when molten or dissolved in water Exist in crystal lattice structure Made up of formula units

Lattice structure

Metallic Bonding Bonding that results from the attraction between the metal atoms (their nuclei) and the surrounding sea of e-s Empty p and d orbitals allow electrons to be VERY mobile (“sea” of electrons)

Properties caused by “sea” of e-s High electrical and thermal conductivity e-s become excited and when return to ground state emit light - luster Malleable Ductile

Intermolecular forces (IMF) Forces of attraction between molecules Weaker than “real” bonds Cause property differences between compounds (boiling point, surface tension …) 3 types you need to know

1. Dipole-Dipole Must have POLAR molecule Created by attraction between opposite ends of polar molecules Example: Br-F (polar) -20°C bp F-F (not) -188 °C bp ɗ+ ɗ- ɗ+ ɗ- ɗ- ɗ+ ɗ- ɗ+

2. Hydrogen Bonding EXTREMELY strong type of dipole-dipole Occurs when H is bonded to a highly electronegative atom (FON) and there is a nearby lone pair on another highly electronegative atom (FON) H2O ɗ- ɗ+ Causes many of water’s properties (high boiling point, surface tension) ɗ- ɗ+ Hydrogen bond ɗ- ɗ+

3. London Dispersion Forces Weakest IMF Exists in ALL molecules Caused by temporary dipoles created as electrons move More electrons = stronger force

Chapter Review Pg. 209 #’s 1-6, 10-17, 19-21, 24, 28, 30, 46-48