Slide 1

Slide 2 Slide 3 Covalent Bonds (2 nonmetals) atoms share e to get a full valence shell C1s 2 2s 2 2p 2 F1s 2 2s 2 2p 5 *Both need 8 v.e for a full outer shell (octet rule)!* o 4 valence e- 7 valence e- o x o o C x x x x x x F Bonding Review Slide 4 Draw the Lewis dot structure for the following elements (write e- config first): Si O P B Ar Br 1s 2 2s 2 2p 6 3s 2 3p 3 1s 2 2s 2 2p 4 1s 2 2s 2 2p 1 1s 2 2s 2 2p 6 3s 2 3p 6 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 5 1s 2 2s 2 2p 6 3s 2 3p 2 4 valence e- 6 valence e- 5 valence e- 3 valence e- 8 valence e- 7 valence e- Slide 5 1 2 3 4 5 6 7 8 H He Li Be B C N O F Ne Na Mg Al Si P S Cl Ar K Ca Se Br Kr Rb Sr Te I Xe Cs Ba Notice any trends? TRANSITION METALS The group # corresponds to the # of valence e Slide 6 C F F F F Lets bond two F atoms together Each F has 7 v. e and each needs 1 more e F F F F Now lets bond C and F atoms together C F F F F carbon tetrafluoride (CF 4 ) F2F2 Slide 7 Lewis Structures: 2D Structures NH 3 CH 2 O CO 2 SO 2 CH 4 Slide 8 1.Sum the # of valence electrons from all atoms Anions: add e (CO 3 2- : add 2 e ) Cation: subtract e (NH 4 + : minus 1 e ) 2.Predict the arrangement of the atoms Usually the first element is in the center (often C, never H) 3.Make a single bond (2 e ) between each pair of atoms 4.Arrange remaining e to satisfy octets (8 e around each) Place electrons in pairs (lone pairs) Too few? Form multiple bonds between atoms: double bond (4 e ) and triple bond (6 e ) 5.Check your structure! All electrons have been used All atoms have 8e- Exceptions: Remember that H only needs 2e ! Drawing Lewis Structures Slide 9 Lewis Structure Practice CH 4 H 2 O NF 3 HBr OF 2 HCN NO 3 - CO 3 2- Draw a Lewis Structure for the following compounds: NHC Slide 10 NHC Slide 11 Lewis Structure Trends Here are some useful trends C group Forms a combo of 4 bonds and no LP (Lone Pairs) i.e. CO 2 N group Forms a combo of 3 bonds and 1 LP i.e. NH 3 O group Forms a combo of 2 bonds and 2 LP i.e. CH 2 O F group (halogens) Forms 1 bond and 3 LP i.e. OF 2 Note that these are NOT always true! Slide 12 Carbonite Carbonate? CO 3 2- CO 2 2- Slide 13 Resonance Structures Resonance structures differ only in the position of the electrons The actual structure is a hybrid (average) of the resonance structures Technically NOT two single bonds and one double bond All 3 Oxygen atoms share the double bond 3 equal bonds (somewhere between a double and single) Arrow formalism: curved arrows show electron movement Show resonanceShow movement of e - Slide 14 Predicting Molecular Shape: VSEPR (Valence Shell Electron Pair Repulsion) Electrons repel each other The molecule adopts a 3-D shape to keep the electrons (lone pairs and bonded e - ) as far apart as possible Different arrangements of bonds/lone pairs result in different shapes Shapes depend on # of bonds/lone pairs (things) and LP around the central atom Slide 15 Selected Shapes and Geometries using VSEPR Things Slide 16 Carbon Dioxide: CO 2 Two things (bonds or lone pairs) Linear geometry 0 LP Linear Shape 180 o Bond angle COO Lewis Structure Slide 17 C H H O Formaldehyde: CH 2 O Three things Trigonal planar geometry 0 LP Trigonal planar shape 120 bond angles Lewis Structure Slide 18 Sulfur Dioxide: SO 2 Three things Trigonal planar geometry 1 LP Bent shape 120 bond angles Lewis Structure S O O B A A A Slide 19 Methane: CH 4 Lewis Structure Four things (bonds/LP) Tetrahedral geometry 0 LP Tetrahedral shape 109.5 o bond angles Slide 20 Ammonia: NH 3 Lewis Structure Four things (bonds/LP) Tetrahedral geometry 1 LP Trigonal pyramid shape 107 o bond angles Slide 21 Water: H 2 O Lewis Structure 4 things (bonds/LP) Tetrahedral Geometry 2 LP Bent Shape 104.5 o bond angle Slide 22 Hydrogen Chloride: HCl Four things (bonds/LP) Tetrahedral geometry 3 LP Linear Shape No Bond angle Lewis Structure ClH Slide 23 A special note For any molecule having only two atoms e.g. N 2, CO, O 2, Cl 2, HBr, etc. Geometry = Linear Shape = Linear Bond Angle(s)? = None It is much like geometry what is formed by connecting two points? a line. NN HBr Cl OO Slide 24 You will need to commit these to memory! Things Slide 25 VSEPR Practice (w/o aid of yellow sheet) CO 2 G: S: Angle: ClO 2 - G: S: Angle: NO 2 - G: S: Angle: CH 3 COO - G: S: Angle: PBr 3 G: S: Angle: AsO 4 3- G: S: Angle: Slide 26 Electronegativity and Bond Type The electronegativity difference between two elements helps predict what kind of bond they will form. Bond type Covalent Polar covalent Ionic Definition e- are evenly shared e- are unevenly shared e- are exchanged (gained or lost) Electronegativity difference 0.4 0.5 1.8 > 1.8 Slide 27 Electronegativity difference 0.4 0.5 1.8 > 1.8 Practice with Bond Types Bond type Covalent Polar covalent Ionic H 2.1 Li 1.0 Be 1.5 B 2.0 C 2.5 N 3.0 O 3.5 F 4.0 Na 0.9 Mg 1.2 Al 1.5 Si 1.8 P 2.1 S 2.5 Cl 3.0 Br 2.8 I 2.5 K 0.8 Ca 1.0 Sample Bonds NaCl Cl-Cl C-O C-H Electronegativity Difference 3.0 0.9 = 2.1 3.0 3.0 = 0 3.5 2.5 = 1.0 2.5 2.1 = 0.4 Bond Type? Ionic Covalent Polar covalent Covalent Slide 28 Dipole Moments and Polarity Arrow points toward partially - end Occurs in polar covalent bonds Uneven distribution of e - Atoms become partially charged Partially + charged end -- ++ Slide 29 Polarity Examples HCN CO 2 CO 3 2- CH 2 O SO 2 CH 4 CH 3 F C 3 H 8 CO NH 3 1.Check molecule for dipole moments (polar bonds) 2.When determining overall polarity, an imbalanced structure will likely be polar (at least partially) 3.Even with polar bonds, a balanced structure is non-polar overall 4. Any structure with lone pairs on the central atom is automatically polar! Try these with your neighbors Polar Non-polar Polar Non-polar Polar Non-polar Polar