Unit 4: Covalent Bonding

Bonding Review C F Covalent Bonds (2 nonmetals) …atoms share e– to get a full valence shell C 1s2 2s2 2p2 F 1s2 2s2 2p5 *Both need 8 v.e – for a full outer shell (octet rule)!* 4 valence e- 7 valence e- o x x C x F o o x x x x o

Draw the Lewis dot structure for the following elements (write e- config first): Si O P B Ar Br 1s2 2s2 2p6 3s2 3p2 4 valence e- 1s2 2s2 2p4 6 valence e- 1s2 2s2 2p6 3s2 3p3 5 valence e- 1s2 2s2 2p1 3 valence e- 1s2 2s2 2p6 3s2 3p6 8 valence e- 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p5 7 valence e-

Notice any trends…? TRANSITION METALS 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 TRANSITION METALS K Ca Se Br Kr Rb Sr Te I Xe Cs Ba The group # corresponds to the # of valence e–

F2 Let’s bond two F atoms together… Each F has 7 v. e– and each needs 1 more e– F F F F2 Now let’s bond C and F atoms together… carbon tetrafluoride (CF4) F C F F C F F

Lewis Structures: 2D Structures NH3 CH2O CO2 SO2 CH4

Drawing Lewis Structures Sum the # of valence electrons from all atoms Anions: add e– (CO32- : add 2 e– ) Cation: subtract e– (NH4+: minus 1 e– ) Predict the arrangement of the atoms Usually the first element is in the center (often C, never H) Make a single bond (2 e–) between each pair of atoms 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– ) Check your structure! All electrons have been used All atoms have 8e- Exceptions: Remember that H only needs 2e– !

Lewis Structure Trends Here are some useful trends… C group Forms a combo of 4 bonds and no LP (Lone Pairs) i.e. CO2 N group Forms a combo of 3 bonds and 1 LP i.e. NH3 O group Forms a combo of 2 bonds and 2 LP i.e. CH2O F group (halogens) Forms 1 bond and 3 LP i.e. OF2 Note that these are NOT always true!

Lewis Structure Practice Draw a Lewis Structure for the following compounds: CH4 H2O NF3 HBr OF2 HCN NO3- CO32- N H C

Resonance structures differ only in the position of the electrons Show resonance Show movement of e- 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

Electronegativity and Bond Type The electronegativity difference between two elements helps predict what kind of bond they will form. 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 Bond type Covalent  Polar covalent Ionic

Practice with Bond Types 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 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 Si 1.8 P S Cl Br 2.8 I K 0.8 Ca Electronegativity difference ≤ 0.4 0.5 – 1.8  > 1.8 Bond type Covalent  Polar covalent Ionic

Dipole Moments and Polarity Occurs in polar covalent bonds Uneven distribution of e- Atoms become partially charged Partially “+” charged end Arrow points toward partially “-” end δ+ δ-

Polarity Examples HCN CO2 CO32- CH2O SO2 Polar CH4 CH3F C3H8 CO NH3 Check molecule for dipole moments (polar bonds) When determining overall polarity, an imbalanced structure will likely be polar (at least partially) Even with polar bonds, a balanced structure is non-polar overall Any structure with lone pairs on the central atom is automatically polar! Try these with your neighbors… HCN CO2 CO32- CH2O SO2 Polar CH4 CH3F C3H8 CO NH3 Non-polar Non-polar Polar Non-polar Non-polar Polar Polar Polar Polar

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

Selected Shapes and Geometries using VSEPR “Things” 16

Carbon Dioxide: CO2 Lewis Structure Two “things” (bonds or lone pairs) Linear geometry 0 LP → Linear Shape 180o Bond angle

Formaldehyde: CH2O Lewis Structure Three “things” Trigonal planar geometry 0 LP → Trigonal planar shape 120° bond angles

Sulfur Dioxide: SO2 Lewis Structure Three “things” B A Three “things” Trigonal planar geometry 1 LP → Bent shape 120° bond angles

Methane: CH4 Lewis Structure Four “things” (bonds/LP) Tetrahedral geometry 0 LP → Tetrahedral shape 109.5o bond angles

Ammonia: NH3 Lewis Structure Four “things” (bonds/LP) Tetrahedral geometry 1 LP → Trigonal pyramid shape 107o bond angles

Water: H2O Lewis Structure 4 “things” (bonds/LP) Tetrahedral Geometry 2 LP → Bent Shape 104.5o bond angle

Hydrogen Chloride: HCl Lewis Structure Four “things” (bonds/LP) Tetrahedral geometry 3 LP → Linear Shape No Bond angle Cl Cl H

Cl N O H Br A special note… For any molecule having only two atoms… e.g. N2, CO, O2, Cl2, HBr, etc. Geometry = Linear Shape = Linear Bond Angle(s)? = None It is much like geometry… what is formed by connecting two points? …a line. Cl N O H Br

You will need to commit these to memory! “Things” 25

VSEPR Practice (w/o aid of yellow sheet) CO2 G: S: Angle: ClO2- NO2- CH3COO- G: S: Angle: PBr3 AsO43-