Unit 5B: Covalent - PowerPoint Presentation

Unit 5B: Covalent Bonding

Covalent Bonds (2 nonmetals) …atoms share e– to get a full valence shell C 1s2 2s2 2p 2 F 1s2 2s2 2p5 *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

Draw the Lewis dot structure for the following elements (write e- config first) : SiOPBArBr 1s 2 2s2 2p6 3s2 3p 3 1s 2 2s2 2p4 1s2 2s2 2p1 1s2 2s2 2p 6 3s2 3p6 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p5 1s2 2s2 2p6 3s2 3p2 4 valence e- 6 valence e- 5 valence e- 3 valence e- 8 valence e- 7 valence e-

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 –

C F F F F Let’s bond two F atoms together… Each F has 7 v. e – and each needs 1 more e – F F F F Now let’s bond C and F atoms together… C F F F F carbon tetrafluoride (CF 4 ) F 2

Lewis Structures: 2D Structures NH 3 CH 2 O CO 2 SO 2 CH 4

Sum the # of valence electrons from all atomsAnions: 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 atomsArrange 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 – !Drawing Lewis Structures

Lewis Structure PracticeCH4 H2ONF3 HBrOF2 HCNNO3-CO 32-Draw a Lewis Structure for the following compounds: N H C

N H C

Lewis Structure TrendsHere are some useful trends… C groupForms a combo of 4 bonds and no LP (Lone Pairs) i.e. CO2N groupForms a combo of 3 bonds and 1 LP i.e. NH 3O groupForms a combo of 2 bonds and 2 LP i.e. CH2OF group (halogens) Forms 1 bond and 3 LP i.e. OF2Note that these are NOT always true!

CarboniteCarbonate?CO32- CO22-

Resonance Structures Resonance structures differ only in the position of the electrons The actual structure is a hybrid (average) of the resonance structuresTechnically NOT two single bonds and one double bondAll 3 Oxygen atoms share the double bond 3 equal bonds (somewhere between a double and single)Arrow formalism: curved arrows show electron movement Show resonance Show movement of e -

Predicting Molecular Shape: VSEPR (Valence Shell Electron Pair Repulsion) Electrons repel each otherThe molecule adopts a 3-D shape to keep the electrons (lone pairs and bonded e-) as far apart as possibleDifferent 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”

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

C H H O Formaldehyde: CH 2 O Three “things” Trigonal planar geometry 0 LP → Trigonal planar shape 120° bond angles Lewis Structure

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

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

Ammonia: NH 3 Lewis Structure Four “things” (bonds/LP) Tetrahedral geometry 1 LP → Trigonal pyramid shape 107 o bond angles

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

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

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. N N H Br Cl Cl O O

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

VSEPR Practice (w/o aid of yellow sheet)CO2 G:S:Angle:ClO2-G: S:Angle:NO2-G:S:Angle: CH3COO-G: S: Angle:PBr3G:S:Angle:AsO43-G:S:Angle:

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 sharede- are unevenly sharede- are exchanged (gained or lost) Electronegativity difference ≤ 0.40.5 – 1.8> 1.8

Electronegativity difference ≤ 0.4  0.5 – 1.8> 1.8 Practice with Bond Types Bond type Covalent  Polar covalent  Ionic H2.1 Li1.0Be1.5B2.0C2.5N3.0 O3.5F4.0Na0.9Mg1.2Al1.5 Si1.8P2.1S2.5Cl3.0Br2.8 I2.5K0.8Ca1.0 Sample BondsNaClCl-ClC-OC-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?IonicCovalentPolar covalent Covalent

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 δ - δ +

Polarity ExamplesHCNCO 2CO32-CH2OSO 2CH4CH3F C3H8CONH3 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… Polar Non-polar Non-polar Polar Polar Non-polar Polar Non-polar Polar Polar