Covalent BondingChapter 8Chemistry 2

Molecular Compounds 8.1

Molecules and Molecular Compounds 8.1

• Covalent bond – SHARE e-• Molecule – neutral group of atoms joined

by covalent bond – Diatomic = 2 atoms = O2

– Molecular Compound – compound composed of molecules = CO or CO2

• Lower mp & bp than ionic compounds• Normally 2 or more nonmetals

Molecular Formula 8.1• Chemical formula• H20 or CO2 or O2

– 1 is omitted if there is only 1 atom• No structure or arrangement of atoms

– Use diagrams

The Nature of Covalent Bonding 8.2

The Octet Rule in Covalent Bonding 8.2

• Share electrons to attain e- conf. of noble gas = 8 e-

• Combination s of Group 4A, 5A, 6A & 7A likely to form covalent bonds

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The Electron Probability The Electron Probability Distribution for the HDistribution for the H22

Molecule Molecule

Single Covalent Bonds 8.2• 2 atoms held together by single pair of e-• 2 dots or H-H by each other represent• Structural Formula = H-H

– Represent bonds and arrangement– Unshared pair = valence e- that is not shared

Double or Triple Covalent Bonds 8.2

• Share 2 pairs or 3 pairs of e-

Practice ProblemPractice Problem• Write a lewis structure for CClWrite a lewis structure for CCl22FF22 Step 1: Arrange Atoms (Carbon is “Central Atom” because is has the lowest Step 1: Arrange Atoms (Carbon is “Central Atom” because is has the lowest

group number and lowest electronegativity.group number and lowest electronegativity.Step 2: Determine total number of valence electronsStep 2: Determine total number of valence electrons

1 x C(4) + 2 x Cl(7) + 2 x F(7) = 321 x C(4) + 2 x Cl(7) + 2 x F(7) = 32 Step 3: Draw in valence electronsStep 3: Draw in valence electrons Step 4: Draw single bonds in replace of 2 electrons Step 4: Draw single bonds in replace of 2 electrons between 2 atoms and between 2 atoms and

subtract 2 e- for subtract 2 e- for each single bond (4 x 2 = 8) so 32 – 8 = 24 each single bond (4 x 2 = 8) so 32 – 8 = 24 remainingremaining

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Coordinate Covalent Bonds 8.2

• Covalent bond in which one atom contributes both bonding e-

• Molecular Formula = CO• Structural Formula = C O• Polyatomic Ion – tightly bound group of

atoms that has a + or – charge and behaves as a unit– H+ attaches to NH3’s unshared e-

– LOOK at page 225 SO3-2

Bond Dissociation Energies 8.2

• E required to break the bond between 2 covalently bonded atoms

• H + H H2 = gives off large amount of heat– Product more stable than reactants

• Big b.d.e. = strong covalent bond = normally unreactive– C-C = 347 kJ/mol– C C = 657 kJ/mol– C C = 908 kJ/mol

Resonance• 2 or more possible e- dot structures

– No back and forth changes actually occur– Just a way to vision

• Drawing – Must adhere to octet rule– Sigma bonds not altered, pi and nonbonding

e- are altered

Exceptions to the Octet Rule• Can occur when odd number of valence

e-• Atom requires less than octet of 8 e-

– BF3-NH3

• Some expand octet to 10 or 12 (esp w/ P and S

Bonding Theories 8.3

Molecular Orbitals 8.3• Orbitals overlap

– REMEMBER: atomic orbitals are orbitals in s,p,d,f• Bonding orbital – molecular orbital that can be

occupied by 2 e- of covalent bond• SIGMA BONDS σ

– 2 atomic orbitals combine– Directly between 2 nuclei– Single bonds

• p overlaps end to end

• Pi Bonds π– 2nd bond of double bond, 2nd and 3rd bond of a triple

bond (sigma is 1st of double)– Makes up 2 lobes

• Tend to be weaker than sigma bonds• Orbital overlapping is less

VSEPR Theory• Valence-shell electron pair repulsion

theory = explains 3-D shape

Hybrid Orbitals• Provided info about molecular bonding and

molecular shape• Atomic orbitals mix to form same total

number of equivalent hybrid orbitals• Single Bonds

– CH4 = sp3

http://www.mikeblaber.org/oldwine/chm1045/notes/Geometry/Hybrid/Geom05.htm

http://www.chemguide.co.uk/atoms/bonding/covalent.html

Hybrid Orbitals• Double Bonds

Molecule # of electron pairs Shapes with, and without non-

bonding e pair

Hybridization of central atom

BeH2 2 linear, linear sp

BF3 3 trigonal planar, trigonal planar

sp2

CH4 4 tetrahedral, tetrahedral

sp3

NH3 4 tetrahedral, trigonal pyramidal

sp3

H2S 4 tetrahedral, bent sp3

PF5 5 trigonal bipyramidal, trigonal

bipyramidal

dsp3

BrF3 5 trigonal bipyramidal, T-shaped

dsp3

TeCl4 5 trigonal bipyramidal, Seesaw

dsp3

SF6 6 octahedral, octehedral

d2sp3

XeF4 6 octahedral, square planar

d2sp3

XeF2 5 trigonal bipyramidal, linear

dsp3

Polar Bonds and Molecules 8.4

Bond Polarity 8.4• Nonpolar covalent bond – equally share

electrons• Polar Covalent bond – unequal sharing

– The more electronegative, the more strongly pulls on e-

• Less electronegative atom = slightly δ+ charge• More electronegative atom = slightly δ- charge

• Use table 6.2 in Chapter 6 for electronegativity of elements– HCl

• H = 2.1• Cl = 3• Electronegativity = .9

• Conceptual Problem Page 239 # 30-31

Polar Molecules 8.4• Often in a polar bond One end of

molecule is slightly – and other end slightly +– Call DIPOLE– Ex: HCl

Attractions Between Molecules 8.4• Intermolecular attractions weaker than ionic or

covalent bonds… but they are important!!! HOW?– Determine if solids, liquids, and gases– Surface tension

• Van der Waals ForceS– Dipole Interactions – polar molecules attracted to one

another• Similar to ionic but weaker

– Dispersion Forces – caused by motion of e-• Temporarily attractive force that results when the e- in 2

adjacent atoms occupy positions that make them temporarily dipole

• Weakest of all interactions• Occurs even in non-polar

Hydrogen Bonds 8.4•Attractive forces in which a hydrogen covalently bonded to a very electronegative atom is weakly bonded to unshared e- pair

Intermolecular Attractions and Molecular Properties 8.4

• Physical properties = depends on type of bonding– Melting/boiling point = lower for covalent

compared to ionic– Few covalent bonds have high mp

• Most network solids (crystals) – solids in which all the atom s are covalently bonded

– Ex: diamond» Each C is attracted to 4 other C’s