1 molecular geometry mr. bruder 2 covalent bonding a metal and a nonmetal transfer electrons –an...

1

Molecular Geometry

Mr. Bruder

2

Covalent BondingA metal and a nonmetal transfer

electrons– An ionic bond

Two metals just mix and don’t react– An alloy

What do two nonmetals do?– Neither one will give away an electron– So they share their valence electrons– This is a covalent bond

3

Covalent bondingMakes molecules

– Specific atoms joined by sharing electrons

Two kinds of molecules:Molecular compound

– Sharing by different elementsDiatomic molecules

– Two of the same atom– O2 N2

4

Diatomic elementsThere are 8 elements that always form

molecules

H2 , N2 , O2 , F2 , Cl2 , Br2 , I2 , and At2

Oxygen by itself means O2

The –ogens and the –ines

1 + 7 pattern on the periodic table

5

1 and 7

6

Molecular compoundsTend to have low melting and boiling

points

Have a molecular formula which shows type and number of atoms in a molecule

Not necessarily the lowest ratio

C6H12O6

Formula doesn’t tell you about how atoms are arranged

7

Polar BondsWhen the atoms in a bond are the

same, the electrons are shared equally.

This is a nonpolar covalent bond.

When two different atoms are connected, the electrons may not be shared equally.

This is a polar covalent bond.

How do we measure how strong the atoms pull on electrons?

8

ElectronegativityA measure of how strongly the atoms

attract electrons in a bond.The bigger the electronegativity

difference the more polar the bond.Use table 12-3 Pg. 2850.0 - 0.4 Covalent nonpolar0.5 - 1.0 Covalent moderately polar1.0 -2.0 Covalent polar>2.0 Ionic

9

Covalent BondingElectrons are shared by atoms.

These are two extremes.

In between are polar covalent bonds.

The electrons are not shared evenly.

One end is slightly positive, the other negative.

Indicated using small delta

10

How to show a bond is polar Isn’t a whole charge just a partial charge

means a partially positive

means a partially negative

The Cl pulls harder on the electrons

The electrons spend more time near the Cl

H Cl

11

H - F+ -

12

H - F+ -

H - F

+-H - F+

-

H - F

+-

H - F +-

H - F+-

H - F

+-

H - F

+-

13

H - F+ -

H - F

+-H - F+

-

H - F

+-

H - F +-

H - F+-

H - F

+-

H - F

+-

+-

14

Polar MoleculesMolecules with ends

15

Polar MoleculesMolecules with a partially positive end

and a partially negative endRequires two things to be true The molecule must contain polar bonds This can be determined from

differences in electronegativity.Symmetry can not cancel out the

effects of the polar bonds.Must determine geometry first.

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Polar MoleculesSymmetrical shapes are those without

lone pair on central atom – Tetrahedral– Trigonal planar– Linear

Will be nonpolar if all the atoms are the same

Shapes with lone pair on central atom are not symmetrical

Can be polar even with the same atom

17

Is it polar?HF

H2O

NH3

CCl4

CO2

CH3Cl

18

Intermolecular ForcesWhat holds molecules to each other

19Chapter 11: States of Matter and Intermolecular Forces 19

Intermolecular Forces

Intramolecular forces determine such molecular properties as molecular geometries and dipole moments

EOS

Intermolecular forces determine the macroscopic physical properties of liquids and solids

IntermolecularForces

© 2009, Prentice-Hall, Inc.


The attractions between molecules are not nearly as strong as the intramolecular attractions that hold compounds together.




They are, however, strong enough to control physical properties such as boiling and melting points, vapor pressures, and viscosities.




These intermolecular forces as a group are referred to as van der Waals forces.



van der Waals Forces

• Dipole-dipole interactions

• Hydrogen bonding

• London dispersion forces

+Intermolecular forces and melting/boiling point



Ion-Dipole Interactions

• Ion-dipole interactions (a fourth type of force), are important in solutions of ions.

• The strength of these forces are what make it possible for ionic substances to dissolve in polar solvents.

Chapter 11: States of Matter and Intermolecular Forces 26

Dispersion ForcesDispersion forces are forces of attraction between an instantaneous dipole and an induced dipole

… also called London forces after Fritz London who offered a theoretical explanation in 1928

The polarizability of an atom or molecule is a measure of the ease with which electron charge density is distorted by an external electrical field

EOS

Dipoles can be induced in molecules

London forcesLondon forces

Instantaneous dipole: Induced dipole:

Eventually electrons are situated so that tiny dipoles form

A dipole forms in one atom or molecule, inducing a

dipole in the other



London Dispersion Forces

• These forces are present in all molecules, whether they are polar or nonpolar.

• The tendency of an electron cloud to distort in this way is called polarizability.

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Dispersion force

H H H HH H H H

+ -

H H H H

+ - +



Factors Affecting London Forces

• The strength of dispersion forces tends to increase with increased molecular weight.

• Larger atoms have larger electron clouds which are easier to polarize.


Dispersion ForcesThe greater the polarizability of molecules, the stronger the intermolecular forces between them

EOS

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Dispersion ForceDepends only on the number of

electrons in the molecule

Bigger molecules more electrons

More electrons stronger forces

• F2 is a gas

• Br2 is a liquid

• I2 is a solid


Dipole–Dipole ForcesDipole–dipole forces arise when permanent dipoles align themselves with the positive end of one dipole directed toward the negative ends of neighboring dipoles

EOS

A permanent dipole in one molecule can induce a dipole in a neighboring molecule, giving rise to a dipole–induced dipole force

34

Dipole interactionsOccur when polar molecules are

attracted to each other.

Slightly stronger than dispersion forces.

Opposites attract but not completely hooked like in ionic solids.



Dipole-Dipole Interactions

• Molecules that have permanent dipoles are attracted to each other.– The positive end of one is

attracted to the negative end of the other and vice-versa.

– These forces are only important when the molecules are close to each other.

36

Dipole interactionsOccur when polar molecules are

attracted to each other.

Slightly stronger than dispersion forces.

Opposites attract but not completely hooked like in ionic solids.

H F

H F



Dipole-Dipole Interactions

The more polar the molecule, the higher is its boiling point.

39

Hydrogen bondingAre the attractive force caused by

hydrogen bonded to F, O, or N.F, O, and N are very electronegative so

it is a very strong dipole.They are small, so molecules can get

close togetherThe hydrogen partially share with the

lone pair in the molecule next to it.The strongest of the intermolecular

forces.


Hydrogen Bonds

A hydrogen bond is an intermolecular force in which a hydrogen atom covalently bonded to a nonmetal atom in one molecule is simultaneously attracted to a nonmetal atom of a neighboring molecule

EOS

The strongest hydrogen bonds are formed if the nonmetal atoms are small and highly electronegative – e.g., N, O, F



Hydrogen Bonding

• The dipole-dipole interactions experienced when H is bonded to N, O, or F are unusually strong.

• We call these interactions hydrogen bonds.

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Hydrogen Bonding

HH

O+ -

+

H HO+-

+

43

Hydrogen bonding

HH

O H HO

HH

O

H

H

OH

HO

H

HO HH

O



Summarizing Intermolecular Forces

MolecularGeometries

and Bonding


Valence Shell Electron Pair Repulsion Theory (VSEPR)

“The best arrangement of a given number of electron domains is the one that minimizes the repulsions among them.”

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VSEPRValence Shell Electron Pair Repulsion.

Predicts three dimensional geometry of molecules.

Name tells you the theory.

Valence shell - outside electrons.

Electron Pair repulsion - electron pairs try to get as far away as possible.

Can determine the angles of bonds.

And the shape of molecules

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VSEPRMolecules take a shape that puts electron

pairs as far away from each other as possible.

Have to draw the Lewis structure to determine electron pairs.

bondingnonbonding lone pairLone pair take more space.Multiple bonds count as one pair.

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VSEPRThe number of pairs determines

– bond angles

– underlying structure

The number of atoms determines

– actual shape

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VSEPRElectronpairs

BondAngles

UnderlyingShape

2 180° Linear

3 120° Trigonal Planar

4 109.5° Tetrahedral

590° &120°

Trigonal Bipyramidal

6 90° Octagonal

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Actual shape

ElectronPairs

BondingPairs

Non-BondingPairs Shape

2 2 0 linear

3 3 0 trigonal planar

3 2 1 bent4 4 0 tetrahedral4 3 1 trigonal pyramidal4 2 2 bent

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Actual Shape

ElectronPairs

BondingPairs


5 5 0 trigonal bipyrimidal

5 4 1 See-saw

5 3 2 T-shaped5 2 3 linear

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Actual Shape

ElectronPairs

BondingPairs


6 6 0 Octahedral

6 5 1 Square Pyramidal

6 4 2 Square Planar6 3 3 T-shaped6 2 1 linear

MolecularGeometries

and Bonding


What Determines the Shape of a Molecule?

• Simply put, electron pairs, whether they be bonding or nonbonding, repel each other.

• By assuming the electron pairs are placed as far as possible from each other, we can predict the shape of the molecule.

MolecularGeometries

and Bonding


Electron Domains

• We can refer to the electron pairs as electron domains.

• In a double or triple bond, all electrons shared between those two atoms are on the same side of the central atom; therefore, they count as one electron domain.

• The central atom in this molecule, A, has four electron domains.

MolecularGeometries

and Bonding


Electron-Domain Geometries

These are the electron-domain geometries for two through six electron domains around a central atom.

MolecularGeometries

and Bonding


Electron-Domain Geometries

• All one must do is count the number of electron domains in the Lewis structure.

• The geometry will be that which corresponds to the number of electron domains.

MolecularGeometries

and Bonding


Molecular Geometries

• The electron-domain geometry is often not the shape of the molecule, however.

• The molecular geometry is that defined by the positions of only the atoms in the molecules, not the nonbonding pairs.

MolecularGeometries

and Bonding


Molecular Geometries

Within each electron domain, then, there might be more than one molecular geometry.

MolecularGeometries

and Bonding


Linear Electron Domain

• In the linear domain, there is only one molecular geometry: linear.

• NOTE: If there are only two atoms in the molecule, the molecule will be linear no matter what the electron domain is.

MolecularGeometries

and Bonding


Trigonal Planar Electron Domain

• There are two molecular geometries:– Trigonal planar, if all the electron domains are

bonding,– Bent, if one of the domains is a nonbonding pair.

MolecularGeometries

and Bonding


Nonbonding Pairs and Bond Angle

• Nonbonding pairs are physically larger than bonding pairs.

• Therefore, their repulsions are greater; this tends to decrease bond angles in a molecule.

MolecularGeometries

and Bonding


Multiple Bonds and Bond Angles

• Double and triple bonds place greater electron density on one side of the central atom than do single bonds.

• Therefore, they also affect bond angles.

MolecularGeometries

and Bonding


Tetrahedral Electron Domain

• There are three molecular geometries:– Tetrahedral, if all are bonding pairs,– Trigonal pyramidal if one is a nonbonding pair,– Bent if there are two nonbonding pairs.

MolecularGeometries

and Bonding


Trigonal Bipyramidal Electron Domain

• There are two distinct positions in this geometry:– Axial– Equatorial

MolecularGeometries

and Bonding



Lower-energy conformations result from having nonbonding electron pairs in equatorial, rather than axial, positions in this geometry.

MolecularGeometries

and Bonding



• There are four distinct molecular geometries in this domain:– Trigonal bipyramidal– Seesaw– T-shaped– Linear

MolecularGeometries

and Bonding


Octahedral Electron Domain

• All positions are equivalent in the octahedral domain.

• There are three molecular geometries:– Octahedral– Square pyramidal– Square planar

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Molecular OrbitalsThe overlap of atomic orbitals from

separate atoms makes molecular orbitals

Each molecular orbital has room for two electrons

Two types of MO

– Sigma ( σ ) between atoms

– Pi ( π ) above and below atoms

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Sigma bonding orbitals From s orbitals on separate atoms

+ +

s orbital s orbital

+ ++ +

Sigma bondingmolecular orbital

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Sigma bonding orbitals From p orbitals on separate atoms

p orbital p orbital

Sigma bondingmolecular orbital

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Pi bonding orbitalsP orbitals on separate atoms

Pi bondingmolecular orbital

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Sigma and pi bondsAll single bonds are sigma bonds

A double bond is one sigma and one pi bond

A triple bond is one sigma and two pi bonds.

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Hybrid Orbitals

Combines bonding with geometry

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HybridizationThe mixing of several atomic orbitals to

form the same number of hybrid orbitals.

All the hybrid orbitals that form are the same.

sp3 -1 s and 3 p orbitals mix to form 4 sp3 orbitals.

sp2 -1 s and 2 p orbitals mix to form 3 sp2 orbitals leaving 1 p orbital.

sp -1 s and 1 p orbitals mix to form 2 sp orbitals leaving 2 p orbitals.

75

Hybridization109.5º with s and p

Need 4 orbitals.

We combine one s orbital and 3 p orbitals.

Make sp3 hybrid

sp3 hybridization has tetrahedral geometry.

78

sp3 geometry

109.5º

This leads to tetrahedral shape.

Every molecule with a total of 4 atoms and lone pair is sp3 hybridized.

Gives us trigonal pyramidal and bent shapes also.

79

How we get to hybridizationWe know the geometry from

experiment.

We know the orbitals of the atom

hybridizing atomic orbitals can explain the geometry.

So if the geometry requires a 109.5º bond angle, it is sp3 hybridized.

80

sp2 hybridization

C2H4

double bond counts as one pair

Two trigonal planar sectionsHave to end up with three blended

orbitalsuse one s and two p orbitals to make

sp2 orbitals. leaves one p orbital perpendicular

C C

H

HH

H

83

Where is the P orbital?Perpendicular

The overlap of orbitals makes a sigma bond ( bond)

84

Two types of BondsSigma bonds ( from overlap of orbitalsbetween the atomsPi bond ( bond) between p orbitals.above and below atomsAll single bonds are

bondsDouble bond is 1

and 1 bond Triple bond is 1

and 2 bonds

85

CCH

H

H

H

86

sp2 hybridizationwhen three things come off atom

trigonal planar

120º

one bond

87

What about twowhen two things come off

one s and one p hybridize

linear

88

sp hybridizationend up with two lobes 180º

apart.

p orbitals are at right angles

makes room for two bonds and two sigma bonds.

a triple bond or two double bonds

89

CO2

C can make two and two O can make one and one

CCOO OO

90

N2

91

N2

1 molecular geometry mr. bruder 2 covalent bonding a metal and a nonmetal transfer electrons –an...

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