Ch8MolecularBondingNotes(12A).ppt

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This table summarizes some of the characteristic differences between ionic and covalent (molecular) substances.

Characteristics of Ionic and Molecular Compounds

Characteristic Ionic Compound Molecular Compound

Representative unit Formula unit Molecule Bond formation Transfer of one or more

electrons between atoms

Sharing of electron parts between atoms

Type of elements Metallic and nonmetallic Nonmetallic

Physical state Solid Solid, liquid, or gas

Melting point High (usually above 300°C) High (usually below 300°C) Solubility in water Usually high High to low

Electrical conductivity

of aqueous solution Good conductor Poor to nonconducting

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Chapter 8:

Covalent Bonding

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How does H

₂

form?

+ + The nuclei repel

But they are attracted to electrons

(4)

Single Covalent Bond

• Sharing of two valence electrons. • Only nonmetals and hydrogen.

• Different from an ionic bond because they actually form molecules.

(5)

The Chemical Bond

Why do atoms sometimes take and

sometimes share to be come stable?

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Bond Type

Comparing the

electronegativities will help decide who wins, in the tug of

war for electrons!

In NaCl, chlorine is strong enough to take the electron from sodium, but a

hydrogen molecule shares the electrons because they have equal

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Electronegativity and Polarity

• Nonpolar covalent - equal sharing of electrons.

(9)

Solubility of Molecules

• Solubility: “like dissolves like” Example:

• Polar solutes dissolve in polar solvents.

– A polar molecule will dissolve in water.

• Nonpolar solutes dissolve in nonpolar solvents.

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How to show a polar bond:

• Isn’t a whole charge just a partial charge

 means a partially positive

 means a partially negative

• The Cl pulls harder on the electrons because it has a higher electronegativty value than hydrogen.

• The electrons spend more time near the Cl Another method: use arrows

• Arrow points toward the more electronegative atom.

H

Cl

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I’m not stealing, I’m sharing unequally

• We described ionic bonds as stealing electrons • In fact, all bonds share – equally or unequally. • Note how bonding electrons spend their time:

• Bonding electrons are shared in each compound, but are NOT always shared equally.

H

₂

HCl

LiCl



+



–



0



0

+

–

covalent

(non-polar) polar covalent ionic

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Bond Polarity

Electronegativity Differences and Bond Types

Electronegativity difference range

Most probable type of bond

Example

0.0–0.39 Nonpolar covalent H—H (______ -______ = ______ ) 0.4–1.0 Moderately polar

covalent (Polar)

δ+ δ–

H—Cl (______ -______ = ______ ) 1.1–1.9 Very polar covalent

(Polar)

δ+ δ–

H—F (______ -______ = ______ ) >2.0 Ionic Na+_Cl–_{(______ -______ = ______ )}

Determining the Bond Type

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Bond Type

Use your notes and let’s try a couple.

What is the bond type for H

₂

?

H =

H=

2.20 – 2.20 = 0

It is nonpolar covalent.

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Bond Type

What is the bond type for HCl?

H =

Cl=

2.20 – 3.16 = 0.96

It is polar covalent.

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Bond Type

Use your notes and let’s try a couple.

What is the bond type for NaCl?

Na =

Cl =

3.16 - .93 = 2.23

It is ionic.

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Bond Type

What is the bond type in a water molecule, H₂O?

Water has two H-O bonds!

H = O =

3.44 – 2.20 = 1.24

The H-O bonds are Polar Covalent

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Lewis Structure

Lewis structure:

•a model of a covalent molecule that shows all of the valence electrons

•Octet Rule: Each atom needs a full outer shell, i.e., 8 electrons.

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Covalent Vocabulary

• Single covalent bond - sharing of 2 electrons (1 pair) between atoms.

• Unshared pair (lone pairs) – nonbonded pair of electrons that are not shared

between atoms.

Cl Cl

_{Cl Cl}

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Covalent Vocabulary

• Double covalent bond - sharing of 4 electrons (2 pairs) between atoms.

• Triple covalent bond - sharing of 6 electrons (3 pairs) between atoms.

O O

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Bond Strength and Length

• Bond length – distance between 2

bonding nuclei.

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Bond dissociation energy

• Bond dissociation energy (bond strength)

– energy required to break the bond

between two covalently bonded atoms.

– Sum of the bond dissociation values is the amount of chemical potential energy in a molecule

– Bond strength is inversely proportional to bond length

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How to show how bonds are

formed:

• It’s like a jigsaw puzzle.

• I have to tell you the final formula.

• You put the pieces together to end up with the right formula.

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Water

H

O

Each hydrogen has 1 valence electron

Each hydrogen wants 1 more

The oxygen has 6 valence electrons

The oxygen wants 2 more

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Water

• Put the pieces together • The first hydrogen is happy

• The oxygen still wants one more

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Water

• The second hydrogen attaches • Every atom has full energy levels • A pair of electrons is a single bond

H

O

H

O

Single covalent bond

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How do you draw molecules?

• For the simple molecules, the key to drawing is to find the number of bonds.

• For octet structures, an easy way to find the number of bonds is to use the NASB method !

H-Cl

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Drawing Covalent Structures

1. First, find the number of bonds

(shared e- pairs) using the following formula:

N – A = S

N = Total number of e- needed for all to have a full

octet

(Remember while most things need 8, H and He need 2)

A = Total number of e- available

(For polyatomic ions, it equals A - charge)

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Drawing Covalent Structures

2. Draw molecule without bonds.

– Draw central atom(s).

• Usually the least electronegative element (never H or halogen)

• or the single element

C  N → B  O

– Draw outer atoms spaced around central atom(s).

S O

O

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Drawing Covalent Structures

3. Draw a single bond between all atoms.

– If there are left over ones, then you will need to have multiple covalent bonds!

– Multiple bonds will go between carbons (if there are 2) or between central atom and B, N, or O groups(usually O).

SO₃ has 4

bonds! S

O

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Drawing Covalent Structures

4. Fill in dots (draw

them in pairs) so all atoms have full

octets. YOU MUST USE ALL

AVAILABLE ELECTRONS!

5. If it’s a polyatomic ion, put the whole

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Let’s do a few together:

N

(needed e-) (available eA -)

S

(shared e_)

#

Bonds Electron dot

8(1)+8(1)

16

7(1)+7(1)

14

16-14

2

2/2

1 H₂

F₂ HBr

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Let’s do a few together:

N

S

(shared e_)

#

8(1)+2(3) 14 5(1)+1(3) 8 14-8 6 6/2 3 NH₃

Cl₂O

O₂

C₃H₈ HCN

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Structural Examples

H C N

C O

H

_{C has 8 electrons}

because each

line is 2 electrons

_{Ditto for N}

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Extra Hints for NASB

H Be B C N O F Ne

#bonds desired

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Coordinate Covalent Bond

• A bond in which one atom contributes both bonding electrons.

• Carbon monoxide • CO

O

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Coordinate Covalent Bond

A bond in which one atom contributes both bonding electrons.

Carbon monoxide CO

O

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Coordinate Covalent Bond

Coordinate Covalent Bond: a bond in

which one atom contributes both bonding electrons.

Carbon monoxide CO

O

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How do we know if we have

coordinate covalent bonds?

• You have to draw the diagram and see what happens.

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Polyatomic Ions

• Polyatomic ions are covalently bonded atoms that have gained or lost electrons to become stable.

• Because they are covalent, we can use N-A=S to draw these structures.

• The only difference is that the gain or loss of electrons has effected how we calculate A

– Find A like before, then subtract the charge to find the actual value of A.

• Draw NO₂- N =

A = S =

8(3) = 24

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Resonance

• When one Lewis structure does not correctly represent the molecule you have resonance.

• Resonance structures are structures that occur when it is possible to draw two or more valid electron dot

structures that have the same number of electron pairs for a molecule or ion.

• This occurs when there is an equal choice for placing multiple bonds.

• Draw all structures connected by double arrows.

• It is a mixture of both, like a mule.

N

O O

N

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N A S B Electron Dot

SO

₃

SO

₃

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Non-Octet Structures

When you have a structure that doesn’t follow the octet rule, you can’t use

N-A=S!!!!

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Non-Octet Structures

Non-Octet Structures

1. Less than an octet

a. Beryllium

• Stable with only 4 shared electrons (2 bonds)

b. Boron

• Stable with only 6 shared electrons (3 bonds)

2. Expanded octets (more than an octet)

 _{Central atom has more than 4 bonds}  PCl₅ and SF₆

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Non-Octet Structures

N

S

(shared e_)

#

BeCl₂

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Bonding

• Sigma bonds

– A bond formed when two atomic orbitals overlap on the bond axis

– Its symbol is the Greek letter sigma (σ).

s atomic orbital

Bond axis Sigma-bonding

molecular orbital

 represents the nucleus

2p 2p

Bond Axis of σ bond

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Sigma bonding orbitals

• From p orbitals on separate atoms

p orbital p orbital

Sigma bonding molecular orbital

 

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Pi bonding orbitals

• P orbitals on separate atoms

• Double bonds

– A bond formed above and below the bond axis of the bonded atoms

– Termed “pi” or π bonds

– Tend to be weaker than sigma because they overlap less than sigma.

p atomic orbital

Pi-bonding molecular orbital

 represents the nucleus

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Sigma and pi bonds

Type of Bond # of sigma bonds ()

# of pi bonds ()

Single 1 0

Double 1 1

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Structural Examples

H C N

C O

H

=

=

=

= 2 2

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V

alence

S

hell

E

lectron

P

air

R

epulsion

Theory

Trigonal Planar

Tetrahedral

Trigonal bipyramidal

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Molecular Shapes

AB₂ Linear AB₃ Trigonal planar AB₄ Tetrahedral AB₅ Trigonal bipyramidal AB₆ Octahedral

AB₂E Bent

AB₃E Trigonal pyramidal

AB₂E₂ Bent

AB₄E (see saw)

AB₃E₂ T-shaped

AB₂E₃ Linear

AB₅E

Square pyramidal

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VSEPR-Geo. of Molecules

Handout

• 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.

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VSEPR Theory

• Types of e- Pairs

– Bonding pairs - form bonds

– Lone pairs - nonbonding electrons

Lone pairs repel

more strongly than

bonded pairs!!!

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VSEPR Theory

• Lone pairs reduce the bond angle between atoms.

Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

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• Draw the Lewis Diagram.

• Tally up e- pairs on central atom.

– double/triple bonds = ONE pair

• Shape is determined by the # of bonding pairs and lone pairs.

Know the 8 common shapes

& their bond angles!

Determining Molecular Shape

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Common Molecular Shapes

2 total 2 bond 0 lone

LINEAR

180°

BeH

₂

Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

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TRIGONAL PLANAR

120°

BF

₃

Common Molecular Shapes

B

B A

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Common Molecular Shapes

BENT

<120°

SO

₂

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TETRAHEDRAL

109.5°

CH

₄

Common Molecular Shapes

B

A

B B

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TRIGONAL PYRAMIDAL

107°

NH

₃

Common Molecular Shapes

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BENT

104.5°

H

₂

O

Common Molecular Shapes

H

x

H

x

O

H

x

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TRIGONAL

BIPYRAMIDAL

120°/90°

PCl

₅

Common Molecular Shapes

A Be

Be

Ba

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OCTAHEDRAL

90°

SF

₆

Common Molecular Shapes

B

B B B

B

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Review:

Electronegativity and Polarity

• Polar: unequal distribution of electrons

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Polar Molecules

Requires 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.

• Does symmetry cancel the polar bond(s)? • If yes, nonpolar.

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Bond Polarity

Electronegativity Differences and Bond Types

Electronegativity difference range

Most probable type of bond

0.0–0.39 Nonpolar covalent

0.4–1.0 Moderately polar covalent (Polar) 1.1–1.9 Very polar covalent (Polar)

>2.0 Ionic

Determining the Bond Type

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BF

₃

F

F F

B

Determining Molecular Polarity

• _{Nonpolar Molecules}

– Dipole moments are symmetrical and cancel out.

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H

₂

O

H _H

O

Determining Molecular Polarity

• _{Polar Molecules}

– Dipole moments are asymmetrical and don’t cancel .

net

dipole moment

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CHCl

₃

H

Cl _Cl

Cl

Determining Molecular Polarity

• _{Therefore, polar molecules have...}

– asymmetrical shape (lone pairs) or – asymmetrical atoms

net

dipole moment

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Is it polar?

• HCl

• CO₂

• NH₃

(72)

..

.. ..

Polar Bonds

H Cl

Polar

A molecule has a zero dipole moment because their dipoles cancel one another.

H H

O

Polar F F

B F Nonpolar H H H N Polar Polar Nonpolar

F Cl F

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Intermolecular Forces

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Holding it together

Q: Consider a glass of water.

Why do molecules of water stay together? A: There must be attractive forces.

Intramolecular forces occur between atoms

Intermolecular forces occur

between molecules

• Intermolecular forces are not considered in ionic bonding because there are no molecules.

• The type of intramolecular bond determines the type of intermolecular force.

Intramolecular forces are much

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Intermolecular Forces

London dispersion forces London dipole-dipole

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Intermolecular Forces

• They are what make solid and liquid molecular compounds possible.

• The weakest are called (London Forces) Van der Waal’s forces - there are two kinds

• Dispersion forces

– caused by momentary dipoles of constantly moving electrons between NONPOLAR molecules

• Dipole Interactions

– depend on the number of electrons – more electrons stronger forces

– Bigger molecules

• Fluorine is a gas

• Bromine is a liquid

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Dipole interactions

• Occur 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

(78)

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Hydrogen bonding

• Are 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.

– _{Examples - HF, NH}₃_{, and H}₂_O – Is it strong “H-NOF” (enough)