08a Bonding 6.pdf

(1)

Chapters 8 and 9

Part 1 representing molecules

Why are these compounds so different?

SiO

2

(Sand and Quartz) Versus CO

2

Types of Chemical Bonds

What

is meant by the term “chemical bond?”

Attractive forces that hold groups of atoms together and make them function as a unit. The electrons involved in

bonding are usually those in the outermost (valence) shell.

Why

do atoms bond with each other to form molecules?

A bond will form if the energy of the aggregate is lower than that of the separated atoms.

How

do atoms bond with each other to form molecules?

We will answer this question in Ch 8 and Ch 9.

Two Types of Chemical Bonds

•Ionic Bonding (formula units)

•Results from electrostatic attractions among ions, which are formed by the transfer of one or more electrons from one atom to another.

•Formed between elements with large differences in electronegativity (often a metal and a nonmetal)

•CovalentBonding (molecules)

•Results from sharing one or more electrons between atoms.

•Formed between elements with similar electronegativities (usually two or more nonmetals)

Ionic Compounds Form Extended 3-D Arrays

of Oppositely Charged Ions

•Ionic compounds have high melting points because the Coulomb force, which holds ionic compounds together, is strong.

•Na+_Cl-

Coulomb’s Law Describes the Attraction Between

Positive and Negative Ions

•Coulomb’s Law E a

Where E = energy of attraction between ions q = magnitude of charge on ions d = distance between center of ions

Small ions with high ionic chargeshave

large

Coulombic forces of attraction.

Large ions with small ionic chargeshave small Coulombic forces of attraction.

(q+)(q-)

d*d

(2)

Arrange these compounds in order of

increasing attractions among ions.

HINT: consider the

SIZE

of the ions and the

CHARGE

of the ions.

KCl, Al

2

O

3

, CaO

Al

3+

O

2-

> Ca

2+

O

2-

> K

1+

Cl

1-

The Reason for the Formation of Ionic Compounds

Lies in the Electron Configurations of the Ions

1s 2s

2p

Li

F

Li

+

F

–

Same configuration as [He]

Same configuration

as [Ne]

Isoelectronic

: two or more species that

contain the same number of electrons.

Cations

(

positive

ions) become

isoelectronic with the

preceding

noble gas.

Anions

(

negative

ions) become

isoelectronic with the

following

noble gas.

Common Ions with Noble Gas

Configurations in Ionic Compounds

Sn: [Kr]5s2₄_d10₅_p2__Sn2+_{: [Kr]5}_s2₄_d10

(Pb)

Sn: [Kr]5s2₄_d10₅_p2__Sn4+_{: [Kr]4}_d10

Pb2+_{, Pb}4+

In+_{, In}3+

Tl+_{, Tl}3+

Bi3+_{, Bi}5+

In general, for a series of isoelectronic ions, the

size decreases

as the

nuclear charge increases

.

•Consider the isoelectronic ions:

O2-_{, F}-_{, Ne, Na}+_{, Mg}2+_{, and Al}3+

•How do the sizes of these ions vary?

Hint: consider the number of electrons and the number of protons in each.

•Since these ions are isoelectronic, the number of electrons is the same (10 e-_).

•The number of protons increases from 8 to 13 from O2- _to

Al3+_{, thus the 10 e}-_{experience a greater attraction as the}

positive charge on the nucleus increases, causing the ions to become smaller.

What Factors Influence the Stability and the

Structures of Binary Ionic Compounds?

•Ionic compounds form extended

3-D arrays (lattices) of oppositely charged ions.

(3)

Lattice energy

(follows Coulomb’s Law) is the change in

energy that takes place when separated gaseous ions are

packed together to form an ionic solid.

M

+

(

g

) + X

-

(

g

)



MX(

s

)

•

Coulomb

’

s Law E

a

Small ions with high ionic chargeshave

large

lattice energy.

Large ions with small ionic chargeshave small lattice energy.

(q+)(q-)

d*d

æ è ç ö_ø÷

Energy Changes of the Steps Involved in

the Formation of an Ionic Solid, MX

1. Sublimation of the solid metal

•

M(

s

)



M(

g

)

[endothermic]

2. Ionization of the metal atoms

•

M(

g

)



M

+

(

g

) + e



[endothermic]

3. Dissociation of the nonmetal

• 1

/2

X

2

(

g

)



X

(

g

)

[endothermic]

4. Formation of X_{ions in the gas phase:}

•

X

(

g

) + e





X



(

g

) [exothermic]

5. Formation of the solid MX

•

M

+

(

g

) + X



(

g

)



MX(

s

)

[quite

exothermic]

Sublimation

Ionization Ionization

Dissociation

Electron Affinity

Lattice Energy

Note the large lattice energy for MgO

(LEMgO = -3916 kJ/mol, where doubly charged ions are combining)

compared with that for NaF

(LENaF = -923 kJ/mol, where singly charged ions are combining)

Which would likely have the larger

negative value for its crystal lattice energy

CaO or KF and predict approximately

how much larger your choice will be over

the compound not selected?

1. CaO; approximately two times larger

2. CaO; approximately four times larger

3. KF; approximately two times larger

4. KF; approximately four times larger

How does a bonding force develop between

two identical atoms?

•3 energy terms

•p+_{- p}+_{repulsion (U)} •e-_{- e}-_{repulsion (U)} •p+_{- e}-_{attraction (F)}

•Bond Length

(4)

The Interaction of Two Hydrogen Atoms

zero point of E

(E = 0) Atoms are at

infinite separation.

Repulsive forces

Bond length: Distance at which the system as minimum energy.

Too far apart; weak attractive force; no bond Too close together, repulsive force dominate s

Attractive and repulsive forces are equal; most stable arrangement

(bond)

Octet

Rule =

8

Valence Electrons

Atoms will gain, lose or

share

sufficient electrons

to

achieve

the

same number of valence electrons as a

noble gas

(i.e.

8

electrons).

•

• • • •

F

• •

•

+

•

F

• • • •

F

• •

•

F

•

Each F counts both shared electrons to “feel” as though it has a Ne

configuration.

Octet

Rule =

8

Valence Electrons

Only two valence electrons

participate in the formation

of the F

2

bond.

F

•

F

•

• • • •

• •

or

• • • •

F

• •

•

F

•

F

•

F

•

• • • •

• •

Pairs of valence electrons

not involved in bonding

are called

lone pairs

.

Lewis structures

are

electron-dot structures

for molecules. They

show the location of

all valence e

–

.

Shared electron pairs are shown either as dashes

or as pairs of dots between two atoms. Unshared electron pairs are

called nonbonded electron pairs or lone pairs. Lone pairs are shown as pairs of dots on individual

atoms.

Covalent

bonds are formed when atoms

share

electrons

.

•If the atoms share

•2 electrons a single covalent bond is formed.

•4 electrons a double covalent bond is formed.

H O H

••

C with 8 e−

O with 8 e− O with 8 e−

one shared pair of electrons results in a single bond

••

O C O

••

2 shared pairs of electrons result in double bonds

O C O

= =

H O H

••

Covalent

bonds are formed when atoms

share

electrons

.

•

If the atoms share

•

6 electrons

a

triple

covalent bond is formed.

each N has 8 e−

3 shared pairs of electrons result in a triple bond

••

N N

••

N N

••

(5)

Bond length

is the distance between the

nuclei of two covalently bonded atoms

Single > Double > Triple

C-C

C=C

Relationship Between

Bond

Type

, Bond

Length

, and Bond

Energy

Bond Length: single > double > triple Bond Strength: triple > double > single

Shortest  Strongest  Longest  Weakest 

According to our current bond formation theory,

number of shared electron pairs, bond length, and

bond energy are related. Which of the following

correctly states the relationship among these three?

1. More shared

e

–

pairs ensures a shorter bond length

with a lower bond energy.

2. Fewer shared

e

–

pairs ensures a longer bond length

3. More shared

e

–

pairs ensures a longer bond length,

4. Fewer shared

e

–

pairs ensures a shorter bond length,

with a higher bond energy.

Steps for Drawing Lewis Structures of

Covalent Compounds

1. Count and total up all valence electrons.

•The valence electrons are in the outermost shell (the highest number

n) of an atom. The valence electrons are available to participate in bonding.

•# valence electrons = A-group number

•For every negative charge, add one electron to the total.

•For every positive charge, subtract one electron from the total.

Examples: CH4, H2O, CS2

2. Figure out which atom goes in the middle.

•Hydrogen never goes in the middle.

•Carbon usually goes in the middle.

•The least electronegative element goes in the middle.

•The element with only one atom usually goes in the middle.

3. Attach outer atoms to middle atom and subtract electrons used from total.

•Each bond ( –– ) equals two electrons ( : ).

4. Distribute electrons on outer atoms until outer atoms are full. Subtract e-_{used from total.}

•Remember the octet rule: Atoms will gain, lose, or share sufficient electrons to achieve the same number of valence electrons as a noble gas (i.e. 8 electrons).

•H only needs 2 electrons.

•All other atoms need 8 electrons to satisfy the octet rule.

5. If any electrons remain, place them on the middle atom.

•Nonmetal central atoms with empty d-orbitals available (Period 3 and higher) can accommodate more than 8 electrons.

•PCl5, SF4, ClF3, XeF4

6. If the middle atom does not have 8 electrons, make the outer atoms share their electrons with the middle atom.

•CO2, HCN

•Exceptions to this include:

•Be only needs 4 electrons, BeCl2

•B, Pb, Sn, only need 6 electrons, BF3

(6)

7. If a polyatomic ion, place the entire Lewis structure in brackets and write the charge as an upper right superscript.

•NH4+, SO42-, NO3-, I3-

8. Some molecules can have more than one Lewis structure (called resonance contributors).

•Sometimes the resonance contributors are equivalent. Other times you have to evaluate how good the resonance contributor is by assigning formal charges, The resonance contributor with the lowest separation of formal charge is the best one.

How to Draw a Lewis Structure, CO

2

1. Count and total up all valence e–_.

2. Determine which atom goes in the middle.

1. H never goes in the middle.

2. C usually goes in the middle.

3. The least electronegative element goes in the middle.

Attach all outer atoms and subtract the number of e–_used

from the total # of valence e–_.

Total 16valence electrons C 4 O 6 O 6 +

O _C O

O C O

12 electrons remaining C 4 O 6 O 6 16

4

+ –

How to Draw a Lewis Structure, CO

2

3. Distribute e

–

on the outer atoms

until the outer atoms have a full shell.

(8e

–

except H only needs 2e

–

)

Subtract the # of e

–

used from the total # of e

–

.

O C O

C 4 O 6 O 6 16 4 12

12 0

+ – –

How to Draw a Lewis Structure, CO

2 4. If any e–_{remain, place them on the central atom.}

5. If the central atom does not have 8e–

and no more valence e–_remain,

make the outer atoms share a lone pair(s) with the central atom.

O C O Octet

filled Octet not Octet filled

filled

O C O

O C O O C O

How to Draw a Lewis Structure for a

Polyatomic Ion

If a negative ion, add 1e–_{for each negative charge.}

If a positive ion, subtract 1e–_{for each positive charge.}

NH4+ 9 – 1 = 8 valence electrons

OH–₇_{+ 1}_{= 8 valence electrons}

CO32– 22 + 2 = 24 valence electrons

How to Draw a Lewis Structure for a

Polyatomic Ion

6. If a polyatomic ion, place the entire structure in brackets and

write the charge of the ion

as an upper right superscript.

(7)

Draw the Lewis Structure for ClO

₃–

If a negative ion, add 1e–_{for each negative charge.}

If a positive ion, subtract 1e–_{for each positive charge.}

Cl: 1 x 7 = 7 e− O: 3 x 6 = 18 e− Charge: 1 e− Total: 26 valence e−

How to Draw a Lewis Structure for a

Polyatomic Ion

6. If a polyatomic ion, place the entire structure in brackets and

write the charge of the ion

as an upper right superscript.

How many Covalent Bonds will an Atom Form?

•Atoms with one, two, or three valence e–_{form one, two, or}

three bonds, respectively.

•Atoms with four or more valence e–_{form enough bonds to}

give an octet (8e–_).

Number of bonds + Number of lone pairs = 4

Multiple Bonds

•A single bond contains two electrons in one 2e–_bond.

•A double bond contains four electrons in 2 2e-_bonds.

•A triple bond contains six electrons in 3 2e-_bonds.

O O

N N

How to Draw a Lewis Structure, C

₂

H

₄ 1. Count and total up all valence e–_.

2. Determine which atom goes in the middle.

1. H never goes in the middle.

2. C usually goes in the middle.

3. The least electronegative element goes in the middle.

Attach all outer atoms and subtract the number of e–_used

from the total # of valence e–_.

2 C x 4 e−_{= 8 e}−

4 H x 1 e−_{= 4 e}−

12e− total

C H

H C H H

How to Draw a Lewis Structure, C

2

H

2

3. Distribute e

–

on the outer atoms

until the outer atoms have a full shell.

(8)

How to Draw a Lewis Structure, C

2

H

2 4. If any e–_{remain, place them on the central atom.}

5. If the central atom does not have 8e–

and no more valence e–_remain,

make the outer atoms share a lone pair(s) with the central atom.

Each C now has an octet.

Answer C H

H C H H C

H H

C H H

Draw a Lewis structure for SiF

₄

SiF

4

=

32

valence electrons

F Si F Step 1

F F

Step 2 F Si F

F F

Step 3 F Si F

F F

Step 4 F Si F

F F

Draw a Lewis structure for SO

₃

SO

3

=

24

valence electrons

O S O Step 1

O

Step 2 O S O

O

Step 3 O S O

O

Step 4 O S O

O

Draw a Lewis structure for CN

–

When drawing Lewis structures for polyatomic ions: •Add one e−_{for each negative charge.}

•Subtract one e− _{for each positive charge.}

For CN– :

C N 1 C x 4 e−_{= 4 e}−

1 N x 5 e−_{= 5 e}−

–1 charge = 1 e−

10 e−_total

All valence e−

are used, but

Clacks an octet.

−

Each atom has an octet.

Answer

C N C N

Exceptions to the Octet Rule

•Most of the common elements generally follows the octet rule.

•H is a notable exception, because it needs only two e–_.

•Elements in group 3A do not have enough valence e–_{to form an}

octet in a neutral molecule.

only 6 e− on B B F

F F

Exceptions to the Octet Rule

•Elements in the third row and higher have empty d

orbitals available to accept more than 8 e–_.

•Thus, elements such as Pand Smay have more than 8e– around them.

(9)

Drawing Resonance Structures

•Resonance structures are two Lewis structures having

the same arrangement of atoms but a different arrangement of electrons.

•Two resonance structures of HCO3–:

•NeitherLewis structure is the true structure of HCO3–.

Drawing Resonance Structures

•The true structure is a hybridof the resonance structures.

•Resonance stabilizesa molecule by spreading out lone pairs and a electron pairs in multiple bonds over a larger region of space.

•A molecule or ion that has two or more resonance structures is resonance – stabilized.

Drawing Resonance Structures for NO

₃–

•The correct Lewis Structure is not any of the three possible resonance contributors, but is instead an average of the three

structures.

•Instead of two single N–O bonds and one double N=O bond, all three NO bonds are considered to be 1 1/3 bonds, that is, the electrons of the double bond are delocalized and can therefore move around the entire molecule.

N

O

N

O

N

O

Formal Charge Can be Used to Determine

the Most Plausible Lewis Structure When

More Than One Possibility Exists for a Compound

When several Lewis structures are possible, assign formal charges (a bookkeeping method) to all atoms in the compound.

The best Lewis structure is the one in which:

1. All atoms in the molecule have a formal charge of zero.

2. Small formal charges are preferred to large formal charges.

3. Any negative charges reside on the more electronegative atoms.

Guidelines for Assigning

Formal Charge

Remember, formal charge is a hypothetical charge on an atom. It is just a bookkeeping method used to evaluate

which resonance contributor best describes the actual bonding in the molecule or ion.

To determine associated electrons:

1) All the atom’s nonbonding electrons are associated with the atom.

2) Half the atom’s bonding electrons are associated with the atom. Formal charge = valence electrons – associated electrons

Determine the formal charges on each oxygen

atom in the ozone molecule (O

3

).

O O O

= − ••

••

•• ••

e−_{associated with atom}

Valence e−

Difference (formal charge)

6 6 6

6 5 7

0 +1 −1

4 unshared + 4 shared = 6 e−

2

(10)

Based on formal charge, identify the best and

the worst structures for the isocyanate ion

4 5 6

−2 +1 0

Ve−

Ae−

FC

4 5 6 4 5 6

6 4 6 5 4 7 7 4 5

−1 +1 −1 −3 +1 +1

Best structure:

Small formal charges Formal charges more consistent with electronegativities

Worst structure:

Large formal charges Formal charges inconsistent with electronegativities

Summary: Rules Governing Formal Charge

To calculate the formal charge on an atom:

1.

Take the sum of the lone pair electrons and

one-half the shared electrons.

2.

Subtract the number of associated electrons from

the number of valence electrons on the free,

neutral atom.

3.

The sum of the formal charges of all atoms in a

given molecule or ion must equal the overall

charge on that species.

Using Formal Charge To

Evaluate Resonance Structures

Draw the 3 possible Lewis structures of N2O,

assign formal charge to all atoms, and evaluate the structures to determine which one

most closely resembles the bonding in N2O.

Predict the ordering of the C–O bond lengths in CO, CO2, and CO32–.

Hint: draw Lewis structures,

assign formal charges if necessary, and remember that single bonds > double > triple.

Electronegativity is a measure of an

atom’s attraction for e

–

in a bond.

Electronegativity and Bond Polarity

If the electronegativities of two bonded atoms are equal

or similar, then the electrons are shared equally. These bonds are called

nonpolar covalent bonds.

If the electronegativities of two bonded atoms are different,

then the electrons are

unequally shared. These bonds are called

polar covalent bonds.

Electron density maps show

the distributions of charge.

Electrons spend a lot of time in red and very little time in blue.

Electrons are shared equally

nonpolar covalent

Electrons are not shared equally and

are more likely to be associated with F

polar covalent

Electrons are not shared but rather transferred from Na to F

(11)

Dipole Moment and Partial Charges

Regions where electrons

spend little time Regions where electrons spend a lot of time

An arrow is used to indicate the direction of electron shift in polar covalent molecules.

The consequent charge separation can be represented as:

Deltas (δ) denote a partial positive or negative charge.

H− F

••

H− F

••

δ+ δ−

The

bond polarization

(shift in the average

bonding-electron location) will be

toward

the atom with the

higher

electronegativity

.

H (EN = 2.1), F (EN = 4.0)

e–_{are pulled toward}_{the F, the more electronegative element; this is}

indicated by the symbol d–

e– _{are pulled away}_{from H, the less electronegative element; this is}

indicated by the symbol d+

H F

d

+

d



Greek lowercase delta,

d, means “_partial”

When the field is turned on, the molecules tend to line up with their negative ends toward the positive pole and their

positive ends towards the negative pole.

The Effect of an Electric Field on HF Molecules

When no electric field is present, the HF molecules are randomly oriented.

d, delta, is used to indicate a fractional charge.

Bond Polarity and Dipole Moments

When a molecule has two different ends like HF it is called

polar, meaning 2 poles, and is said to have a dipole moment.

H––F

d

+

d

-

Polar covalent bonds have a separation of

centers of positive and negative charge, an

asymmetrical charge distribution.

e

-

rich

region

e

-

poor

region

d

+

d

-

Which of the following bonds are polar?

Identify the negative and positive ends of each bond

by using

d

–

and

d

+

.

a. I—F, I (EN = 2.5), F (EN = 4.0) so the bond polarization is toward the F

I—F d– d+

b. O—H, O (EN = 3.5), H (EN = 2.1) so the bond polarization is toward the O

O—H d– d+

(12)

The Three Possible

Types of Bonds

Ionic

:

Electrons are completely lost or gained by the atoms to give oppositely charged ions which are attracted to

one another in an ionic bond.

Nonpolar Covalent

: covalent bonds in which the electrons are shared

equally; have a symmetrical charge distribution.

Polar Covalent

: Covalent bonds in which the electrons are not shared