chem281- chapter 3: covalent bonding bonding theories

8/3/2019 Chem281- Chapter 3: Covalent Bonding Bonding Theories

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CHEM281. Chapter 3 Covalent Bonding

Bonding Theories

1. VSEPR Theory 2. Valence Bond theory (with hybridization)

3. Molecular Orbital Theory ( with molecualr orbitals)

To date, we have looked at three different theories of molecular boning. They are the VSEPR Theory (with Lewis Dot Structures), theValence Bond theory (with hybridization) and Molecular OrbitalTheory. A good theory should predict physical and chemical propertiesof the molecule such as shape, bond energy, bond length, and bondangles.Because arguments based on atomic orbitals focus on thebonds formed between valence electrons on an atom, they are often

said to involve a valence-bond theory. The valence-bond model can'tadequately explain the fact that some molecules contains twoequivalent bonds with a bond order between that of a single bond anda double bond. The best it can do is suggest that these molecules aremixtures, or hybrids, of the two Lewis structures that can be written for these molecules.

This problem, and many others, can be overcome by using a moresophisticated model of bonding based on molecular orbitals.Molecular orbital theory is more powerful than valence-bond theorybecause the orbitals reflect the geometry of the molecule to which theyare applied. But this power carries a significant cost in terms of theease with which the model can be visualized. One model does not describe all the properties of molecular bonds.Each model desribes a set of properties better than the others. Thefinal test for any theory is experimental data.

Introduction to Molecular Orbital Theory The Molecular Orbital Theory does a good job of predicting elctronicspectra and paramagnetism, when VSEPR and the V-B Theories don't.The MO theory does not need resonance structures to describemolecules, as well as being able to predict bond length and energy.The major draw back is that we are limited to talking about diatomicmolecules (molecules that have only two atoms bonded together), or

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the theory gets very complex.

The MO theory treats molecular bonds as a sharing of electronsbetween nuclei. Unlike the V-B theory, which treats the electrons aslocalized baloons of electron density, the MO theory says that the

electrons are delocalized. That means that they are spread out over the entire molecule. Forming Molecular Orbitals Molecular orbitals are obtained by combining the atomic orbitals on theatoms in the molecule. Consider the H

2molecule, for example. One of

the molecular orbitals in this molecule is constructed by adding themathematical functions for the two 1s atomic orbitals that cometogether to form this molecule. Another orbital is formed by subtracting

one of these functions from the other, as shown in the figure below.

One of these orbitals is called a bonding molecular orbital becauseelectrons in this orbital spend most of their time in the region directlybetween the two nuclei. It is called a sigma ( ) molecular orbitalbecause it looks like an s orbital when viewed along the H-H bond.Electrons placed in the other orbital spend most of their time away

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from the region between the two nuclei. This orbital is therefore anantibonding, or sigma star ( *), molecular orbital.

The bonding molecular orbital concentrates electrons in the regiondirectly between the two nuclei. Placing an electron in this orbitaltherefore stabilizes the H

2molecule. Since the * antibonding

molecular orbital forces the electron to spend most of its time awayfrom the area between the nuclei, placing an electron in this orbitalmakes the molecule less stable. The MO Theory has five basic rules:

1. The number of molecular orbitals = the number of atomic orbitalscombined

2. Of the two MO's, one is a bonding orbital (lower energy) and one isan anti-bonding orbital (higher energy)

3. Electrons enter the lowest orbital available 4. The maximum # of electrons in an orbital is 2 (Pauli Exclusion

Principle) 5. Electrons spread out before pairing up (Hund's Rule)

Calculating Bond Order

In molecular orbital theory, we calculate bond orders by assuming thattwo electrons in a bonding molecular orbital contribute one net bondand that two electrons in an antibonding molecular orbital cancel theeffect of one bond. We can calculate the bond order in the O

2

molecule by noting that there are eight valence electrons in bondingmolecular orbitals and four valence electrons in antibonding molecular

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orbitals in the electron configuration of this molecule. Thus, the bondorder is two.

Although the Lewis structure and molecular orbital models of oxygen

yield the same bond order, there is an important difference betweenthese models. The electrons in the Lewis structure are all paired, buttherecould be unpaired electrons in the molecular orbital description of a molecule. As a result, we can test the predictions of these theoriesby studying the effect of a magnetic field on certain molecules.

Homo Nuclear Diatmic Molecules Molecular Orbitals of the Second Energy Level (1s and 2s only) Molecular Orbitals for Period 1 Diatomic Molecules

Electrons are added to molecular orbitals, one at a time, starting withthe lowest energy molecular orbital. The two electrons associated witha pair of hydrogen atoms are placed in the lowest energy, or bonding,molecular orbital, as shown in the figure below. This diagram suggeststhat the energy of an H

2molecule is lower than that of a pair of

isolated atoms. As a result, the H2

molecule is more stable than a pair

of isolated atoms.

We can put the Molecular Orbital Theory to use!! Would you predictthat dilithium or diberylium is more likely to form, based on the diagrambelow?

The answer is dilithium because it has a bond order of 1 which isstable and diberylium has a BO of 0 which is unstable and thereforewill not form.

Using the Molecular Orbital Model to Explain WhyHe

2MoleculeDo Not Exist

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This molecular orbital model can be used to explain why He2

molecules don't exist. Combining a pair of helium atoms with 1s2

electron configurations would produce a molecule with a pair of electrons in both the bonding and the * antibonding molecular

orbitals. The total energy of an He2 molecule would be essentially thesame as the energy of a pair of isolated helium atoms, and there wouldbe nothing to hold the helium atoms together to form a molecule.

The fact that an He2

molecule is neither more nor less stable than a

pair of isolated helium atoms illustrates an important principle: Thecore orbitals on an atom make no contribution to the stability of themolecules that contain this atom. The only orbitals that are important in

our discussion of molecular orbitals are those formed when valence-shell orbitals are combined. Molecular Orbitals for Period 2 Diatomic Molecules and Li

2and Be

2

(does not exsit)

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Homo Nuclear Diatmic Molecules Molecular Orbitals of the Second Energy Level (2s and 2p

together) Molecular Orbitals of the Second Energy Level

The 2s orbitals on one atom combine with the 2s orbitals on another toform a

2sbonding and a

2s* antibonding molecular orbital, just like

the1s

and1s

* orbitals formed from the 1s atomic orbitals. If we

arbitrarily define the Z axis of the coordinate system for the O2

molecule as the axis along which the bond forms, the 2 pz

orbitals on

the adjacent atoms will meet head-on to form a2 p

bonding and a2 p

*

antibonding molecular orbital, as shown in the figure below. These arecalled sigma orbitals because they look like s orbitals when viewedalong the oxygen-oxygen bond.

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The 2 px

orbitals on one atom interact with the 2 px

orbitals on the other

to form molecular orbitals that have a different shape, as shown in the

figure below. These molecular orbitals are called pi ( ) orbitals becausethey look like p orbitals when viewed along the bond. Whereas and *orbitals concentrate the electrons along the axis on which the nuclei of the atoms lie, and * orbitals concentrate the electrons either above or below this axis.

The 2 px

atomic orbitals combine to form ax

bonding molecular orbital

and ax* antibonding molecular orbital. The same thing happens when

the 2 py

orbitals interact, only in this case we get ay

and ay*

antibonding molecular orbital. Because there is no difference betweenthe energies of the 2 p

xand 2 p

yatomic orbitals, there is no difference

between the energies of thex

andy

or thex* and

y* molecular

orbitals.

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The interaction of four valence atomic orbitals on one atom (2s, 2 px,

2 py

and 2 pz) with a set of four atomic orbitals on another atom leads to

the formation of a total of eight molecular orbitals:2s

,2s

*,2 p

,2 p

*,x,

y,

x*, and

y*.

There is a significant difference between the energies of the 2s and 2 p orbitals on an atom. As a result, the

2sand *

2sorbitals both lie at

lower energies than the2 p

,2 p

*,x,

y,

x*, and

y* orbitals. To sort out

the relative energies of the six molecular orbitals formed when the 2 p atomic orbitals on a pair of atoms are combined, we need tounderstand the relationship between the strength of the interactionbetween a pair of orbitals and the relative energies of the molecular orbitals they form. Because they meet head-on, the interaction between the 2 p

zorbitals is

stronger than the interaction between the 2 px

or 2 py

orbitals, which

meet edge-on. As a result, the2 p

orbital lies at a lower energy than

thex

andy

orbitals, and the2 p

* orbital lies at higher energy than the

x* and

y* orbitals, as shown in the figure below.

Unfortunately an interaction is missing from this model. It is possiblefor the 2s orbital on one atom to interact with the 2 p

zorbital on the

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other. This interaction introduces an element of s-p mixing, or hybridization, into the molecular orbital theory. The result is a slightchange in the relative energies of the molecular orbitals, to give thediagram shown in the figure below. Experiments have shown that O

2

and F2

are best described by the model in the figure above, but B2, C

2,

and N2

are best described by a model that includes hybridization, as

shown in the figure below.

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Construct a Molecular orbital diagram for the O

2molecule.

There are six valence electrons on a neutral oxygen atom andtherefore 12 valence electrons in an O

2molecule. These electrons are

added to the diagram in the figure below, one at a time, starting withthe lowest energy molecular orbital.

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Because Hund's rules apply to the filling of molecular orbitals,molecular orbital theory predicts that there should be two unpaired

electrons on this molecule one electron each in thex* and

y*

orbitals.

When writing the electron configuration of an atom, we usually list theorbitals in the order in which they fill.

We can write the electron configuration of a molecule by doing the

same thing. Concentrating only on the valence orbitals, we write theelectron configuration of O2

as follows.

Molecular Orbital diagram fro N2

Pb: [Xe] 6s2

4f 14

5d 10

6 p2

O2:

2s

2

2s*2

2 p

2

x

2

y

2

x*1

y*1

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Bond Order

The number of bonds between a pair of atoms is called the bondorder . Bond orders can be calculated from Lewis structures, which arethe heart of the valence-bond model. Oxygen, for example, has a bondorder of two.

When there is more than one Lewis structure for a molecule, the bondorder is an average of these structures. The bond order in sulfur

dioxide, for example, is 1.5 the average of an S-O single bond inone Lewis structure and an S=O double bond in the other.

In molecular orbital theory, we calculate bond orders by assuming thattwo electrons in a bonding molecular orbital contribute one net bondand that two electrons in an antibonding molecular orbital cancel theeffect of one bond. We can calculate the bond order in the O

2

molecule by noting that there are eight valence electrons in bondingmolecular orbitals and four valence electrons in antibonding molecular orbitals in the electron configuration of this molecule. Thus, the bondorder is two.

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Although the Lewis structure and molecular orbital models of oxygenyield the same bond order, there is an important difference betweenthese models. The electrons in the Lewis structure are all paired, butthere are two unpaired electrons in the molecular orbital description of the molecule. As a result, we can test the predictions of these theories

by studying the effect of a magnetic field on oxygen.

Atoms or molecules in which the electrons are paired are diamagneticrepelled by both poles of a magnetic. Those that have one or more

unpaired electrons are paramagnetic attracted to a magnetic field.Liquid oxygen is attracted to a magnetic field and can actually bridgethe gap between the poles of a horseshoe magnet. The molecular orbital model of O

2is therefore superior to the valence-bond model,

which cannot explain this property of oxygen. Look at the following MO diagrams for some of the period twoelements. Can you tell which molecules are paramagnetic? Whichmolecules have the highest bond energy, which has the lowest? Ranksingle, double, and triple bonds in order of bond energy and bondlength. (hint a BO of 1 is a single bond, 2 a double...) Delocalized Bonding Finaly, it was mentioned earlier that the MO Theory did not needresonance structures to explain anything. Because the MO theory

holds that electrons are not held to only one position. Instead they arespread across the entire molecule. Below is a picture of Benzene andOzone. You can see that rather than having two resonance structures,we can picture one structure with the electrons dispersed over theentire molecule.

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Molecular Orbitals for Heteronuclear Diatomic Molecules

The most striking molecular orbital difference between a heteronuclear

diatomic, such as HF, and the homonuclear diatomics is that theorbitals are no longer have equally density on each atom. Note that the "1-sigma" and "2-sigma" molecular orbitals of Hydrogen

Fluoride are mostly derived from the Fluorine 2s and 2p atomic orbitalsrespectively. Qualitatively, the high fluorine character of these orbitalsis a consequence of the high electronegativity of fluorine as comparedto hydrogen. Mathematically, the molecular orbitals have largecoefficients on Fluorine and small coefficients on Hydrogen. The 1 pi

orbitals are non-bonding and Fluorine 2p in character. Finally, the anti-bonding "3-sigma" orbital is primarily H 1s in character. The use of nodes to give a general idea of the energy ordering still

works. "1-sigma" has no nodes, "2-sigma" and "1-pi" have one node,and "3-sigma" has two nodes. Carbon monoxide CO molecular orbital diagram

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Iodine monochloride ICl molecular orbital diagram

Lewis Theory Octet Rule All elements except hydrogen ( hydrogen have a duet of electrons)have octet of electrons once they from ions and covalent compounds. The Lewis dot symbols for atoms and ions shows how manyelectrons are need for a atom to fill the octet. Noramlly there are octetof electrons on most monoatomic ions Basic rules drawing Lewis dot symbols: 1. Draw the atomic symbol.

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2. Treat each side as a box that can hold up to two electrons. 3. Count the electrons in the valence shell. Start filling box - don’t make pairs unless you need to.

Elements try to complete their valence shell to achieve noble-gas electronconfigurations. The stability of noble/inert gases must be due to their filledvalence shell. Elements either loose/gain or share electrons for this purpose.Electron transfer is associated with the formation of ionic compounds.Covalent compounds are formed when electrons are shared between atoms.A Lewis symbol is a symbol in which the electrons in the valence shell of anatom or simple ion are represented by dots placed around the letter symbol of the element. Each dot represents one electron.

A covalent bond is a chemical bond formed by the sharing of a pair of electrons between two atoms.The Lewis structure of a covalent compound or polyatomic ion shows howthe valence electrons are arranged among the atoms in the molecule to show

the connectivity of the atoms.

Hydrogen

Oxygen

Chlorine

Chloride ion

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Instead of using two dots to indicate the two electrons that comprise thecovalent bond, a line is substituted for the two dots that represent the twoelectrons.

Below is shown the Lewis structure for water. Two hydrogens (H) areseparately covalently bonded to the central oxygen (O) atom. The bonding electrons are indicated by the dashes between the oxygen (O) and eachhydrogen (H) and the other two pairs of electrons that constitute oxygensoctet, are called non-bonding electrons as they are not involved in a covalent

bond.

RULES FOR LEWIS STRUCTURES

1. Determine whether the compound is covalent or ionic. If covalent, treat the

entire molecule. If ionic, treat each ion separately. Compounds of lowelectronegativity metals with high electronegativity nonmetals (DEN > 1.6)are ionic as are compounds of metals with polyatomic anions. For amonoatomic ion, the electronic configuration of the ion represents the correctLewis structure. For compounds containing complex ions, you must learn torecognize the formulas of cations and anions.2. Determine the total number of valence electrons available to the moleculeor ion by:(a) summing the valence electrons of all the atoms in the unit and

(b) adding one electron for each net negative charge or subtracting oneelectron for each net positive charge. Then divide the total number of available electrons by 2 to obtain the number of electron pairs (E.P.) available. 3. Organize the atoms so there is a central atom (usually the leastelectronegative) surrounded by ligand (outer) atoms. Hydrogen is never thecentral atom.

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4. Determine a provisional electron distribution by arranging the electron pairs(E.P.) in the following manner until all available pairs have been distributed:a) One pair between the central atom and each ligand atom.

b) Three more pairs on each outer atom (except hydrogen, which has noadditional pairs), yielding 4 E.P. (i.e., an octet) around each ligand atom when

the bonding pair is included in the count.c) Remaining electron pairs (if any) on the central atom.

Examples: Consider carbon dioxide CO2

1. The first step in drawing Lewis structures is to determine the number of electrons to be used to connect the atoms. This is done by simply addingup the number of valence electrons of the atoms in the molecule.

carbon (C) has four valence electrons x 1 carbon = 4 e-

oxygen (O) has six valence electrons x 2 oxygens = 12 e-

There are a total of 16 e-

to be placed in the Lewis structure. 1. Connect the central atom to the other atoms in the molecule with single

bonds.Carbon is the central atom, the two oxygens are bound to it and electronsare added to fulfill the octets of the outer atoms.

2. Complete the valence shell of the outer atoms in the molecule.

3. Place any remaining electrons on the central atom.There are no more electrons available in this example.

4. If the valence shell of the central atom is complete, you have drawn anacceptable Lewis structure.Carbon is electron deficient - it only has four electrons around it. This is

not an acceptable Lewis structure.

5. If the valence shell of the central atom is not complete, use a lone pair onone of the outer atoms to form a double bond between that outer atomand the central atom. Continue this process of making multiple bonds

between the outer atoms and the central atom until the valence shell of the central atom is complete.

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There are two possible Lewis structures for this molecule. Each has the samenumber of bonds. We can determine which is better by determining which hasthe least formal charge. It takes energy to get a separation of charge in themolecule (as indicated by the formal charge) so the structure with the least

formal charge should be lower in energy and thereby be the better LewisstructureThe two possible Lewis structures are shown below. They areconnected by a double headed arrow and placed in brackets. The non-zeroformal charge on any atoms in the molecule have been written near the atom.

The two structures differ only in the arrangement of the valence electrons inthe molecule. No atoms have been moved. These are called resonance

structures. The better Lewis structure or resonance structure is that which hasthe least amount of formal charge. If the central atom formal charge iszero or is equal to the charge on the species, the provisional electrondistribution from (4) is correct. Calculate the formal charge of the ligandatoms to complete the Lewis structure. If the structure is not correct, calculate the formal charge on each of the ligand atoms. Then to obtain the correct structure, form a multiplebond by sharing an electron pair from the ligand atom that has themost negative formal charge. a) For a central atom from the second (n = 2) row of the periodic tablecontinue this process sequentially until the central atom has 4 E.P. (anoctet).

becomes

The central atom is still electron deficient, so share another pair

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b) For all other elements, continue this process sequentially until theformal charge on the central atom is reduced to zero or two doublebonds are formed. Recalculate the formal charge of each atom to complete the Lewisstructure.

Draw the Lewis structures for the following molecules: H

2O (H

2S), CH

4(SiH

4), CCl

4(CF

4, CBr

4), CO

2, NH

3(PH

3), PCl

3

(PF3, NCl

3)

b) CH4

i) Sum of valence electrons:

4 ( from one C) + 4 (from four H) = 8 electrons = 4electron pairs

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SiH4 has the same Lewis Structure as CH4 since Si and C are in the

same group.

c) CCl4

i) Valence electrons: 4 (from one carbon)+ 28 (7 from eachchlorine) = 32 = 16 electron pairs ii) Central atom iii) Octet on C is already complete iv) Count electrons 4 x 3 lone pairs = 12 pairs 4 x 4 bond pairs = 4 pairs

16 electron pairs Lewis Structure of CCl4. CF

4and CBr

4have similar structures.

d) CO2

i) Valence electrons: 4 + 2 x 6 = 16 ( 8 pairs) ii) Central atom iii) Give octet to carbon

Try to fill octet to O

iv) Count electrons: 4 bond pairs = 4 pairs 4 lone pairs = 4 pairs 8 electron pairs

Lewis Structure of CO2

ii) Central atom is C iii) Octet on C atom and duet on H are already complete.

iv) Count valence electrons on the Lewis structure: bond pairs = 2 x 4 = 8 = 4 electron pairs

lone pairs = 0

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e) NH3

i) Valence electrons: 5 + 3 x 1 = 8 = 4 electron pairs ii) N is the central atom: iii) Add terminal Hs vi) Give octet to N

Check duet on Hs v) Count electrons:

3 bond pairs = 3 pairs 1 lone pairs = 1 pairs 4 electron pairs

Lewis Structures of NH3, PH

3has the same Lewis Structure

f) PCl3

i) Valence electrons: 5 + 3 x 7 = 26 (13 pairs) ii) Central atom is P iii) Connect to terminal atoms to central atom vi) Give octet to P

give octets to Cl v) Count electron pairs:

3 bond pairs = 3 pairs 1 + 3 x 3 = 10 lone pairs = 10 pairs

13 pairs

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PF3

and NCl3

have similar Lewis Structures.

10) How many lone pairs and bond pairs of electrons are foundaround the central atoms of the following molecules? CH

4, NH

3,

PCl3, H2O

a) Lewis Structure of CH4

i) 4 + 1 x 4 = 8 = pairs ii) C central

iii) octet to C, duet to H

iv) count pairs 4 pairs 4 bond pairs. Bond pair: electron pair shared between two atoms. Lone pair: electron pair found on a single atom.

CH4

has only 4 bond pairs on central C.

b) Lewis Structure of NH3

i) Valence electrons: 5 + 3 x 1 = 8 = 4 electron pairs

ii) N is the central atom: iii) Add terminal Hs vi) Give octet to N

Check duet on Hs v) Count electrons: 3 bond pairs = 3 pairs

1 lone pairs = 1 pairs

4 electron pairs Lewis Structures of NH

3, PH

3has the same Lewis Structure

3 bond pairs and 1 lone pair are found on central N.

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c) PCl3

i) Valence electrons: 5 + 3 x 7 = 26 = 13 electron pairs ii) P is the central atom: iii) Add terminal Cls

vi) Give octet to P add octet to Cls

v) Count electrons: 3 bond pairs = 3 pairs 13 lone pairs = 13 pairs

26 electron pairsLewis Structures of PCl3 Three bond pairs and 13 lone pairs,one on P and 12 on Cls.

d) H2O. Lewis structure of H

2O

Two bond pairs and two lone pairs are found oncentral O.

11) Draw the Lewis structures of the following molecules andshow that the electron count around the central atoms violate theoctet rule: BF

3(BH

3, AlCl

3), SF

6, PCl

5(AsCl

5)

a) BF3

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i) Count total valence electrons: 3 + 3 x 7 = 24 (12 electronpairs) ii) Central atom is B iii) Connect central atom terminal atoms: iv) 3 bonds pairs on B.

3 lone pairs on 3F 3 + 3 x 3 = 12 pairs

Boron do not have a octet, therfore, BF3

is an electron deficient

compound. BF3

violate octet rule.

Similarly BH3

has the Lewis Structure:

B has on 3 electron pairs. BH3

violates octet rule

b) SF6

i) Count total valence electrons:6 + 6 x 7 = 48 (24 pairs) ii) S is the central atom

sulfur already have excess of 4 electrons to octet iii) fill octet to Fs iv) count electrons:

S - 6 bond pairs = 6 6F - 3 x 6 lone pairs = 18

24 pairs

Similarly AlCl3

has the Lewis Structure:

Al has only 3 electron pairs. AlCl

3also violates octet rule.

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Sulfur violates octet rule, is an exception to octet rule having morethan eight electrons on S. S has 12 electrons.

c) PCl5

i) Count the total valence lectrons: 5 + 5 x 7 = 40 (20 pairs) ii) P is central atom: iii) P already have 5 electron pairs then give octets to Cls

iv) count electron pairs. 5 pairs on P = 5 5 x 3 pairs on Cl = 15

20 electron pairs

AsCl5

have similar structure. PCl5

and AsCl5

violates octet rule

having 10 electrons on P and As.

Resonance Structures

Necessity to draw several Lewis structures (resonancestructures) to adequately describe the structures of the following

molecules: CO3

2-, NO

3

-, NO

2

-

a) Lewis Structure of CO32-:

i) valence electrons:4 + 3 x 6 + 2(negative charge) = 24 (12 pairs) ii) central atom is carbon

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iii) fill octet to C; fill octet to O

One oxygen does not have an octet. Share lone pair on C withthe oxygens. The double bond could be on any other

oxygen atoms.

Therefore, there are three resonance structures for CO3

2-ion

b) NO3

-Lewis Structure:

i) valence electrons: 5 + 3 x 6 + 1(negative charge) = 24 (12 pairs)

iii) N is the central atom:

iii) Fill octet to N; Fill octet to Os

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one oxygen has only 6 electrons share lone pair on N with the oxygen

The double bond could be on any other oxygen atoms. Therefore,there are three resonance structures for the

NO3

-ion.

c) NO2

-Lewis Structure

i) valence electrons: 5 + 2 x 6 + 1(negative charge) = 18 (9 pairs)

ii) central atom is nitrogen. iii) Fill octet to N ; Fill octet to O

one oxygen has only six electrons share lone pair on N with oxygen.

The double bond could be on any of the two oxygen atoms.

Therefore, there are two resonance structures for NO2

-ion.

Predicting the Molecular Structure from Lewis Structure

Valence Shell Electron Pair Repulsion Theory VSEPR stands for Valence Shell Electron Pair Repulsion. It is a theory whichallows the user to predict the shapes of simple polyatomic molecules byapplying a set of straight forward rules. VSEPR Theory is one method that

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chemists use to predict the shapes of molecules. This theory predicts thatelectron pairs, whether involved in bonds or as non-bonding pairs, will adopt ageometry in which they maximize the distance from one another in order tominimize repulsions. This will result in a geometry with the lowest possibleenergy.

VSEPR theory states that each region of electron density on the central atomhas an impact on the final shape of a molecule, and these regions want to be asfar apart as possible. Why do these regions want to be as far apart as possible?Let us consider each region of electron density as a point charge. Since allthese regions have the same charge they should be repelled and the lowestinteraction would be when the point charges are as far apart as possible.Another question which must be answer is: How do I represent a structure?The answer to this question depends on what you want to show. If you want toemphasize the geometry of individual atoms in a molecule you would useeither stick or ball and stick structures. On the other hand, you want toemphasize the overall shape of the molecule, a space filling structure is moreappropriate.The ideal electronic symmetry of a molecule consisting of a centralatom surrounded by a number of substituents (bonded atoms and non-bonding electrons) is characteristic of the total number of substituents,and is determined solely by geometric considerations -- thesubstituents are arranged so as to maximize the distances amongst

them. VSEPR is useful for predicting the shape of a molecule whenthere are between 2 and 6 substituents around the central atom (thecase of one substituent is not discussed because it is trivial -- the onlypossible shape for such a molecule is linear). That means that thereare only five unique electronic geometries to remember. For eachelectronic geometry, there may be a number of different molecular geometries (the shape of the molecule when only bonded atoms, notnon-bonding electrons are considered). Molecular geometries arereally just special cases of the parent electronic geometry -- this will

hopefully be evident from the models shown on the pages linked to thisone. Since the molecular geometry is determined by how many bonding andnon-bonding electron groups surround the central atom, the first thingone needs to do is count how many of each there are. There is anotation that simplifies this bookkeeping:

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ABxE

y

The A represents the central atom, B represents the electron groupsthat form bonds to other atoms, and E represents the non-bondingelectron groups. The subscripts, x and y, indicate how many of each

kind are present.

Note that bonding "electron groups" does not necessarily imply singlebonds; it can mean double or triple bonds as well:

An incidental benefit of using the ABE notation is that it provides aconvenient way of remembering the hybridization at the central atom.The total number of substituents (bonding plus non-bonding groups) is

AB4

bonding groups: 4 non-bonding groups: 0

AB3

bonding gnon-bonding

CH4

methane

NH

amm

AB2

bonding groups: 2 non-bonding groups: 0

CO

2

Carbon Dioxide

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equal to the number of atomic orbitals that participate in the hybridorbital.

Displayed in the following table are the five most important electronicsymmetries. Each row in the table is linked to a page that shows thedifferent molecular symmetries possible for that electronic symmetry.

Predicting the Shapes of Molecules

There is no direct relationship between the formula of a compound andthe shape of its molecules. The shapes of these molecules can bepredicted from their Lewis structures, however, with a modeldeveloped about 30 years ago, known as the valence-shell electron-pair repulsion (VSEPR) theory.

The VSEPR theory assumes that each atom in a molecule will achieve

Molecule ABE representation # of substituents Hybridization

CH4

AB4 4 sp

3

NH3

AB3E 4 sp

3

H2O AB

2E

2 4 sp3

CO2

AB2 2 sp

SF6

AB6 6 sp

3d

2

I

3

-ion AB

2E

3 5 sp3d

Class Hybridization Electronic Symmetry

AB2 sp linear

AB3 sp2

trigonal planar

AB4 sp

3 tetrahedral

AB5 sp

3d trigonal bipyramidal

AB6 sp

3d

2 octahedral

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a geometry that minimizes the repulsion between electrons in thevalence shell of that atom. The five compounds shown in the figurebelow can be used to demonstrate how the VSEPR theory can beapplied to simple molecules.

There are only two places in the valence shell of the central atom inBeF

2where electrons can be found. Repulsion between these pairs of

electrons can be minimized by arranging them so that they point inopposite directions. Thus, the VSEPR theory predicts that BeF

2should

be a linear molecule, with a 180o

angle between the two Be-F bonds.

There are three places on the central atom in boron trifluoride (BF3)

where valence electrons can be found. Repulsion between these

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electrons can be minimized by arranging them toward the corners of an equilateral triangle. The VSEPR theory therefore predicts a trigonalplanar geometry for the BF

3molecule, with a F-B-F bond angle of

120o.

BeF2

and BF3

are both two-dimensional molecules, in which the atoms

lie in the same plane. If we place the same restriction on methane(CH

4), we would get a square-planar geometry in which the H-C-H

bond angle is 90o. If we let this system expand into three dimensions,however, we end up with a tetrahedral molecule in which the H-C-H

bond angle is 109o28'.

Repulsion between the five pairs of valence electrons on thephosphorus atom in PF

5can be minimized by distributing these

electrons toward the corners of a trigonal bipyramid. Three of thepositions in a trigonal bipyramid are labeled equatorial because they liealong the equator of the molecule. The other two are axial becausethey lie along an axis perpendicular to the equatorial plane. The angle

between the three equatorial positions is 120o, while the angle

between an axial and an equatorial position is 90o

.

There are six places on the central atom in SF6

where valence

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electrons can be found. The repulsion between these electrons can beminimized by distributing them toward the corners of an octahedron.The term octahedron literally means "eight sides," but it is the sixcorners, or vertices, that interest us. To imagine the geometry of anSF

6molecule, locate fluorine atoms on opposite sides of the sulfur

atom along the X , Y , and Z axes of an XYZ coordinate system.

Predict the molecular structure of the following molecules usingVSEPR theory: H

2

O, NH

3

, CO

2

, SF

6

, PCl

5

, XeF

4

a) Molecular Structure of H2O :

First get the Lewis structure of H2O, problem 9a.

2 bond pairs = 2 x 2 = 4 2 lone pairs = 2 x 2 = 4

Total 8 = 4 electron pairs (an Octet) Basic Structure: Arrangement of electron pairs arounf the central

atom Tetrahedral for H2O

Molecular Structure: Arrangement of atom or groups around

central atom. Ignore lone pairs and look at the basic structure Basic Structure Molecular Structure: Angular

Basic(electronic) structure Molecular structure

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b) molecular structure of NH3:

Lewis Structure:

Lewis Structure Basic(electronic) structureMolecular structure

tetrahedralpyramidal structure c) Molecular Structure CO

2

Lewis Structure of CO2

electron pairs but double bond is counted as single electron pair =2 pairs total

2 electron pairs linear basic structure Since there are no lonepairs on C basic

structure is same as molecular structure. Basic structure molecular structure

d) Molecular Structure of: SF6

Lewis Structure

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i) count electrons around central atom: 6 electron pairs

ii) 6 electron pairs given octahedral basic structure

iii) Since there are no lone pairs basic structure is same asmolecular structure.

Basic structure and molecular structure is the same since thereare no lone pairs on S. molecular structure of SF

6is octahedral

e) Molecular Structure of PCl5:

Lewis Structure:

i) Count electron pairs on the central atom: 5 electron pairs. ii) Basic structure is trigonal bipyramidal iii) There are no lone pairs on central P atom. Therefore, basic structure is same as molecular structure.

Molecular structure of PCl5

- Trigonal bipyramidal

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f) Molecular Structure of XeF4.

First get the Lewis structure i) Count the total valence lectrons: 8 + 4 x 7 = 36 (18 pairs) ii) Xe is central atom and connect Xe to terminal Fs iii) Give octets to Fs vi) Place extra electrons on Xe iv) count electron pairs.

6 pairs on Xe = 6 4 x 3 pairs on Fl = 12

18 electron pairs

Molecular Structure: i) Count electron pairs on central atom: 6 pairs ii) Basic structure is octahedral. Molecular

structure

The lone pairs occupy two opposite corners of the octahedron toavoid stronger repulsion. When you ignore lone pairs the molecular structure becomes a square plane.

Predicting Polarity of molecules

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H2O has a net dipole moment and SO

3has nodipole moment?

However, both the H-O and S-O bonds in these molecules have apolarity. Molecular Structure of H

2O is bent:

There is polarity alone O - H bonds:

Polarity adds like other mechanical forces. If there are equal

dipoles opposite (1800

) to each other they cancel. In water two dipoles are added up to give water molecule a

stronger net dipole. Therefore, water molecule is polar and have adipole moment.

Molecular structure of SO3:

First get the Lewis structure of SO3

i) Count the total valence lectrons: 6 + 3 x 6 = 24 (12 pairs) ii) S is central atom and connect S to terminal Os iii) Give octets to Os v) count electron pairs.

4 pairs on S = 4 8 pairs on Os = 8

12 electronpairs

Molecular structure of SO3: Trigonal planer

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There is a dipole along each S - O bond

They are at 1200

to each other. Therefore, they cancel out. Thereis no net dipole moment in SO

3molecule.

Predict the polarity (net dipole moment or zero dipole moment) of following molecules : H

2O, NH

3, CO

2, SO

3, SF

6, PCl

5, XeF

4

a) H

2

O: Lewis Structure of H

2

O:

Polarity adds like other forces. If there are equal dipoles opposite

(1800) to each other they cancel. In water two dipoles are added up to

give water molecule a stronger net dipole. Therefore, water molecule ispolar and have a dipole moment. Two bonds are at about 105

0and the

polarities adds up to a stronger dipole. H2O has a dipole moment.

b) NH3: Lewis Structure of NH

3:

Molecular Structure: Pyramidal N - H bond is polar:

Bonds are not distributed symmetrically. Therefore, ammonia has a dipole moment.

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c) CO2: Lewis Structure of CO

2:

Molecular Structure:Linear

C - O bond has a polarity: C - O However, bonds are aligned 1800

toeach other canceling each other. There is no dipole moment in themolecule.

d) SO3: Lewis Structure of SO

3:

i) Count the total valence lectrons: 6 + 3 x 6 = 24 (12 pairs) ii) S is central atom and connect S to terminal Os iii) Give octets to Os v) count electron pairs.

4 pairs on S = 4 8 pairs on Os = 8

12 electronpairs

Molecular structure of SO3: Trigonal planer

There is a dipole along each S - O bond

They are at 1200

to each other. Therefore, they cancel out. There is nonet dipole moment in SO

3molecule.

Molecular Structure of SO3: trigonal planer

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dipoles are arranged at 1200

to each other. Dipoles cancel and there is no dipole moment in the molecule. e) SF

6: Lewis Structure of SF

6

Molecular Structure of SF6: octahedral

S - F has a dipole. However, bonds are arranged symmetricallyopposite to each other. There is no dipole moment in the molecule.

f) PCl5: Lewis Structure of PCl5:

Molecular Structure: Trigonal bipyramid P - Cl: there is dipole moment in the bond

However P - Cl bond are arranged symmetrically around P cancelingeach other PCl

5molecule do not have a dipole.

e) XeF4: Lewis Structure of XeF

4:

Molecular Structure of XeF4: Square planer

Xe - F bond is polar. However bond are arranged symmetricallyopposite to each

other, cancelling the dipoles. There is no dipole moment in the

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molecule.

Predicting the hybridization of the central atom(s) in the following

molecules: H2O (H

2S), CH

4(SiH

4), CCl

4(CF

4, CBr

4), CO

2, NH

3

(PH3), PCl

3(PF

3, NCl

3)

a) H2O:

Basic structure: tetrahedral number of electrons pairs around oxygen is four: four pairs arrange tominimize electron pair repulsion in a tetrahedron. Four electron pairs

are in four sp3

hybridized orbitals in oxygen. The hybridization of the

oxygen atom is sp3. H

2S is similar to H

2O Hybridization of S atom is

sp3.

b) CH4:

Basic structure: tetrahedral number of electron pairs around the central atom is four.

The hybridization of the C atom of CH4

is sp3.

SiH4

- hybridization of Si atom is sp3.

c) CCl4:

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Basic structure: tetrahedral There are four electron pairs on C atom.

The hybridization of the carbon atom is sp3.

In CF4

and CBr 4

the hybridization of C is sp3.

d) CO2:

Basic structure: linear number of electron pairs 2 (what is really making directional

bonds)

Hybridization of the C atom is sp2.

e) NH3:

Four electron pairs around central N atom. The hybridization of atomic

orbitals of N atom is sp3.

PH3: sp3 hybridization of P.

f) PCl3:

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Four electron pairs around central P atom. Hybridization of P is sp3.

PF3, NCl

3: hybridization of central atom is sp

3

Inclusion of atomic orbitals into Lewis Structure and VSEPR

Theory

Valence-Bond Theory

Covalent bonding is the result of orbital overlap, the overlap of atomicorbitals from two atoms, and the sharing of electrons in the region of overlap. This is the central idea of Valence Bond Theory, whichdescribes how covalent bonding occurs. Molecular Orbital Theory, onthe other hand, postulates the combination of atomic orbitals of different atoms to form molecular orbitals, so that the electrons inthem belongs to the molecule as a whole. These two have strengths

and weaknesses that are complementary. VB theory is easy tovisualize, but MO theory gives a better descriptions of bond energiesand magnetic properties. The Heitler-London model of covalent bondswas the basis of the valence-bond theory. The last major step in theevolution of this theory was the suggestion by Linus Pauling that

atomic orbitals mix to form hybrid orbitals, such as the sp, sp2, sp

3,

dsp3, and d

2sp

3orbitals.

Covalent bonding and orbital overlap of hybridized orbitals of the centralatom Hybridization in Simple Molecules Hybrid orbitals Hybrid orbitals are combinations of different types of atomic orbitalson the same atom. In forming a molecule, the combination of valenceorbitals of the center atom and the other atoms, which gives amaximum overlap will be favored. This is the original idea to mix different atomic orbitals of the center atom.

sp hybrid orbitals sp hybridization

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Hybridization involving d orbitals It is easy to apply the valence-bond theory to some coordination

complexes, such as the Co(NH3)6

3+ion. We start with the electron

configuration of the transition- metal ion.

Co3+

: [Ar] 3d 6

We then look at the valence-shell orbitals and note that the 4s and 4 p orbitals are empty.

Co3+

: [Ar] 3d 6

4s0

4 p0

Concentrating the 3d electrons in the d xy, d xz, and d yz orbitals in thissubshell gives the following electron configuration.

sp2

hybrid orbitals

sp3

hybrid orbitals

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The 3d x

2

-y

2, 3d

z

2, 4s, 4 p

x, 4 p

yand 4 p

zorbitals are then mixed to form

a set of empty d 2sp

3orbitals that point toward the corners of an

octahedron. Each of these orbitals can accept a pair of nonbondingelectrons from a neutral NH

3

molecule to form a complex in which the

cobalt atom has a filled shell of valence electrons.

At first glance, complexes such as the Ni(NH3)6

2+ion seem hard to

explain with the valence-bond theory. We start, as always, by writing

the configuration of the transition-metal ion. Ni

2+: [Ar] 3d

8

This configuration creates a problem, because there are eightelectrons in the 3d orbitals. Even if we invest the energy necessary topair the 3d electrons, we can't find two empty 3d orbitals to use to form

a set of d 2sp

3hybrids.

There is a way around this problem. The five 4d orbitals on nickel are

empty, so we can form a set of empty sp3d

2hybrid orbitals by mixing

the 4d x

2

-y

2, 4d

z

2, 4s, 4 p

x, 4 p

yand 4 p

zorbitals. These hybrid orbitals

then accept pairs of nonbonding electrons from six ammoniamolecules to form a complex ion.

The valence-bond theory therefore formally distinguishes between"inner-shell" complexes, which use 3d , 4s and 4 p orbitals to form a set

of d 2sp

3hybrids, and "outer-shell" complexes, which use 4s, 4 p and 4d

orbitals to form sp3d

2hybrid orbitals.

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Multiple Bonding

Hybrid orbitals and multiple bonds Double Bonds When working geometry problems by the VSEPR method, a double or

triple bond acts just like a single bond for geometry purposes: theadded electrons don't have any effect on the molecular geometry. Thisis because the extra electrons in multiple bonds do not reside in hybridorbitals For example, the ethylene molecule has the Lewis diagram below.

The bond angles around carbon in ethylene are 120o, which implies an

AX3

geometry and sp2

hybridization. One of the two shared electron

pairs on a carbon in the molecule (blue) is in an sp2

hybrid orbital asare the two electron pairs that form bonds to the hydrogen, the other (red) is not. The bonding pair that is in the hybrid orbital is said to

reside in a σ bond. So what type of orbital are the other two electrons in? It turns out thatthey are in the remaining unhybridized p orbitals on the carbon. (The

sp2

hybrid uses one s and two p orbitals, leaving a p orbital on each

carbon.) These two p orbitals overlap, forming a π bond. This orbitalsticks up out of the plane of the molecule, as in the picture below

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Triple bonds In the case of acetylene, HCCH, we have the Lewis structure below

Here, we have sp hybridization and an AX2

geometry around the

carbons. There are two p orbitals left over on each carbon: bothoverlap and thus form two π bonds (red) along with the (blue) σ bondthat results from the electrons in sp hybrid orbitals. Some pictures of

the two sets of π bonds are below

Delocalized bonding ORBITAL PICTURE OF THE CYCLIC OVERLAP IN BENZENE. Notethat all six carbons in benzene are trigonally hybridized and the entiremolecule is coplanar. Each carbon therefore contributes a 2p

zAO to the pi

bonding system. A key aspect of this is that each of these orbitals overlapswith two other 2p

zAO's on adjacent carbon atoms, one to the left and one to

the right, not just to a single one. Thus the system of overlapping orbitals acts

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as a unit, all electrons being entirely delocalized over the ring. Becauseoverlap is much more extensive, bonding is stronger than in a simple alkene pi

bond.

Whenever more than one reasonable valence structure can be writtenfor a species. ( In this case, this normally means that the electrons are moredelocalized than can be shown by any one structure).

Network Covalent Substances Network Covalent Network covalent solids are one common type of solids. In a network covalentsolid, the atoms are

bound to each other by a network of covalent bonds. Common properties of network covalentsubstances are:

High melting point and boiling point. Covalent bonds are very strong.Insoluble in most solvents. Breaking the covalent bonds to dissolve the

substance takes a lotof energy.Poor electrical conductivity. The electrons are localized in covalent bonds,

not spread outlike in a metal.

Examples of network covalent substances are diamond, quartz and graphite.Models of diamondand graphite are shown below if you have Chime installed. Graphite is an

interesting case: it's made

up of sheets of sp2

hybridized carbon atoms in a network covalentarrangement, but the sheetsthemselves aren't bonded to other sheets and are held together only withLondon dispersion forces.Intermolecular Forces Molecules exist as distinct, separate collections of matter. The bonds within a

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molecule are typicallyquite strong, such that it's usually necessary to heat a molecule to very highenergies before the

bonds begin to break. For example, water is stable to decomposition tohydrogen and oxygen up

to temperatures well above 500oC. In contrast, the forces between moleculestend to be relativelyweak. If we chose an arbitrary scale in which the bonds between atoms withina molecule are set at100, then the forces between molecules range between 0.001 and 15. In other words it generallytakes far less energy to separate molecules from one another than it does totake molecules apart.

The forces between molecules are called intermolecular forcesIntermolecular forces are the attractions that exist among particles or molecules of matter inany of the three distinct phases of matter, gases, liquids and solids, found onearth. The strength of these forces increases as the matter change in phase from gas to a liquid andthen to a solid. Thusgases have the weakest intermolecular forces compare to liquids and solidsand solids have the

strongest forces. The intermolecular forces in ionic solids are so strong thatthey exist only as solidseven at very higher temperatures. ionic solids have ionic interactions, thestrong electrostaticattractions between oppositely charged ions to form a ionic lattice as insodium chloride, NaCl(s).Other forms of weaker intermolecular forces resulting from polar covalent

bonds in molecules andthe polarization of electrons on non-polar covalent molecules. These weaker intermolecular forcesare classified according to their strength as shown below:Intramolecular forces: Intramolecular forces are the attractions that existamong particles of amolecules. These are found in gaint molecules such as proteins.

Dipole-dipole (or simply polar) forces

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Dipole-Dipole Intermolecular forces: This include the attraction between all polar molecules through the dipoles (except F-H, O-H and N-H dipoles)dipolar covalent bond formed by unequal sharing of electrons of bonds in amolecule. A molecule can be non-polar even though it may have polar covalent bond because of the symmetry of the molecular structure canceling

the dipoles. Therefore, there are no dipole-dipole interactions in non-polar molecules.Hydrogen bonding Hydrogen bonding are three special cases of stronger dipole-dipoleinteractions resulting from F-H, O-H and N-H dipolar covalent bond arecalled hydrogen bonding. This is because F, O and N have the largestelectronegativities among other elements. Hydrogen bonding (between O-Hand O-H dipoles) in water allows water to exist as a liquid at roomtemperature. N-H dipole in proteins allows the formation of doubles helixstructure in DNA and other complex structures found in living cells.London dispersion forces London Dispersion Forces are the weakest of all intermolecular attractionsand occur in non-polar molecules without dipoles or dipolar covalent bonds.The London dispersion forces results from instantaneous shifts of electroncloud of non-polar molecules. These shifts in electron cloud createinstantaneous dipoles with very short lifetime. A weaker attractive forceresults because of the short lived dipole attractions between two molecules.

Non-polar molecules such as H2 and N2 can be cooled to liquids at very lowtemperature due to the existence of London Dispersion forces. As the number of electrons increases or the molecular weight of a substance increases theLondon Dispersion forces tend to increase making these substances to existsas solids. Even though London dispersion forces exist in ionic solids and polar covalent compounds, their effect is masked by stronger ionic interactions anddipole-dipole interactions.To fully understand intermolecular forces you may wish to review discussionson molecular shape

and VSEPR Theory.

Electronegativity: Deciding Factor of PureCovalent, Polar Covalent and Ionic Bondingbetween two atoms

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In LiH molecule, it would seem that the bonding orbital places moreelectron density on the hydrogen than on the lithium since the orbitalshape describes the probability of finding the electrons. As a result, thehydrogen end of the moelcule would be slightly negative and thelithium end would be slightly positive.This situation is called a polar

bond in which the electrons in the bond are being shared, but notequally shared.

In almost every case in which a bond is formed between two differentatoms the resulting bond will be polar.

In the 1930's, Linus Pauling (1901 -1994), an American chemist who wonthe 1954 Nobel Prize, recognized that

bond polarity resulted from the relativeability of atoms to attract electrons.Pauling devised a measure of thiselectron attracting power which hecalled "electronegativity " which hedefined as the "power of an atom in amolecule to attract electrons to itself."Electronegativity only has meaning ina bond.

The table below presents theelectronegativities for the main groupelements.

Electronegativity

H =2.1

x x x x x x

Li =1.0

Be=1.5

B =2.0

C =2.5

N =3.0

O =3.5

F =4.0

Na=0.9

Mg=1.2

Al =1.5

Si =1.8

P =2.1

S =2.5

Cl =3.0

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Generally, the electronegativity increases moving left to right across arow, and decreases going down the table. Notice that this trend is

violated by the Group 13 metals for which the electronegativity dropsfrom B to Al as expected, but hen rises slightly going down to Tl. Thiseffect is due to the intervention of the d electrons and other effects thatcome into play with very large atoms.

The transition metals are not presented in this chart to conserve room,but their values range from 1.0 to about 2.4.

When we are looking at a chemical bond to predict whether it is polar

(and polar in which direction) we must compare the electronegativitiesof the atoms in the bond. The atom with the higher EN will have thenegative end of the polar bond, and the difference in EN will give arough measure of the extent of the polarity. n the molecules H

2, F

2and O

2which are called homonuclear (same

atoms) diatomic molecules there is no bond polarity. Electrons in the bond are pulled equally andfound in the center between

the two atoms. How you decide whether a bond is pure covalent, polar covalent, or ionic?

0-0.6 0.6 - 1.5 > 1.5 covalent Polar-covalent ionic

bond bond bond

K =0.8

Ca=1.0

Ga=1.6

Ge=1.8

As=2.0

Se=2.4

Br =2.8

Rb

=0.8

Sr =

1.0

In =

1.7

Sn

=1.8

Sb

=1.9

Te

=2.1

I =

2.5

Cs=0.7

Ba=0.9

Tl =1.8

Pb=1.9

Bi =1.9

Po=2.0

At =2.2

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Polarity of a molecular substance is measured as dipole moment. If molecule has a dipole moment it means may have polar covalentbonds If polar covalent bonds are symmetrical, they may lead to zero dipolemoment

Consider the following examples:

Pure Covalent, Polar Covalent, and Ionic Bonds It's not unusual for covalent bonds and ionic bonds to be presented tointroductory students as two clearly distinct categories. In the usualdefinition, covalent bonds involve electron sharing and ionic bondsinvolve electron transfer to form ions. These definitions are correct, butthey miss the reality that pure covalent and ionic bonds sit on a

continuum dictated by the difference in electronegativity. The bigger this difference, the more polar the bond. At some point we wouldexpect the bond to have become so polar that it constitutes a transfer of an electron from one atom to the other. The chart below gives anidea of how this trend works.

Bond Polarities

Bond Diff in EN Negative atom Type of Bond

H - H 0.0 N/A pure covalent

C - H 0.4 C (weakly) polar covalent

O - H 1.4 O polar covalent

H - F 1.9 F polar covalent

S - O 1.0 O polar covalent

C - O 1.0 O polar covalent

Al - C 1.0 C polar covalent

Na - Cl 2.1 Cl ionic

Mg - O 2.3 O ionic

Mg - C 1.3 C polar covalent

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Using some simple calculations of % ionic and % covalent character, itis possible to show that a difference of electronegativity of about 1.7 isthe point where a bond may be thought of as ionic. There are a fewexceptions to this rule such as HF, in which the electronegativitydifference is 1.9, but the molecular properties are decidedly covalent. Again, in many introductory courses, students are presented with athumb rule:

It's easy to see that this rule is only a rough approximation and workswell for compounds of the Group 1 or 2 metals with the halides of Group 17, but doesn't reflect reality when we consider bonds betweencarbon and most metals. Carbon is certainly a non-metal, but it formscovalent bonds (sometimes highly polar, but covalent none the less)

with almost all of the elements on the periodic table. Covalent Bonding and the Periodic Table

In general chemistry you learned that having eight electrons (an octet) inthe valence shell of thenoble gases (group 8A) gives those elements special stability. We also haveseen that thereactions of most elements are driven by their desire to attain a noble gas

Covalent bond = non-metal + non-metalIonic bond = metal + non-metal

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electronic configuration.The alkali metals in group 1A have a single electron in their valence shell andusually give up thiselectron in reactions to achieve a noble gas electronic configuration. Unlikethe alkali metals which

tend to give up their valence electron to achieve a noble gas configuration, thehalogens in group 7A(with 7 valence electrons) tend to gain an electron in reactions to attain a fulloctet of electrons.Elements on the right side of the periodic table (nonmetals) react by gainingelectrons and elementson the left side of the periodic table (metals) react by losing electrons. Thetype of bonding thatoccurs when a nonmetal accepts an electron from a metal is referred to asionic bonding and the

bond is called an ionic bond. This is the type of bonding that occurs in ionicsolids like sodiumchloride, NaCl.Please return to the referring document and note the hypertext link thatled you here.

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chem281- chapter 3: covalent bonding bonding theories

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