C10J Structure Lecture 2

8/3/2019 C10J Structure Lecture 2

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Inorganic Chemistry Structure and Bonding

Learning Objectives

1. Describe how the elements form bonds with each other by

sharing electrons

2. Derive Lewis structures for molecules or polyatomic ions.

3. Predict when resonance structures are possible and draw the

various structures

Reference:

Inorganic Chemistry , 3rd Edition, Chapter 2, pg. 30-33.

Basic Inorganic Chemistry, 3rd Edition, Chapter 3, pg. 73-79


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Bonding

Chemical bonds may be classified into two maincategories: ionic and covalent bonds.

Ionic bond - purely electrostatic attraction betweenoppositely charged particles (ions). Electronegativitydifference, (EN, between elements is relatively large.

Covalent bonding - used to describe the bonds in

compounds that result from the sharing of one ormore pairs of electrons. (EN, between elements iszero or relatively small.

-two center (localized) vs. multicenter

(delocalized) covalent bonding


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The Localized Bond Approach

3 relevant theories work simultaneously toexplain bonding and geometry of molecules andcomplex ions:

- Lewis electron-pair bond theory

- Hybridization theory

- Valence Shell Electron PairRepulsion(VSEPR) theory


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Covalent Bonding Lewis Concepts

One covalent bond between two atoms results

from the sharing of a pair of electrons between

the two atoms.

This pair of electrons is called the bonding pair and

is represented by a line connecting the two atoms


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Covalent Bonding Lewis Concepts

Electrons that are not shared are not involved in the bond and are

localized as lone-pairs (non-bonding electrons) on atoms within

the molecule.

One bonding pair one bond, three bonding pairs three bonds(triple bond)

These diagrams are called Lewis Diagrams

Using the concept that the electronic structure of the entire molecule

is represented by the sum of all the bonding and lone pairs of

electrons, Lewis diagrams can be drawn for a molecule or ion to

represent the approximate arrangement of atoms and the valenceelectrons within the structure.


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Deriving Lewis structures:

Here are nine points to take into account in deriving Lewis structures:

(i) only valence electrons are used

(ii) a simple (symmetrical) geometry is used to represent the shape

of the molecule, or polyatomic ion

(iii) metals must be the central atom, and H is peripheral. The least

electronegative element is usually the central element

(iv) the bond is represented by a dash or line

(v) each non hydrogen atom is given an octet (8 electrons). This is

the OCTET RULE. There are exceptions to the Lewis octet rule.

(vi) if any atom has less than an octet, multiple bonds may be used

so that all atoms have an octet(vii) electrons not shared between atoms remain as lone pairs of

electrons on the parent atom

(viii) for a cation, subtract 1 electron for each +ve charge, and in the

case of an anion, add 1 electron for each ve charge

(ix) the entire structure of the molecule (ion) is represented by the

sum of all bonding and lone pairs of electrons.


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A stepwise approach to writing Lewis

structures

Step 1: Determine the total number of valenceelectrons.

Step 2:W

rite the skeleton structure of the molecule.

Step 3: Use two valence electrons to form eachbond in the skeleton structure.

Step 4: Try to satisfy the octets of the atoms bydistributing the remaining valence electronsas nonbonding electrons (lone pairs).

Step 5: Use FC (formal charges) to decide which of

more than 1 possible structures is correct.


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A stepwise approach to writing Lewis

structures

Example

Derive the Lewis Structure for the chlorate ion

ClO3-.


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Step 1: Determine the total number of

valence electrons.

ClO3-: Cl has 7 valence electrons = 7 e-

Each O atom has 6 valence electrons (3 x 6)= 18 e-

add 1 electron for the negative charge= + 1e-

Total 26 e-


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Step 2: Write the skeleton structure of

the molecule.

a simple (symmetrical) geometry is used to represent

the shape

Cl

O

OO


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Step 3: Use two valence electrons to

form each bond in the skeleton structure.

Two electrons for each bond is 3x2 = 6 e-

This leaves 26 e (total) 6 e (bonding) =

20 nonbonding electrons

Cl

O

OO


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Step 4: Satisfy the octets of the atoms

lone pairs

Cl

O

OO

..

..

....

.. ..

..

....

..

-


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Molecules that Contain Too Few

Electrons Multiple Bonds Sometimes the molecule does not have enough

valence electrons

Cant get satisfactory Lewis structure by sharinga single pair of electrons, then share two or eventhree pairs of electrons.

Result double or triple bonds

Elements that form strong double or triple bondsare C, N, O, P, and S


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Electrons Multiple Bonds

Example

Derive the Lewis Structure for the nitrate ion

NO3-.


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Electrons Multiple Bonds

NO3-:N has 5 valence electrons = 5 e-

Each O atom has 6 valence electrons (3 x 6) = 18 e-

add 1 electron for the negative charge = 1e-

Total 24 e-

If we use only single bonds then the N atom would only have 6 valence

electrons ( 2 from each bond)

Lewis Structure

N

O

O O

-


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Electrons Electron Deficient

Consider BF3.

Note neither B nor F tends to form multiple bonds so Bwill have only six valence electrons and the Lewis

structure is written with less than an octet on B.

Molecules that feature electron deficiency usually

involve B, Be and sometimes Al.

What is the difference b/w an unsaturated system and an

electron deficient one?


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Molecules that Contain Too Many

Electrons Valence Shell Expansion Consider SF4:

6 + 4(7) = 34 valence electrons

4 covalent bonds requires 8 e so S has its octet

26 nonbonding valence electrons remaining (34e 8e = 26e)

Each fluorine atom needs 6e to satisfy its octet total 24 e


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Molecules that Contain Too Many

Electrons Valence Shell Expansion But there are 26 nonbonding electrons in this molecule.

We have an additional pair of valence electrons (26e 24 e = 2e)

So we expand the valence shell of the sulfur atom to hold more than

eight electrons.

Lewis Structure

How does sulfur hold 10 e

in its valence shell?


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Other examples where octet rule does

not apply

Coordinate covalent bonds in coordination

compounds:

-ligands as Lewis bases (electron-pair donors)

-metals as Lewis acids (electron-pair acceptor)

more in C10K!!!


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Example of the use of FC in writing

correct Lewis Structures

Class Exercise:

Write the most favorable Lewis Formula,

(Draw the most favorable Lewis Structure) for

nitrosyl chloride, NOCl. (3 marks)

Problems you might experience with this question:

Two possible structures. Hint: which is preferable?

Where do I put the double bond?

Advice: Use formal charges (FC) to decide


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More than 1 possible Lewis Formula??

Use Formal Charges (FC) to get correct formula

FC is the hypothetical charge on an atom in a molecule or polyatomicion. Formal charges must add up to the actual charge on the species,that is 0 for molecules and the ionic charge for ions.

The most energetically favorable Lewis Formula is the one inwhich FC on each atom is 0 or as near to 0 as possible.

-structures placing negative formal charges on more

electronegative elements are usually favored.

TO CALCULATE FC

FC = (group number)-[(number of bonds) + (number of unsharedelectrons)]

Complete Class Exercise:

Calculate the FC involved and then write the correct Lewis Formula fornitrosyl chloride, NOCl.

Go to the board !


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More than 1 correct Lewis Formula??

Consider Resonance

Resonance structures have the same relativeplacement of atoms but different locations ofbonding and lone electron pairs. Below are the 3

resonance forms ofCO32-

.

C

O

O O

2-

C

O

O O

2-

C

O

O O

2-

RESONANCE STRUCTURES ARE NOT REAL DEPICTIONS!

For example, the diagram above implies that 2 single bonds and 1 double bond

are present in the carbonate ion and the double arrow implies that the structure

switches from one resonance form to the next. NOTTRUE!


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Resonance - Example CO32- .

IN FACT experimentally, it is observed that all C-O bonds areequal, and all the O-C-O bond angles are equal.

The real molecule has an electron distribution corresponding tothe average of 3 contributing structures, and is said to be a

resonance hybrid of the three structures. Each individual resonance form has no real existence, but their

average gives the actual structure:

C

O

O O

2-

resonance hybrid

Note this is different from the resonance forms of NOCl discussed in class. NOCl

does not exist as a resonance hybrid since the atoms surrounding the central atom

are different. Hence its structure is dominated by one of the resonance forms which

is determined by the formal charges.


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Have we achieved our objectives?

CAN YOU:

1. Describe how the elements form bonds with each other by sharing

electrons

2. Derive Lewis structures for molecules or polyatomic ions.

3.Predict when resonance structures are possible and draw the

various structures?????????

Attempt the next tutorial.

Remember Lewis structures does not necessarily show shapesso Next Class

Hybridization and Valence Shell Electron PairRepulsion(VSEPR) Theory

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C10J Structure Lecture 2