2013 Lect2b Chemical Properties- Electronic Structure and Chemical Bonding

7/27/2019 2013 Lect2b Chemical Properties- Electronic Structure and Chemical Bonding

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Structure BASICS

of Organic Molecules


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Objectives

Objectives

Know how to use the periodic table

Understand atomic structure of an atom including its massnumber, isotopes, and orbitals

Know how atomic orbitals overlap to form molecular orbitals Understand orbital hybridization

Using the VSEPR model, predict the geometry of molecules

Understand the formation of molecular orbitals

Know how to draw Lewis structures

Predict the direction and approximate strength of a bond dipole

Using a Lewis structure, find any atom or atoms in a moleculethat has a formal charge

Understand how to draw resonance structures


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ELECTRON CONFGURATION

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The basic unit of matter is an atom. It consists dense centralnucleus surrounded by negatively charged electrons.

The nucleus contains a mix of positively charge proton and electrically

neutral neutrons.

The electrons of an atom are bounded to the nucleus by

electromagnetic force

The electrons determine the chemical properties of an

element .

ATOM AND ITS ELECTRONS


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It means the arrangement of electrons of an atom.

The knowledge of electron configuration of different

atoms is useful in understanding the structure of

elements in periodic table.

The concept also useful for describing chemical bonds

that hold atoms together

ELECTRON CONFIGURATION


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Chemical BondFormation


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A chemical bondis an attraction between atoms that allowsthe formation of chemical substances that contain two or more atoms.

The bond is caused by the electrostatic force of attractionbetween opposite charges, either between electrons and

nuclei, or as the result of a dipole attraction.


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the negatively charged electrons that are orbiting the nucleus and the

positively charged protons in the nucleus attract each other.

Also, an electron positioned between two nuclei will

be attracted to both of them

These electrons cause the nuclei to be attracted toeach other, and this attraction results in the bond

However, this assembly cannot collapse to a size dictated by the

volumes of these individual particles. Due to the matter wave natureof electrons and their smaller mass, they occupy a much larger

amount of volume compared with the nuclei, and this volume

occupied by the electrons keeps the atomic nuclei relatively far apart,

as compared with the size of the nuclei themselves.


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In general there are two types of chemical bonds:

Ionic Bonds - Bonds formed when one or more

electrons are transferred from one atom

to another atom.

This attraction may be seen as the result of different

behaviors of the outermost electrons (VE) of atoms.

The strength of chemical bonds varies considerably;

there are "strong bonds" such as covalent or ionic bonds

and "weak bonds" such as dipoledipole interactions, theLondon dispersion force and hydrogen bonding.


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In this type of bond, the outer atomic orbital of one atom has a vacancy

which allows addition of one or more electrons.

These newly added electrons potentially occupy a lower energy-state than

they experience in a different atom. Thus, one nucleus offers a more tightly

bound position to an electron than does another nucleus, with the result that

one atom may transfer an electron to the other.

This transfer causes one atom to assume a net positive charge, and theother to assume a net negative charge. The bondthen results from

electrostatic attraction between atoms, and the atoms become positive or

negatively charged ions.

Often, such bonds have no particular orientation in space, since they result

from equal electrostatic attraction of each ion to all ions around them.

Ionic bonds are strong since the forces between ions are short-range, and do

not easily bridge cracks and fractures. This type of bond gives a charactistic

physical character to crystals of classic mineral salts, such as table salt.

a simplified view of an ionicbond,


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Ionic Bonding

Na F

Sodium Atom Fluorine Atom


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Ionic Bonding (2)

Na F

Attraction between the two ions is electrostatic --

I onic Bond

Sodium ion Fluoride ion


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Covalent Bonds -

Bonds formed when two

atoms share one or more

pairs of valence electrons.


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a pair of electrons are drawn into the space between the two atomic

nuclei. Here the negatively charged electrons are attracted to the

positive charges ofboth nuclei, instead of just their own. This

overcomes the repulsion between the two positively charged nuclei of

the two atoms, and so this ing attraction holds the two nuclei in a fixed

configuration of equilibrium,

Thus, covalent bonding involves sharing of electrons in which thepositively charged nuclei of two or more atoms simultaneously attract

the negatively charged electrons that are being shared between them.

These bonds exist between two particular identifiable atoms, and have

a direction in space, allowing them to be shown as single connectinglines between atoms in drawings, or modeled as sticks between

spheres in models.

In a polar covalent bond, one or more electrons are unequally shared

between two nuclei.

the simplest view of a so-called 'covalent' bond,


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Covalent bonds often result in the formation of small collections of better-

connected atoms called molecules, which in solids and liquids are bound toother molecules by forces that are often much weaker than the covalent

bonds that hold the molecules internally together. Such weak intermolecular

bonds give organic molecular substances, such as waxes and oils, their

soft bulk character, and their low melting points (in liquids, molecules must

cease most structured or oriented contact with each other).

When covalent bonds link long chains of atoms in large molecules, however

(as in polymers such as nylon), or when covalent bonds extend in networks

though solids that are not composed of discrete molecules (such as

diamond or quartz or the silicate minerals in many types of rock) then the

structures that result may be both strong and tough, at least in the direction

oriented correctly with networks of covalent bonds. Also, the melting points

of such covalent polymers and networks increase greatly.


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A SIMPLE COVALENT BOND

H . H.

A pair of electrons is shared between the two bonded atoms.


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When two atoms move into close proximity, they experience a change in energy. At the

distance ofthe bond length, they achieve minimum energy.

energy is released during the formationof the bond.

DISTANCE AND ENERGY RELATIONSHIP


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BOND DISSOCIATION ENERGIES.

Conversely, breaking bond of two atoms in a molecule requires an input

of energy because the energy level of the molecule is lower than the

energy level of the two atoms.


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The electrons of the atoms possess a set of stable energy levels

or ORBITALS, and can undergo transitions between orbital by

absorbing or emitting photons that match the energy differences

between the level.

ELECTRON CONFIGURATION


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Electrons of atom occupy a set of allowed states which called SHELL

the first shell can accommodate 2 electrons, the second shell

8 electrons, and the third shell 18 electrons

(An atom's nth electron shell can accommodate 2n2 electrons)

A SUBSHELL is the set of states defined by azimuthalquantum number, l = 0, 1, 2, 3 which correspond to the s,p, d,

and flabels. The max number of electrons which can be placed

in a subshell is given by 2(2l+ 1).

This gives :2 electrons in a s subshell,

6 electrons in a p subshell,

10 electrons in a d subshell

14 electrons in a f subshell.

SHELL AND SUBSHELL of ATOM


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Electron configuration

The outermost electron shell is often referred to as

the "VALENCE SHELL" and determines the chemical

properties


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LEWIS STRUCTURE

OF MOLECULES

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Know how to draw Lewis structures

Using a Lewis structure, find any atom

or atoms in a molecule that has a formal

charge

Understand how to draw resonance

structures


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Lewis structure

diagrams that show the bonding between atoms of a

molecule and the lone pairs of electrons that may exist

in the molecule.

A Lewis structure can be drawn for any covalently

bonded molecule, as well as coordination

compounds.


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C. Lewis Symbols for VE

The elements symbol represents the nucleus and the core

electrons.

Dots are placed around the symbol to represent the VE.

The electrons are paired if there are more than four VE.

A maximum of 4 pairs of electrons (8 e) areaccommodated by a main group element.


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Lewis Structures

show the arrangement of atoms and VE in a molecule.

A pair of electrons that are shared in a covalent bond is

represented by a line between a pair of atoms.

Electrons that are shared by two atoms are called bonding

pairs. Electrons that reside on a single atom and are not shared

are called lone pairsornonbonding electrons.

Some elements can share more than 1 pair of e in a

covalent bond called double bondsand triple bonds.

Most atoms in a molecule are surrounded by 8 electrons to

achieve noble gas configuration. The tendency of these

atoms to be surrounded by 8e is known as the octet rule.


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Construction of the Lewis Dot diagram

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The total number of electrons represented in a Lewis

structure is equal to the sum of the numbers of valence

electrons on each individual atom.

Non-valence electrons are not represented in Lewis

structures.

The octet rule states atoms with eight electrons in theirvalence shell will be stable, regardless of whether these

electrons are bonding or nonbonding.


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Lewis structures for oxygen, fluorine, thehydrogen sulfate anion, and formamide


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Drawing Lewis Structures

1. Decide on the central atom.

2. Determine the total number of valence electrons in the

molecule or ion.

3. Place one pair of electrons between each pair ofbonded atoms to form a single bond.

4. Use the remaining electrons as lone pairs around each

terminal atom (except H) and then the central atom so

that each is surrounded by eight electrons.

5. If the central atom has fewer than eight electrons, move

one or more lone pairs from terminal atoms to form

multiple bonds to the central atom.


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Predicting Lewis Structures

1. Hydrogen Compounds


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Example: Lewis structure of

the nitrite ion

The formula of the nitrite ion is NO2.

Step one: Nitrogen is the least electronegative atom, so it is the

central atom by multiple criteria.

Step two: Count valence electrons. Nitrogen has 5 valence electrons;

each oxygen has 6, for a total of (6 2) + 5 = 17. The ion has a charge of

1, which indicates an extra electron, so the total number of electrons is 18.

Step three: Place ion pairs. Each oxygen must be bonded to the nitrogen,

which uses four electrons two in each bond. The 14 remaining electrons

should initially be placed as 7 lone pairs. Each oxygen may take amaximum of 3 lone pairs, giving each oxygen 8 electrons including the

bonding pair. The seventh lone pair must be placed on the nitrogen atom.


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Step four: Satisfy the octet rule. Both oxygen atoms currently have

8 electrons assigned to them. The nitrogen atom has only 6 electrons

assigned to it. One of the lone pairs on an oxygen atom must form a

double bond, but either atom will work equally well. We therefore must

have a resonance structure.

Step five: Tie up loose ends. Two Lewis structures must be drawn:

one with each oxygen atom double-bonded to the nitrogen atom. The

second oxygen atom in each structure will be single-bonded to thenitrogen atom. Place brackets around each structure, and add the

charge () to the upper right outside the brackets. Draw a double-

headed arrow between the two resonance forms.
http://en.wikipedia.org/wiki/File:Nitrite-ion-lewis-canonical.pnghttp://en.wikipedia.org/wiki/File:Nitrite-ion-lewis-canonical.pnghttp://en.wikipedia.org/wiki/File:Nitrite-ion-lewis-canonical.pnghttp://en.wikipedia.org/wiki/File:Nitrite-ion-lewis-canonical.pnghttp://en.wikipedia.org/wiki/File:Nitrite-ion-lewis-canonical.png


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2. Oxo Acids and Their Anions


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3. Isoelectronic species have the same number of valence

electrons and the same Lewis structure.


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RESONANCE

Some compounds can have more than one valid Lewisstructure.

Multiple Lewis structures that represent the same

molecule are called resonance structures.

Often the actual structure of a molecule

with resonance structures is a

composite of all the resonance

structures called a resonance hybrid.

RESONANCE HYBRIDE


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RESONANCE HYBRIDEMultiple Lewis structures

1) sulfur dioxide

2) nitric acid

3) formaldehyde

resonance

hybrid

major

contributor

minor

contributor

non

contributor


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4) carbon

monoxide

5) azide anion

RESONANCE HYBRIDE

Multiple Lewis structures


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D. Exceptions to the Octet Rule

1. Compounds in which an atom has fewer than eight

electrons. Group 3 atoms (B, Al, Ga, In, Tl) are

commonly surrounded by six

electron or three electron pairs.

These atoms can formcoordinate covalent bonds in

which the pair of electrons in

the covalent bond originate

from one atom.


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2. Compounds in which an atom has more than eight

electrons.

Element from the third period or higher can formmolecules or ions in which the octet rule is exceeded.


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3. Molecules with an Odd Number of Electrons Compounds containing an odd number of nitrogens

will have an odd number of electrons making it

impossible to draw a structure that obeys the octet

rule.

N

O O

N

OO

Chemical species with an unpaired electron are called

free radicals and are highly reactive.


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CHARGE DISTRIBUTION

IN COVALENT BONDS AND

MOLECULES


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We need to determine these loci of charge concentration,

to figure out if an atom is negative, positive, or neutral?

in order to understand chemical reactivity.

Many important organic species are ions.

In these ions, charges appear to be preferentially

concentrated on certain atoms.

C

O

OO

2-


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Calculation of Formal Charges

Formal

charge

Number ofValence

electrons

Number ofnonbonding

electron

Half number ofelectron in

covalent bond

= - -

Heres the formula for figuring out the formal

charge of an atom:


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Determination of Formal Charges

Determine the numberof valence electrons

which are present in the neutralatom.

Subtract the non-bonding electrons. Theseelectrons belong to the atom.

Subtract one-halfof the bonding electrons.

This formula explicitly spells out how many

electron are formally owned by the atom.


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is the charge assigned to an atom in a

molecule, assuming that electrons in a chemicalbond are shared equally between atoms,

regardless of relative electronegativity.

Formal charge


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Charge Distribution in Covalent Bonds

A. Formal Charge

LPE = lone pair electrons

BE = bonding electrons

BE)](Number21LPE[Number-NumberGroupChargeFormal

NH H

H

H N H

H

H

Formal Charge = 5 - [2 + (6)] = 0

Formal Charge = 1 - [0 + (2)] = 0

Sum of Formal Charges = 0 + 0 + 0 + 0 = 0

Formal Charge = 5 - [0 + (8)] = + 1

Formal Charge = 1 - [0 + (2)] = 0

Sum of Formal Charges = 1 + 0 + 0 + 0 = +1



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A. Formal Charge

LPE = lone pair electrons

BE = bonding electrons

BE)](Number21LPE[Number-NumberGroupChargeFormal

Formal Charge = 6 - [4 + (4)] = 0

C

O

OO

2-

Formal Charge = 6 - [6 + (2)] = - 1

Formal Charge = 4 - [0 + (8)] = 0

Sum of Formal Charges = -1 + -1 + 0 + 0 = -2

Lets apply it to some examples.


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http://masterorganicchemistry.files.wordpress.com/2010/09/formal-charge-copy.jpghttp://masterorganicchemistry.files.wordpress.com/2010/09/formal-charge-copy.jpg


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More on Formal Charges

The sum of the formal charges must be equalto the total charge on the ion or molecule

If there arent enough electrons to provideevery atom with an octet, consider double ortriple bonds

All structures must obey the rules of valence


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Still More on Formal Charges

Separated formal charges should be

avoided, if possible

Where there are separated formalcharges, the negative formal charge

should reside on the more

electronegative element


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2013 Lect2b Chemical Properties- Electronic Structure and Chemical Bonding

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