xi chemistry unit 4

8/7/2019 XI Chemistry Unit 4

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OVERVIEW

Kossel-Lewis Approach to chemical bonding Octet Rule

Covalent Bond

Lewis Structures Simple Molecules

Ionic or Electrovalent Bond

Bond Parameters

VSEPR Theory

Orbital Overlap Hybridisation

Molecular Orbital Theory

Summary


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Most bonds are somewhere in between.

Forms of ChemicalBonds

There are TWO extreme forms ofconnecting or bonding atoms:

Ioniccomplete transfer of electronsfrom one atom to another

Covalentelectrons shared betweenatoms


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Electron Distribution in

Molecules Electron distribution is

depicted with Lewis

electron dotstructures

Electrons aredistributed as:

shared or BONDPAIRS and

unshared or LONEPAIRS.

G. N. Lewis1875 - 1946


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Chemical bond: attractive force holding two or more

atoms together.

Covalent bond results from sharing electrons between the

atoms. Usually found between nonmetals.

Ionic bond results from the transfer of electrons from a

metal to a nonmetal.

Metallic bond: attractive force holding pure metalstogether.

Chemical Bonds, Lewis Symbols,and the Octet Rule


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All noble gases except He has an s2p6 configuration.

Octet rule: atoms tend to gain, lose, or shareelectrons until they are surrounded by 8 valence

electrons (4 electron pairs).

Caution: there are many exceptions to the octet

rule.

The Octet Rule


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Usually occurs with B and elementsof higher periods. Commonexceptions are: Be, B, P, S, andXe.

BF3BF3

SF4SF4

Be: 4

B: 6

P: 8 OR 10

S: 8, 10, OR 12

Xe: 8, 10, OR 12

Voilations of Octet Rule


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Central Atoms Having Less than an Octet

Relatively rare.

Molecules with less than an octet are typical for

compounds of Groups 1A, 2A, and 3A.

Most typical example is BF3.

Formal charges indicate that the Lewis structure with

an incomplete octet is more important than the oneswith double bonds.

Exceptions of Octet Rule


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Central Atoms Having More than an Octet

This is the largest class of exceptions.

Atoms from the 3rd period onwards can accommodate

more than an octet.

Beyond the third period, the d-orbitals are low enough in

energy to participate in bonding and accept the extra

electron density.

Exceptions of Octet Rule


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Covalent Bonding

C l


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CovalentBonding

Covalent bond is the sharing of the VALENCEELECTRONS of each atom in a bond

Recall: Electrons are divided between

core and valence electrons.ATOM core

valenceNa 1s2 2s2 2p6 3s1 [Ne]

3s1[Ar] 3d10 4s2 4p5 [Ar] 3d10 4s2 4p5


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Valence Electrons1A

2A 3A 4A 5A 6A 7A

8A

Number of valence electrons is equalto the Group number.

C t


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Cova entBondingThe bond arises from the mutual attraction of 2 nuclei for the same electrons.

HB+ H

A

HB

HA

A covalent bond is a balance

of attractive and repulsive forces.

+

-

-

+


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BondFormation

A bond can result from a head-to-head overlap of

atomic orbitals on neighboring atoms.

H H Cl

Cl

+

Overlap of H (1s) and Cl (2p)

his type of overlap places bonding electrons in aOLECULAR ORBITAL along the line between

the two atoms and forms a SIGMA BOND (s).


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More sharing examples

O2

N2

O O

N N

O O O O

N N N N N N

double bond (2 pairs)

triple bond (3 pairs)

Share until octet is complete.

octet complete


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Bond Energy

F2

single bond BE = 142 kJ/mole

O2 double bond BE = 494

N2

triple bond BE = 942

X2 + energy X + X

increasing

bon

d

stren

gth

Is breaking a bond an endothermic or exothermic process?


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NH3NH

H

H

NH

H

HH+

NH4+ NH3 + H

+ NH4+

coordinate covalent bond(the pair of electronsfrom the same atom)

normal covalent bond(each atom supplies

an electron)

Some more sharing examples


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Follow Step by Step Method1. Total all valence electrons. [Consider Charge]

2. Write symbols for the atoms and guess skeleton

structure [ define a central atom ].

3. Place a pair of electrons in each bond.

4. Complete octets of surrounding atoms. [ H = 2 only ]

5. Place leftover electrons in pairs on the central atom.

6. If there are not enough electrons to give the central atom

an octet, look for multiple bonds by transferring

electrons until each atom has eight electrons around it.

Drawing Lewis Structure


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Resonance Structures

Drawing LewisStructures


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To determine the electron pair geometry: draw the Lewis structure,

count the total number of electron pairs around the central

atom,

arrange the electron pairs in one of the above geometries to

minimize e-e repulsion, and count multiple bonds as one

bonding pair.

VSEPR Model


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VSEPR Model

d l


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VSEPR Model

d l ( i


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VSEPR Model (Domainsaround Central Atom)


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Drawing Lewis Structures

Formal Charge

Consider:

For C:

There are 4 valence electrons (from periodic table).

In the Lewis structure there are 2 nonbonding electrons and 3

from the triple bond. There are 5 electrons from the Lewis

structure.

Formal charge: 4 - 5 = -1.

C N


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Drawing Lewis Structures

Formal Charge Consider:

For N: There are 5 valence electrons. In the Lewis structure there are 2 nonbonding electrons and 3

from the triple bond. There are 5 electrons from the Lewis

structure.

Formal charge = 5 - 5 = 0. We write:

C N

C N


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Bond and LonePairs

Electrons are distributed as shared or BONDPAIRS and unshared or LONE PAIRS.

H Cl

This is a LEWIS ELECTRON DOT structure.

shared or bond pair

Unshared orlone pair (LP)

Rules of Lewis


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his observation is called the OCTET RULE

Rules of LewisStructures

No. of valence electrons of an atom =Group number

xcept for H

(and atoms of 3rd and higher periods),#Bond Pairs + #Lone Pairs = 4

For Groups 5A-7A (N - F),o. of BOND PAIRS = 8 - group No.

or Groups 1A-4A (Li - C),o. of BOND PAIRS = group number

Building a Dot


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2. Count valence electronsH = 1 and N = 5

Total = (3 x 1) + 5= 8 electrons or

Decide on the central atom; never H.entral atom is atom of lowest affinity for electrons

n ammonia, N is central

Building a DotStructure

Ammonia, NH3

4 pairs


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4. Remaining electrons formLONE PAIRS to completeoctet as needed.

3. Form a sigma bondbetween the central atomand surrounding atoms.

H H

H

N

Building a Dot Structure

H H

H

N

3 BOND PAIRS and 1 LONE PAIR.e that N has a share in 4 pairs (8 electrons), whileh H shares 1 pair.


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Step 2. Count valence electronsS = 63 x O = 3 x 6 = 18

Negative charge = 2TOTAL = 6 + 18 + 2 = 26 e-

or 13 pairs

Step 1. Central atom = S

10 pairs of electrons are left.

Sulfite ion, SO32-

Step 3. Form sigma bonds

O O

O

S

Sulfite ion SO 2-


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Remaining pairs become lone pairs,first on outside atomsthen on central atom.

Sulfite ion, SO32

(2)

Each atom is surrounded by an octet ofelectrons.

O O

O

S

TE - must add formal charges (O-, S+) for complet diagram


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Carbon Dioxide, CO2

1. Central atom = __C____2. Valence electrons = _16_ or _8_

pairs

3. Form sigma bonds.O OC

O OC

This leaves __6__ pairs.4. Place lone pairs on outer atoms.


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O OC

Carbon Dioxide, CO2 (2)

4. Place lone pairs on outer atoms.

O OC

O OC

The second bonding pair forms a pi (p) bond.

. To give C an octet, form DOUBLE BONDSetween C and O.


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SO3

H2CO

Double and eventriple bonds arecommonly observedfor C, N, P, O, and S

O OC

C2F4


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Sulfur Dioxide, SO2

1. Central atom = S2. Valence electrons = 6 + 2*6 = 18 electrons

or 9 pairs

O OS

O OS

bring in

left pair

OR bring in

right pair

3. Form pi ( ) bond so that S has an octet note that there are two ways of doing

this.

Sulfur


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SulfurDioxide, SO2

O OS

bring inleft pair

OR bring in

right pair

O OS

O OS

Equivalent structures

called:RESONANCESTRUCTURES

The proper Lewis struct

is a HYBRID of the two.

ETTER representation of SO2

made by forming 2 double bonds

O = S = O

Each atom has- OCTET- formal charge = 0

Violations of the


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Violations of theOctet Rule

Usually occurs with:

Boron

BF3SF

4

Elements of higher periods.

Boron


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BoronTrifluoride

Central atom = B

Valence electrons = 3 + 3*7 = 24

or electron pairs = 12

Assemble dot structure

F

F

F

B

The B atom has a sharein only 6 electrons (or

3 pairs). B atom inmany molecules iselectron deficient.


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Central atom = S

Valence electrons = 6 + 4 * 7 = 34 e -

or 17 pairs.

Form sigma bonds and distributeelectron pairs.

F

F

F

S

F

5 pairs around the S atom. A5 pairs around the S atom. A

common occurrence outsidecommon occurrence outsidethe 2nd period.the 2nd period.

Sulfur Tetrafluoride, SF4

Formal Atom


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Formal charge = Group no.

- 1/2 (no. bond electrons)

- (no. of LP electrons)

Formal AtomCharges

he most important dominant resonance structuref a molecule is the one with formal chargess close to 0 as possible.

Carbon


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04 - (1/ 2)(8) - 0 =

6 - (1/2)(4) - 4 = 0

CarbonDioxide, CO2

At OXYGEN

O C O

At CARBON

Carbon Dioxide


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C atom chargeis 0

6 - (1/2)(6) - 2 = +1

6 - (1/2)(2) - 6 = -1

Carbon Dioxide,CO2 (2)

O C O

n alternate Lewis structure is:

AND the correspondingresonance form

+

O C O

+


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Boron Trifluoride, BF3

F

F

F

B

What if we form a BF doublebond to satisfy the B atomoctet?

Boron Trifluoride


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Boron Trifluoride,BF3 (2)

To have +1 charge on F, with its very high electronaffinity is not good. -ve charges best placed onatoms with high EA.

Similarly -1 charge on B is bad

NOT important Lewis structure

fc = 7 - 2 - 4 = +1 Fluorine

fc = 3 - 4 - 0 = -1 BoronF

F

F

B

+

St t D t i ti b


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Water, H2O

H O H

H O H

2 bond

pairs2 lone

pairs

Structure Determination byVSEPR

The electronpair

geometry isTETRAHEDRAL

Themoleculargeometry isBENT.


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Ammonia, NH3

The electron pair geometry is

TETRAHEDRAL.The Molecular geometry the positionsof the atoms is TRIGONAL PYRAMID.

Structure Determination byVSEPR

H

H

H

lone pair of electronsin tetrahedral position

N


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TAKE A BREAK


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Ionic Bonding


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Ionic Bonds

Ionic compounds

Essentially complete electron transfer froman element of low IE (metal) to an elementof high electron affinity (EA) (nonmetal)

Na(s) + 1/2 Cl2(g) Na+ + Cl-

NaCl (s)

Ionic bonding


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ClNa+ Cl

Ionic bonding Ionic bonding involves 3 steps (3 energies)

1) loss of an electron(s) by one element, 2) gain of

electron(s) by a second element, 3) attraction betweenpositive and negative

Na Cl

e1) 2)

3)

Na+


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NON-DIRECTIONALbonding

via Coulomb (charge)interaction

rimarily between metalsrps 1A, 2A and transition metals)

d nonmetals (esp O and halogens)

Ionic Bonds


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Bond Type Single Double Triple

# of es 2 4 6

Notation =

Bondorder

1 2 3

Bondstrength

Increases from Single to Triple

Bondlength

Decreases from Single to Triple

Chemical Bonds

Average Bond Enthalpies


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Average Bond Enthalpies(KJ/mol)

Average Bond Lengths of some Single


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Average Bond Lengths of some Single,Double and Triple Bonds

L i S b l /


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Lewis Symbols/Electronic Configuration


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Electronegativity: The ability of one atoms in a

molecule to attract electrons to itself.

Pauling set electronegativities on a scale from0.7 (Cs) to 4.0 (F).

Electronegativity increases

across a period and

down a group.

Electronegativity

Electronegativity of


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g yElements

Electronegativity and


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There is no sharp distinction between bonding types.

The positive end (or pole) in a polar bond is represented

+ and the negative pole -.

Electronegativity andBond Polarity


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HCl is POLAR because ithas a positive end and anegative end. (difference

in electronegativity)

Cl has a greater share inbonding electrons than

does H.

Cl has a greater share inbonding electrons than

does H.

Cl has slight negative charge (-d) and H has

slight positive charge (+ d)

H Cl

+ -

H Cl

+ -

Bond Polarity


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why oil and water will not mix! Oil isnonpolar, and water is polar.

The two will repel each other, and so youcan not dissolve one in the other

Bond Polarity


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Valence bond theory

Valence bond theory utilizes orbitals andelectrons coming together to form thecovalent bonds in a molecule.

According to valence bond theory, a bondbetween two atoms is formed when twoelectrons with their spins paired are

shared by overlapping atomic orbitals

Molecular Shapes


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There are five fundamental geometries for molecular

shape:

Molecular Shapes(VSEPR)


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Summary of VSEPR shapes


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e-pairs Notation Name of VSEPRshape

Examples

2 AX2 Linear HgCl2 , ZnI2 , CS2 , CO2

3 AX3 Trigonal planar BF3 , GaI3

AX2E Non-linear (Bent) SO2 , SnCl2

4 AX4 Tetrahedral CCl4 , CH4 , BF4-

AX3E (Trigonal) Pyramidal NH3 , OH3-

AX2E2 Non-Linear (Bent) H2O , SeCl2

5 AX5 Trigonal bipyramidal PCl5 , PF5

AX4E Distorted tetrahedral(see-sawed)

TeCl4 , SF4

AX3E2 T-Shaped ClF3 , BrF3

AX2E3 Linear I3- , ICl2

-

6 AX6 Octahedral SF6 , PF6-

y p


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Sigma Bond Formation by


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Sigma Bond Formation byOrbital Overlap

sigma bond ( )

+HH

Two s AtomicOrbitals (A.O.s)

overlap to form ans (sigma)Molecular Orbital

(M.O.)

Sigma Bond Formation by


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Orbital Overlap

sigma bond ( )

+HH

Two s A.O.s overlap tofrom an s M.O.

Similarly, two p A.O.s

can overlap end-on tofrom a p M.O.

eg F2


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Hybridization

In 1931, Linus Pauling proposed thewave functions

for the s and p atomic orbitals .

The mathematical process ofreplacing pure atomic

orbitals with reformulated atomicorbitals for bonded

atoms is called hybridization. In a hybridization scheme, the

number of hybrid

orbitals equals to the total number of


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H b idi ti


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3

signifies one s and three p orbitals are combined

Mixing one s orbital with three p orbitals yieldsfour equivalent sp3 hybrid orbitals.

Hybridization

Hybridization


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The formation of four sp3 hybrid orbitals by combinationof an atomic s orbital with three atomic p orbitals. Eachsp3 hybrid orbital has two lobes, one of which is largerthan the other. The four large lobes are oriented toward

the corners of a tetrahedron at angles of109.5.

Hybridization

Hybridization


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The bonding in methane. Each of the four C-H bonds

results from head-on (s) overlap of a singly occupiedcarbon sp3 hybrid orbital with a singly occupiedhydrogen 1s orbital. Sigma bonds are formed byhead-to-headoverlap between the hydrogen s orbital

and a singly occupied sp3

hybrid orbital of carbon.

Hybridization

sp hybridization


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sp hybridizationNow consider BeCl

2which has linear molecular

geometry determined experimentally.

The combination of one s and one p orbital gives twosp hybrid orbitals oriented 180 apart. Two

unhybridized p orbitals remain and are oriented at 90

sp2 hybridization


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sp hybridization

E.g. the molecular

geometry is trigonal

planar with bond angle =120. To explain its

geometry, we can use the

following rational.

sp2 signifies one s and two

p orbitals are combined.

sp3d hybrid Orbitals


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sp d hybrid Orbitals

For Hybridization scheme to correspond to the 5- and 6- electron-

group geometries of VSEPR theory, we need to go beyond s and porbitals and traditionally this meant including d orbitals.

We can achieve the five half-filled orbitals and trigonal-bipyramidal

molecular geometry through the hybridization of one s, three p and

one d orbitals of valence shell into five sp3

d hybrid orbitals.

d h b id bi l


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sp3d2 hybrid Orbitals

In the same way, we can achieve the six half-filledorbitals and octahedral geometry through thehybridization ofone s, three p and two d orbitals ofvalence shell into six sp3d2 hybrid orbitals.


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Hybridisation of Elements

Shape ofmolecules/ions

Hybridisation type

AtomicOrbitals

Examples

Square

Planar

dsp2 d + s + p(2) [Ni(CN)4]2-

[Pt(Cl)4]2-

Trigonal

bipyramidal

sp3d s + p(3) + d PF5, PCl5

Squarepyramidal

sp3d2 s + p(3) +d(2)

BrF5

Octahedral sp3

d2

s + p(3) + SF6 ,


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Covalent Bonding and Orbital


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Lewis structures and VSEPR do not explain why a bond

forms.

How do we account for shape in terms of quantum

mechanics?

What are the orbitals that are involved in bonding?

We use Valence Bond Theory:

Bonds form when orbitals on atoms overlap.

There are two electrons of opposite spin in the orbital overlap.

Covalent Bonding and OrbitalOverlap

Covalent Bonding and Orbital


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Covalent Bonding and OrbitalOverlap


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Molecular Shapes

Linear: Atoms lie in a straight line,bond angle is 180o

Planar triangular: Atoms are locatedon the corners of a triangle. Bondangles are all 120o

Tetrahedral: Atoms are located on

the corners of a tetrahedron. Bondangles are all 109.5o


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Trigonal bipyramid: Consists of two triangularpyramids that share a common base.

Consists of two types of bonds: a) equatorial bonds:bond angles are 120o and b) axial bonds: bonds are

180o from each other, but 90o between each equatorialbond

Octahedral: Two square pyramids sharing a commonbase. All bond angles are 90o from each other.

Molecular Shapes

Electron Domain Geometry


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Electron Domain Geometry


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Molecular Shape and Molecular


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Molecular Shape and MolecularPolarity

Molecular Shape and Molecular


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o ecu a S ape a d o ecu aPolarity


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l d i


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Electron domains

Regions in space where groups ofelectrons can be found

Two types of electron domains:a) bonding: Electrons that are

involved between pairs of atoms

b) Nonbonding: contains valenceelectrons bonded to a single atom

B di d i


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Bonding domains

All the electrons within a given

single, double or triple bond areconsidered to be in the samedomain

N b di d i


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Nonbonding domains

These can consist of either alone (unshared) pair of

electrons, or a single unpairedelectron

Nonbonding domains affect the

shape of the molecule

Molecular shapes with four


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Molecular shapes with fourdomains

Tetrahedral

Trigonal pyramidalBent

Molecular shapes with five


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Molecular shapes with fivedomains

Trigonal bipyramid

Distorted tetrahedral(seesaw)

T-shaped

linear

Molecular shapes with six


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Molecular shapes with sixdomains

Octahedral

Square pyramidalSquare planar


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Polarity of polyatomic

molecules/ions

If all the atoms attached to the

central atom are not the same,or there are lone pairs in thecentral atom, the molecule is

usually polar.

M l l Sh


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A molecule is usually non-polar if

a) all the bonds are non-polar or

b) There are n lone pairs in the valenceshell of the central atom and all atomsattached to the central atom are thesame

Molecular Shapes

H d B di


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Hydrogen Bonding

Hydrogen bond is formed when a hydrogen atom findsitself between two highly electronegative atoms suchas F,O,N.

It may be intermolecular (existing between two ormore molecules of the same or different substances)or intramolecular (present within the same molecule).

Hydrogen bonds have a powerful effect on thestructure and properties of many compounds

VSEPR Model-HydrogenBonding


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The Effect of Nonbonding Electrons

By experiment, the H-X-H bond angle decreases on

moving from C to N to O:

Since electrons in a bond are attracted by two nuclei, they do

not repel as much as lone pairs.

Therefore, the bond angle decreases as the number of lone pairs

increases

Bonding

104.5O

107O

NH

HH

C

H

HHH

109.5O

OHH

VSEPR Model


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Shapes of Larger Molecules In acetic acid, CH3COOH, there are three central atoms.

VSEPR Model


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Ionic Bonding

Resonance Structures

VSEPR

Basic Shapes

3-D Notation

Hybridization

Molecular Geometries

Octet Rule Polar Molecules

Lewis Structures Covalent Bonding

Types of Bonds


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THANK YOU

xi chemistry unit 4

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