slide1 n.
Download
Skip this Video
Loading SlideShow in 5 Seconds..
Chemical Bonding II PowerPoint Presentation
Download Presentation
Chemical Bonding II

Loading in 2 Seconds...

play fullscreen
1 / 71

Chemical Bonding II - PowerPoint PPT Presentation


  • 59 Views
  • Uploaded on

Chemical Bonding II. Lattice Energy. Remember, IE and EA are for adding/removing an electron to/from an atom in the gaseous state .

loader
I am the owner, or an agent authorized to act on behalf of the owner, of the copyrighted work described.
capcha
Download Presentation

PowerPoint Slideshow about 'Chemical Bonding II' - jafari


An Image/Link below is provided (as is) to download presentation

Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author.While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server.


- - - - - - - - - - - - - - - - - - - - - - - - - - E N D - - - - - - - - - - - - - - - - - - - - - - - - - -
Presentation Transcript
slide1

Chemical

Bonding

II

slide2

Lattice Energy

  • Remember, IE and EA are for adding/removing an electron to/from an atom in the gaseous state.
  • Ionic compounds are usually solids. The release of energy on forming the solid, called the lattice energy is the driving force for the formation of ionic compounds.
  • Because of high lattice energies, ionic solids tend to be hard and have high melting points. Ionic compounds are insulators in the solid state, because electrons are localized on the ions, but conduct when molten or in solution, due to flow of ions (not electrons).
  • Lattice energies can be calculated using Hess’s law, via a Born-Haber Cycle.
slide3

Figure 9.6

The Born-Haber cycle for lithium fluoride

slide4

Calculating Lattice Energy

  • Step 1: Convert elements to atoms in the gas state
    • e.g. for Li, Li (s)  Li (g) DH1 = DHatomization
    • for F, 1/2 F2 (g)  F (g) DH2 = 1/2 (Bond Energy)
  • Step 2: Electron transfer to form (isolated) ions
    • Li (g)  Li+ (g) + e–DH3 = IE1
    • F (g) + e– F– (g) DH4 = EA1
  • Step 3: Ions come together to form solid
    • Li+ (g) + F– (g) LiF (s) DH5 = Lattice Energy
  • Overall: Li (s) + 1/2 F2 (g) LiF (s) DH = DHf = S(DH1–5)
  • Lattice Energy = DHf – (DH1 + DH2 + DH3 + DH4)
slide5

Periodic Trends in Lattice Energy

Coulomb’s Law

charge A X charge B

electrostatic force a

distance2

charge A X charge B

or, electrostatic energy a

distance

(since energy = force X distance)

  • So, lattice energy increases, as ionic radius decreases (distance between charges is smaller).
  • Lattice energy also increases as charge increases.
slide6

Figure 9.7

Trends in lattice energy

slide7

I. Bonding Theory

A covalent H-H bond is the net result of attractive and repulsive electrostatic forces. When bringing together two atoms that are initially very far apart. Three types of interaction occur:

(1) The nucleus-electron attractions (blue arrows) are greater than the (2) nucleus-nucleus and (3) electron-electron repulsions (red arrows), resulting in a net attractive force that holds the atoms together to form an H2 molecule.

slide8

A graph of potential energy versus internuclear distance for the H2 molecule.

If the hydrogen atoms are too far apart, attractions are weak and no bonding occurs. A zero of energy when two H atoms are separated by great distances.

A drop in potential energy (net attraction) as the two atoms approach each other.

When the atoms are optimally separated, the energy is at a minimum. A minimum in potential energy at particular internuclear distance (74pm) corresponding to the stable H2 molecule and the potential energy corresponds to the negative of the bond dissociation energy.

If the atoms are too close, strong repulsions occur. A increase in potential energy as the atoms approach more closely.

ii valence bond theory
II. Valence-Bond Theory
  • bond formation by overlapping orbitals: A description of covalent bond formation in terms of atomic orbital overlap is called the valence bond theory. It gives a localized electron model of bonding: core electrons and lone-pair valence electrons retain the same orbital locations as in the separated atoms, and the charge density of the bonding electrons is concentrated in the region of orbital overlap.
  • 2)hybridization of atomic orbitals: How do a carbon with a s orbital and three p orbitals combined with four hydrogen (s orbitals) form four bonds and all four bonds are found to be 109.5?
slide10

1s

1s

Overlap of two half-filled orbitals leads to the formation of a covalent bond.

1s-1s overlap gives a H – H single bond

slide11

2s

2p

1s







F

H

The 1s-2p overlap gives a H – F single bond

slide12

2s

2p

1s







F

H

Non-bonding electrons

slide13

2s

2s

2p

2p













F

F

The 2p-2p overlap gives a F – F single bond

slide14

2s

2s

2p

2p













F

F

Non-bonding electrons

Each F atom has three pairs of non-bonding electrons.

slide15

2s

2s

2p

2p









O

O

Q.23 Identify the non-bonding electrons in O2 molecules.

Two 2p-2p overlaps give a O=O double bond

slide16

2s

2s

2p

2p









O

O

Q.23 Identify the non-bonding electrons in O2 molecules.

Non-bonding electrons

Each O atom has two pairs of non-bonding electrons.

slide17

Overlap of anempty orbitalwith afully-filled orbitalleads to the formation of aco-ordinate covalent bondordative bond

hybridization
Hybridization

In 1931, Linus Pauling proposed that the wave functions for the s

and p atomic orbitals can be mathematically combined to form a new

set of equivalent wave functions called hybrid orbitals.

The mathematical process of replacing pure atomic orbitals with

reformulated atomic orbitals for bonded atoms is called

hybridization.

In a hybridization scheme, the number of hybrid orbitals equals to the total number of atomic orbitals that are combined. The symbols identify the numbers and kinds orbitals involved.

slide22

(a) NH4+

By Lewis model, the structure is

 4 single bonds are formed,

one of them is a dative bond.

slide23

2s

2p

1s

1s



N

3H

H+

By VB Theory,

Three 2p-1s(half-filled) overlaps lead to the formation of three N – H single bonds.

slide24

2s

2p

1s

1s



N

3H

H+

By VB Theory,

One 2s(fully-filled)-1s(vacant) overlap leads to the formation of one N  H dative bond.

slide25

(b) HCN

By Lewis model, the structure is H-CN

 one H-C single bond and

one CN triple bond.

slide26

2s

2p

2s

2p



C*

By VB Theory,

C

 Only 2 single bonds can be formed.

Promotion of a 2s electron to a 2p orbital.

slide27

C*

N

H

2s

2s

2p

2p

1s



  • The overlap of one orbital (?) of C* with an 1s orbital of Hgives theC-H single bond.
  • Overlaps of three orbitals (???) of C* with three 2porbitals of N give theCN triple bond.
slide28

C*

N

H

2s

2s

2p

2p

1s



  • The 2s electrons on N are non-bonding electrons.
  • The energy released by forming a stronger triple bond outweighs the energy required for promoting an electron from a 2s orbital to a 2p orbital.
slide29

Most stableno separation of charge.

(c) SO2

By Lewis model, the three possible structures are

OS=O, O=SO, O=S=O

slide30

3s

3p





S

By VB Theory,

 Only two single bonds can be formed.

 One 3p electron has to be promoted to a 3d orbital.

 Expansion of Octet.

slide31

3d

3s

3s

3p

3p







S

S*

octet expansion

By VB Theory,

slide32

3d

2s

3s

2p

3p







S*

2O

 Overlaps of two half-filled orbitals (??) of S* with two half-filled 2p orbitals of an oxygen atom give a S=O double bond.

A total of two S=O bonds are formed with two O atoms

slide33

3d

2s

3s

2p

3p







S*

2O

Non-bonding electrons :

S* 3s2 ;

O 2s2 and 2p2

slide34

3d

3s

3s

3p

3p







S

S*

octet expansion

The energy released by forming of two stronger double bonds outweighs the energy required for promoting an electron from a 3p orbital to a 3d orbital.

the concept of resonance
The Concept of Resonance

According to VB theory, the two less stable structures of SO2,

OS=O and O=SO do ‘exist’.

Each of these structures contributes in certain extent to the real structure of SO2.

slide36

If represents the wave function of the real structure of SO2 molecules, then

where

are the wave functions of the three possible structures and

a > b = c > 0

slide37

More contribution

Less contribution

In other words, the real structure of SO2 is the resonance hydrid of the three possible structures.

O=S=O  OS=O  O=SO

slide38

2s

2s

3s

2p

2p

3p















O

S

O*

Q.24

O=SO

A S=O double bond is formed by 3p(half-filled)-2p(half-filled) overlaps between S and O.

slide39

2s

2s

3s

2p

2p

3p















O

S

O*

Q.24

O=SO

A OS dative bond is formed by 3p(fully-filled)-2p(empty) overlap between S and O*

slide40

2s

2s

3s

2p

2p

3p















O

S

O*

Q.24

O=SO

Formation of dative bond is not favourable because the two unpaired 2p electrons in O are forced to pair up to give O*

slide41

F-S-F

(d) SF2, SF4, SF6

Molecule

SF2

SF4

SF6

Most stable Lewis Structure

slide42

3s

2s

2p

3p











F

S

By VB Theory,

Only two S-F single bonds can be formed by 3p-2p overlaps between one S atom and two F atoms

 SF2 is formed.

F-S-F

slide43

3d

3s

2s

3s

3p

2p

3p













F

S

S*

By VB Theory,

To form four S-F single bonds in SF4, a 3p electron in S has to be promoted to a 3d orbital.

slide44

2s

3s

2p

3p











F

S

3d

3s

3p

S**

By VB Theory,

To form six S-F single bonds in SF6, a 3s electron in S* has to be promoted to a 3d orbital.

slide45

3s

3p





S

3d

3s

3p

S**

By VB Theory,

The energy released by forming more single bonds outweighs the energy required for promoting 3s and 3p electrons to 3d orbitals.

slide46

F-Xe-F

Q.25

Molecule

XeF2

XeF4

XeF6

Most stable Lewis Structure

slide47

5s

2s

5p

2p















Xe

F

5d

5s

5p







Xe*

By VB Theory,

To form two Xe-F bonds in XeF2, a 5p electron in Xe has to be promoted to a 5d orbital.

slide48

5d

5d

5s

5s

5p

5p











Xe**

Xe*

By VB Theory,

To form four Xe-F bonds in XeF4, a 5p electron in Xe* has to be promoted to a 5d orbital.

slide49

5d

5s

5p

5d

5s

5p





Xe**



Xe***

By VB Theory,

To form six Xe-F bonds in XeF6, a 5p electron in Xe** has to be promoted to a 5d orbital.

slide50

5d

5s

5p

5d

5s

5p





Xe**



Xe***

By VB Theory,

The energy released by forming more single bonds outweighs the energy required for promoting 5p electrons to 5d orbitals.

slide51

E.g. sp3 signifies one s and three p orbitals are combined.

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

slide52

The formation of four sp3 hybrid orbitals by combination of an atomic s orbital with three atomic p orbitals. Each sp3 hybrid orbital has two lobes, one of which is larger than the other. The four large lobes are oriented toward the corners of a tetrahedron at angles of 109.5°.

slide53

The bonding in methane. Each of the four C-H bonds results from head-on (s) overlap of a singly occupied carbon sp3 hybrid orbital with a singly occupied hydrogen 1s orbital. Sigma bonds are formed by head-to-head overlap between the hydrogen s orbital and a singly occupied sp3 hybrid orbital of carbon.

sp 2 hybridization
sp2 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.

sp hybridization
sp hybridization

Now consider BeCl2 which has linear molecular geometry determined experimentally.

In hybridization scheme that best describes this compound is that

The combination of one s and one p orbital gives two sp hybrid orbitals oriented 180° apart. Two unhybridized p orbitals remain and are oriented at 90° angles to the sp hybrids.

sp 3 d hybrid orbitals
sp3d hybrid Orbitals

To described hybridization scheme to correspond to the 5- and 6- electron-group geometries of VSEPR theory, we need to go beyond s and p orbitals 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 sp3d hybrid orbitals.

sp 3 d 2 hybrid orbitals
sp3d2 hybrid Orbitals

In the same way, we can achieve the six half-filled orbitals and octahedral geometry through the hybridization of one s, three p and two d orbitals of valence shell into six sp3d2 hybrid orbitals.

molecular orbital theory
Molecular Orbital Theory

The goal of molecular orbital theory is to describe molecules in a similar way to how we describe atoms, that is, in terms of orbitals, orbital diagrams, and electron configurations.

forming a covalent bond
Forming a Covalent Bond
  • Molecules can form bonds by sharing electron
    • Two shared electrons form a single bond
  • Atoms can share one, two or three pairs of electrons
    • forming single, double and triple bonds
  • Other types of bonds are formed by charged atoms (ionic) and metal atoms (metallic).
atomic and molecular orbitals cont d
Atomic and Molecular Orbitals (cont’d)
  • Orbital Mixing
    • When atoms share electrons to form a bond, their atomic orbitals mix to form molecular bonds. In order for these orbitals to mix they must:
      • Have similar energy levels.
      • Overlap well.
      • Be close together.

This is and example of orbital mixing. The two atoms share one electron each from there outer shell. In this case both 1s orbitals overlap and share their valence electrons.

http://library.thinkquest.org/27819/ch2_2.shtml

atomic and molecular orbitals
Atomic and Molecular Orbitals
  • In atoms, electrons occupy atomic orbitals, but in molecules they occupy similar molecular orbitals which surround the molecule.
  • The two 1s atomic orbitals combine to form two molecular orbitals, one bonding (s) and one antibonding (s*).
  • This is an illustration of molecular orbital diagram of H2.
  • Notice that one electron from each atom is being “shared” to form a covalent bond. This is an example of orbital mixing.

http://www.ch.ic.ac.uk/vchemlib/course/mo_theory/main.html

molecular orbital theory1
Molecular Orbital Theory
  • Each line in the diagram represents an orbital.
  • The molecular orbital volume encompasses the whole molecule.
  • The electrons fill the molecular orbitals of molecules like electrons fill atomic orbitals in atoms
molecular orbital theory2
Molecular Orbital Theory
  • Electrons go into the lowest energy orbital available to form lowest potential energy for the molecule.
  • The maximum number of electrons in each molecular orbital is two. (Pauli exclusion principle)
  • One electron goes into orbitals of equal energy, with parallel spin, before they begin to pair up. (Hund's Rule.)
molecular orbital diagram h 2
Molecular Orbital Diagram (H2)

http://www.ch.ic.ac.uk/vchemlib/course/mo_theory/main.html

mo diagram for o 2
MO Diagram for O2

http://www.chem.uncc.edu/faculty/murphy/1251/slides/C19b/sld027.htm

molecular orbital diagram hf
Molecular Orbital Diagram (HF)

http://www.ch.ic.ac.uk/vchemlib/course/mo_theory/main.html

molecular orbital diagram ch 4
Molecular Orbital Diagram (CH4)

So far, we have only look at molecules with two atoms. MO diagrams can also be used for larger molecules.

http://www.ch.ic.ac.uk/vchemlib/course/mo_theory/main.html

conclusions
Conclusions
  • Bonding electrons are localized between atoms (or are lone pairs).
  • Atomic orbitals overlap to form bonds.
  • Two electrons of opposite spin can occupy the overlapping orbitals.
  • Bonding increases the probability of finding electrons in between atoms.
  • It is also possible for atoms to form ionic and metallic bonds.