by: dr. ashish kumar. chemical bonding. valence electrons are the outer shell electrons of an atom....

By: Dr. Ashish Kumar.

Chemical Bonding

Valence electrons are the outer shell electrons of an atom. The valence electrons are the electrons thatparticpate in chemical bonding.

1A 1ns1

2A 2ns2

3A 3ns2np1

4A 4ns2np2

5A 5ns2np3

6A 6ns2np4

7A 7ns2np5

Group # of valence e-e- configuration

By: Dr. Ashish Kumar .

9.1

Lewis Dot Symbols for the Representative Elements &Noble Gases


Chemical bonds

• A chemical bond is the force that holds two or more atoms together.

• Chemical bonds involve the electrons.

• A bond results if a more stable electron configuration results.

• The valence electrons are the electrons in the outer s and p subshells.


Formation of Ions and Ionic Compounds

• Metals can lose electrons to form ions

Na ([Ne]3s1) Na+ ([Ne]) + e-

If a metal loses all of its outer electrons, it acquires the octet of the previous noble gas

• Nonmetals can gain electrons to form ions

Cl ([Ne]3s23p5 + e- Cl ([Ne]3s23p6

• Lewis dot structures of the atoms can be very helpful here.


Forming ionic compounds

• Reaction of Na with Cl

• Na donates an electron to Cl

• Na+ has the previous noble gas structure (Ne)

• Cl- has the next noble gas structure (Ar)


Binary ionic compounds

• In NaCl, each Na+ is surrounded by six Cl-, and each Cl- is surrounded by six Na+.

• Ionic lattice is a three-dimensional array of ions.

• These electrostatic attractions are called ionic bonds.


Ionic Bonds

An ionic bond forms by the transfer of electrons from the valence shell of an atom of lower electronegativity to the valence shell of an atom of higher electronegativity.

We show the transfer of a single electron by a single-headed (barbed) curved arrow.

+ F-(1s22s22p6)Na+(1s22s22p6)F(1s22s22p5)+Na(1s22s22p63s1)


9.2

Li + F Li+ F -

The Ionic Bond

1s22s11s22s22p5 1s21s22s22p6[He][Ne]

Li Li+ + e-

e- + F F -

F -Li+ + Li+ F -


Covalent Bonds

A covalent bond forms when electron pairs are shared between two atoms whose difference in electronegativity is 1.9 or less.

An example is the formation of a covalent bond between two hydrogen atoms.

The shared pair of electrons completes the valence shell of each hydrogen.


.

Formation of a covalent bond between two hydrogen atoms

(a) Two H atoms separated by a large distance. (b) As the atoms approach each other, their electron densities are pulled into the region

between the two nuclei. (c) In the H2 molecule, the electron density is

concentrated between the nuclei. Both electrons in the bond are distributed over both nuclei.


v

Changes in the potential energies of two hydrogen atoms as they form H2

The energy of the molecule reaches a minimum when there is a balance between the attractions

and repulsions.


A covalent bond is a chemical bond in which two or more electrons are shared by two atoms.

Why should two atoms share electrons?

F F+

7e- 7e-

F F

8e- 8e-

F F

F F

Lewis structure of F2

lone pairslone pairs

lone pairslone pairs

single covalent bond

single covalent bond

9.4By: Dr. Ashish Kumar .

8e-

H HO+ + OH H O HHor

2e- 2e-

Lewis structure of water

Double bond – two atoms share two pairs of electrons

single covalent bonds

O C O or O C O

8e- 8e-8e-double bonds double bonds

Triple bond – two atoms share three pairs of electrons

N N8e-8e-

N N

triple bondtriple bond

or


Lengths of Covalent Bonds

Bond Lengths

Triple bond < Double Bond < Single Bond By: Dr. Ashish Kumar .

H F FH

Polar covalent bond or polar bond is a covalent bond with greater electron density around one of the two atoms

electron richregion

electron poorregion e- riche- poor

d+ d-


Polar Covalent Bonds

In a polar covalent bond

• The more electronegative atom has a partial negative charge, indicated by the symbol d-.

• The less electronegative atom has a partial positive charge, indicated by the symbol d+.

In an electron density model

• Red indicates a region of high electron density.

• Blue indicates a region of low electron density.


Nonpolar and polar covalent bonds(a) The electron density of the electron pair in the

bond is spread evenly between the two H atoms in H2, which gives a nonpolar covalent bond. (b) In HCl, the

electron density of the bond is pulled more tightly around the Cl end of the molecule, causing that end of

the bond to become slightly negative. At the same time, the opposite end of the bond becomes slightly

positive. The result is a polar covalent bond.

Nonpolar, polar and ionic bonds


Coordinate or Dative Bond

Both electrons of the shared pair are provided by one of the combining atoms.

Formation of Ammonium (NH4 +) Ion

During the formation of ammonium ion, nitrogen is the donor atom, while H+ is the acceptor ion as shown:


Electronegativity is the ability of an atom to attract toward itself the electrons in a chemical bond.

Electron Affinity - measurable, Cl is highest

Electronegativity - relative, F is highest

X (g) + e- X-(g)


Linus Pauling

Electronegativity is the ability of an atom to attract toward itself the electrons in a chemical bond.


9.5

The Electronegativities of Common Elements


The Pauling electronegativity (EN) scale.


Pauling scale : Pauling scale of electronegativity is the most widely used. It is based on excess bond energies. He determined electronegativity difference between the two atoms and then by assigning arbitrary values to few elements (e.g. 4.00 to fluorine, 2.5 to carbon and 2.1 to hydrogen), he calculated the electronegativity of the other elements. χA – χB = 0.208 √ ∆ where χA and χB are electronegativities of the atoms A and B respectively, the factor 0.208 arises from the conversion of kcal to electron volt (1 eV = 23.0 kcal/mole),

while ∆ = Actual bond energy – √(EA–A × EB–B)


Pauling calculation in SI unit

XA – XB = 0.208 (D(A-B))½ kcal/mol

XA – XB = 0.1017 (D(A-B))½ kJ/mol

D(A-B) is the extra bond energy in kJ/mol (D(A-B) is in eV)

A-B A+ B-

0.1017 is a conversion from kJ/mol to eV

∆ = E(A-B)actual - E(A-B)theort.

∆ = Actual bond energy – √(EA–A × EB–B)

Bond (D kj/mol) 0.1017 √ D

C–H 24.3 0.50 i.e XC – XH = 0.50

H–Cl 102.3 1.02 i.e XCl – XH = 1.02

N–H 105.9 1.04 i.e XN– XH = 1.04

Electronegativity and atomic size.

Covalent

share e-

Polar Covalent

partial transfer of e-

Ionic

transfer e-

Increasing difference in electronegativity

Classification of bonds by difference in electronegativity

Difference Bond Type

0 Covalent

2 Ionic

0 < and <2 Polar Covalent


Classify the following bonds as ionic, polar covalent, or covalent: The bond in CsCl; the bond in H2S; andthe NN bond in H2NNH2.

Cs – 0.7 Cl – 3.0 3.0 – 0.7 = 2.3 Ionic

H – 2.1 S – 2.5 2.5 – 2.1 = 0.4 Polar Covalent

N – 3.0 N – 3.0 3.0 – 3.0 = 0 Covalent


Fajan’s ruleThe nature of bond between these two ions — effect of one ion on the other.

Na+ Cl–Cl–

Positive ion attracts the outermost electron of anion and repels its nucleus, causes distortion or polarisation of the anion.

Na+Cl–

Polarisation leads to the partial sharing of electron cloud.

Partialcovalent nature


Fajan’s rule

1. For a given cation, covalent character increases with increasing anion size.

2. For a given anion, covalent character increases with decreasing cation size.

3. The covalent character increases with increasing charge on either ion.

4. Covalent character is greater for cations with non-noble gas electronic configurations.


Factors responsible for partial covalent nature

(a)

Na+ Cl–Cl–Al+3

Smaller size and increase in charge density on cation, increases its polarising power.

AlCl3 is more covalent than NaCl

mp 180oC 810oC



Decreasing order of covalent nature

NaI > NaBr > NaCl

(b)

Na+ Cl–Cl– Br– I–

Larger anion holds its valence electronloosely — more polarisation occurs.



(c) Cu+ [Ar]3d10

(Pseudo noble gas configuration)

Na+ [Ne](Noble gas configuration)

Na+ Cu+

(0.95 ) (0.96 )oA

oA

No. of protons 11 29

More polarising powerof Cu+ — increase in effective nuclear

charge

Size andcharge arealmostsame.


Question


The decreasing ionic character of NaF, MgF2 and AlF3 is

(a) NaF > MgF2 > AlF3

(b) AlF3 > MgF2 > NaF

(c) MgF2 > AlF3 > NaF

(d) None of these

Illustrative example 1


Solution

Increase in charge on cation increases covalent nature.

Therefore, increasing order of covalent characterNa+ < Mg+2 < Al+3

Therefore, decreasing order of ionic characterNaF > MgF2 > AlF3

Hence, answer is (a).


Factors responsible for pure ionic bond

1. Large cation

2. Small anion

3. Small charge density on ions

4. Noble gas configuration


1. Draw skeletal structure of compound showing what atoms are bonded to each other. Put least electronegative element in the center.

2. Count total number of valence e-. Add 1 for each negative charge. Subtract 1 for each positive charge.

3. Complete an octet for all atoms except hydrogen

4. If structure contains too many electrons, form double and triple bonds on central atom as needed.

Writing Lewis Structures


Write the Lewis structure of nitrogen trifluoride (NF3).

Step 1 – N is less electronegative than F, put N in center

F N F

F

Step 2 – Count valence electrons N - 5 (2s22p3) and F - 7 (2s22p5)

5 + (3 x 7) = 26 valence electrons

Step 3 – Draw single bonds between N and F atoms and complete octets on N and F atoms.

Step 4 - Check, are # of e- in structure equal to number of valence e- ?

3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons


Write the Lewis structure of the carbonate ion (CO32-).

Step 1 – C is less electronegative than O, put C in center

O C O

O

Step 2 – Count valence electrons C - 4 (2s22p2) and O - 6 (2s22p4) -2 charge – 2e-

4 + (3 x 6) + 2 = 24 valence electrons

Step 3 – Draw single bonds between C and O atoms and complete octet on C and O atoms.

Step 4 - Check, are # of e- in structure equal to number of valence e- ?

3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons

9.6

Step 5 - Too many electrons, form double bond and re-check # of e-

2 single bonds (2x2) = 41 double bond = 4

8 lone pairs (8x2) = 16Total = 24


9.7

Two possible skeletal structures of formaldehyde (CH2O)

H C O HH

C OH

An atom’s formal charge is the difference between the number of valence electrons in an isolated atom and the number of electrons assigned to that atom in a Lewis structure.

formal charge on an atom in a Lewis structure

=1

2

total number of bonding electrons( )

total number of valence electrons in the free atom

-total number of nonbonding electrons

-

The sum of the formal charges of the atoms in a molecule or ion must equal the charge on the molecule or ion.


H C O HC – 4 e-

O – 6 e-

2H – 2x1 e-

12 e-



formal charge on C = 4 -2 - ½ x 6 = -1

formal charge on O = 6 -2 - ½ x 6 = +1


=1

2




-

-1 +1


C – 4 e-

O – 6 e-

2H – 2x1 e-

12 e-



HC O

H

formal charge on C = 4 -0 - ½ x 8 = 0

formal charge on O = 6 -4 - ½ x 4 = 0


=1

2




-

0 0


Formal Charge and Lewis Structures

9.7

1. For neutral molecules, a Lewis structure in which there are no formal charges is preferable to one in which formal charges are present.

2. Lewis structures with large formal charges are less plausible than those with small formal charges.

3. Among Lewis structures having similar distributions of formal charges, the most plausible structure is the one in which negative formal charges are placed on the more electronegative atoms.

Which is the most likely Lewis structure for CH2O?

H C O H

-1 +1 HC O

H

0 0


Exceptions to the Octet Rule

The Incomplete Octet

H HBeBe – 2e-

2H – 2x1e-

4e-

BeH2

BF3

B – 3e-

3F – 3x7e-

24e-

F B F

F

3 single bonds (3x2) = 69 lone pairs (9x2) = 18

Total = 24


Exceptions to the Octet Rule

Odd-Electron Molecules

N – 5e-

O – 6e-

11e-

NO N O

The Expanded Octet (central atom with principal quantum number n > 2)

SF6

S – 6e-

6F – 42e-

48e-

S

F

F

F

FF

F

6 single bonds (6x2) = 1218 lone pairs (18x2) = 36

Total = 48


• The approximate shape of a molecule can be predicted from the Lewis dot structure using the Valence-Shell Electron-Pair Repulsion(VSEPR)model which depicts electrons in bonds and lone pairs as “electron groups” or “charge clouds” that repel one another and stay as far apart as possible.

– Every group of electrons, whether in a bond or a lone pair, counts as a charge cloud. (Multiple bonds count as one charge cloud.

– There are five basic shapes in the VSEPR model,

with some variations depending on how many charge clouds are in lone pairs vs. covalent bonds. The three-dimensional shape of a molecule is often important in determining its chemical behaviour, particular for biologically important molecules

VSEPR THEORY


The Basic VSEPR Shapes

A Balloon Analogy for VSEPR Theory


Valence shell electron pair repulsion (VSEPR) model:

Predict the geometry of the molecule from the electrostatic repulsions between the electron (bonding and nonbonding) pairs.

AB2 2 0

Class

# of atomsbonded to

central atom

# lonepairs on

central atomArrangement of electron pairs

MolecularGeometry

linear linear

B B


Shapes: Two Charge Clouds3D animation on BeCl2 molecule

2 bonds, 0 lone pairs (LP): linear, 180°


AB2 2 0 linear linear

Class

# of atomsbonded to

central atom

# lonepairs on


MolecularGeometry

VSEPR

AB3 3 0trigonal planar

trigonal planar


Shapes: Three Charge Clouds3D animation on BF3 molecule

3 bonds, 0 LP: trigonal planar, 120°• 2 bonds, 1 LP: bent, < 120°


Class

# of atomsbonded to

central atom

# lonepairs on


MolecularGeometry

VSEPR


trigonal planar

AB2E 2 1trigonal planar

bent


AB2 2 0 linear linear

Class

# of atomsbonded to

central atom

# lonepairs on


MolecularGeometry

VSEPR


trigonal planar

AB4 4 0 tetrahedral tetrahedral


Shapes: Four Charge Clouds3D animation on CH4 molecule

4 bonds, 0 LP: tetrahedral, 109.5°


bonding-pair vs. bondingpair repulsion

lone-pair vs. lone pairrepulsion

lone-pair vs. bondingpair repulsion> >

• 3 bonds, 1 LP: trigonal pyramidal, <109.5°• 2 bonds, 2 LP: bent, <109.5°


Class

# of atomsbonded to

central atom

# lonepairs on


MolecularGeometry

VSEPR

AB3E 3 1


tetrahedraltrigonal

pyramidal


Class

# of atomsbonded to

central atom

# lonepairs on


MolecularGeometry

VSEPR


AB3E 3 1 tetrahedraltrigonal

pyramidal

AB2E2 2 2 tetrahedral bent


Shapes: Four Charge Clouds3D animation on NH3 molecule

3D animation on H2O molecule


Summary of VSEPR Shapes


Summary of VSEPR ShapesComparsion

between shapes of molecules on the basis of VSEPR theory


http://www.mhhe.com/physsci/chemistry/animations/chang_7e_esp/bom3s2_7.swf

http://www.mhhe.com/physsci/chemistry/animations/chang_7e_esp/bom3s2_7.swf

Hybridization – mixing of two or more atomic orbitals to form a new set of hybrid orbitals.

1. Mix at least 2 nonequivalent atomic orbitals (e.g. s and p). Hybrid orbitals have very different shape from original atomic orbitals.

2. Number of hybrid orbitals is equal to number of pure atomic orbitals used in the hybridization process.

3. Covalent bonds are formed by:

a. Overlap of hybrid orbitals with atomic orbitals

b. Overlap of hybrid orbitals with other hybrid orbitals


Hybridisation:Hybridization may be defined as phenomenon of mixing of orbital's of an atom of nearly equal energy, giving rise to entirely new orbital's equal in number of mixing orbital's and have same energy.

– [For instance, carbon’s electron configuration has 2 half-filled p orbitals 90° from each other, yet carbon forms four bonds to hydrogen atoms to form CH4, with bond angles of 109.5 °.]

• The number of hybrid orbitals equals the number of atomic orbitals which combine; we need one hybrid orbital for each electron pair (bond or lone pair)

sp hybridisationsp2 hybridisationsp3 hybridisation

http://dwb4.unl.edu/chemAnime/SP11ORBD/SP11ORBD.html

http://dwb4.unl.edu/chemAnime/SP21ORBD/SP21ORBD.html

http://dwb4.unl.edu/chemAnime/SP3SORB1/SP3SORB1.html


sp3 Hybridization

• These 4 orbitals point to the corners of a tetrahedron (109.5°).• sp3 orbitals are used to describe the C atoms in CH4 and C2H6, the N in NH3, and the O in H2O.

When the s and all three p orbitals combine, the resulting hybrid orbitals are sp3 hybrid orbitals


sp3 Hybridization


sp3 Orbitals in Methane


sp3 Orbitals in Ammonia


sp3 Orbitals in Ethane, C2H6


sp2 HybridizationWhen the s and two of the three p orbitals combine,the resulting hybrid orbitals are sp2 hybrid orbitals

These 3 orbitals point to the corners of an equilateraltriangle (120°), with the unhybridized p orbital 90°from this plane.• sp2 orbitals are used to describe the B atom in BF3,and the C atoms in ethylene, C2H4.


sp2 Hybridization


sp2 Orbitals in BF3


sp2 Hybridization and π BondsWhen two sp2-hybridized atoms are joined, a doublebond is form from two types of overlap:– a sigma (σ) bond resulting from end-on-end overlap of the sp2 orbitals.– a pi (π) bond resulting from the side-to-side overlap of the p orbitals.


Single Bonds and Double Bonds in VB TheoryA sigma (σ) bond results from the end-to-end overlap of cylindrical (p, sp, sp2, sp3, etc.) or spherical (s) orbitals.– σ-bonds are cylindrically symmetrical, and there is free rotation around them.– All single bonds are σ-bonds.• A pi (π) bond results from the side-to-side overlap of p orbitals on sp2- or sp- hybridized atoms.– π-bonds have regions of electron density above and below the σ-bond axis; free rotation is not allowed around a π-bond.– A double bond is a σ-bond + a π-bond.– A triple bond is a σ-bond + 2 π-bonds.


sp2 Orbitals in Formaldehyde, CH2O


sp2 Hybridization


sp2 Orbitals in Ethylene, C2H4


sp HybridizationWhen the s and one of the three p orbitals combine,the resulting hybrid orbitals are sp hybrid orbitals.

These 2 orbitals point 180° from each other, with the unhybridized p orbitals at 90° angles from this line. sp orbitals are used to describe the Be atom in BeCl2, and the C atoms in acetylene, C2H2.


sp Hybridization


sp Orbitals in BeCl2


sp Orbitals in Acetylene, C2H2

# of Lone Pairs+

# of Bonded Atoms Hybridization Examples

2

3

4

5

6

sp

sp2

sp3

sp3d

sp3d2

BeCl2

BF3

CH4, NH3, H2O

PCl5

SF6

How do I predict the hybridization of the central atom?

1. Draw the Lewis structure of the molecule.

2. Count the number of lone pairs AND the number of atoms bonded to the central atom

Sigma (s) and Pi Bonds (p)

Single bond 1 sigma bond

Double bond 1 sigma bond and 1 pi bond

Triple bond 1 sigma bond and 2 pi bonds

How many s and p bonds are in the acetic acid(vinegar) molecule CH3COOH?

C

H

H

CH

O

O Hs bonds = 6 + 1 = 7

p bonds = 1

Molecular orbital theory – bonds are formed from interaction of atomic orbitals to form molecular orbitals.

O

O

No unpaired e-

Should be diamagnetic

Experiments show O2 is paramagnetic


Molecular Orbital Theory: Hund & MullikenThe valence bond model is easy to visualize, and works well for most molecules, but it does not describe magnetic and spectral properties well.Another model is used to explain these phenomena.

• In Molecular Orbital (MO) theory, electrons occupy molecular orbitals that belong to the entire molecule rather than to an individual atom.

• Atomic or hybrid orbitals overlap with each other to form molecular orbitals. This overlap results in both constructive and destructive interference between the orbitals.

• The number of molecular orbitals formed is the same as the number of atomic orbitals which are combined.


Constructive and Destructive Interference

• Additive combinations (resulting from constructive interference between atomic orbitals) form bonding molecular orbitals (σ, π).– These are lower in energy than the atomic orbitals.– The electrons in these orbitals spend most of their time in the region in between the nuclei, helping to bond the atoms together.

• Subtractive combinations (resulting from destructive interference between atomic orbitals) form antibonding molecular orbitals (σ*, π*).– These are higher in energy than the atomic orbitals.– The electrons in these orbitals can’t occupy the node in the central region and don’t contribute to bonding.

Energy levels of bonding and antibonding molecular orbitals in hydrogen (H2).

A bonding molecular orbital has lower energy and greater stability than the atomic orbitals from which it was formed.

An antibonding molecular orbital has higher energy and lower stability than the atomic orbitals from which it was formed. By: Dr. Ashish Kumar .


MO Theory and Other Diatomic Molecules

For homonuclear diatomic molecules, which are composed of two identical atoms (e.g., H2, N2, O2, F2).

• In the simple VB model, the electrons in the O2 molecule are all paired; experimentally, however, O2 is found to be paramagnetic.

• the σ2s and σ2s* orbitals are produced from overlapping 2s orbitals; the σ2p and σ2p* orbitals are produced from end-to-end overlap of 2p orbitals;

• the π2p and π2p* orbitals are produced from side-to-side overlap of 2p orbitals.

Two Possible Interactions Between Two Equivalent p Orbitals

1. The number of molecular orbitals (MOs) formed is always equal to the number of atomic orbitals combined.

2. The more stable the bonding MO, the less stable the corresponding antibonding MO.

3. The filling of MOs proceeds from low to high energies.

4. Each MO can accommodate up to two electrons.

5. Use Hund’s rule when adding electrons to MOs of the same energy.

6. The number of electrons in the MOs is equal to the sum of all the electrons on the bonding atoms.

Molecular Orbital (MO) Configurations


General molecular orbital energy level diagram for the second-period homonuclear diatomic molecules Li2, Be2, B2, C2, and N2.


bond order = 12

Number of electrons in bonding MOs

Number of electrons in antibonding MOs

( - )

bond order

½ 1 0½By: Dr. Ashish Kumar .


Why Doesn’t He2 Form?

Bond order= (bonding e 's) - (antibonding e 's)/2


MO Energy Diagrams for B2 through Ne2

Delocalized molecular orbitals are not confined between two adjacent bonding atoms, but actually extend over three or more atoms.

Example: Benzene, C6H6

Delocalized p orbitals


Electron density above and below the plane of the benzene molecule.
