Bonding: General Concepts

Chapter 8

8.1 Types of Chemical Bonds8.2 Electronegativity8.3 Bond Polarity and Dipole Moments8.4 Ions: Electron Configurations and Sizes8.5 Energy Effects in Binary Ionic Compounds8.6 Partial Ionic Character of Covalent Bonds8.7 The Covalent Chemical Bond: A Model8.8 Covalent Bond Energies and Chemical Reactions8.9 The Localized Electron Bonding Model8.10 Lewis Structures8.11 Exceptions to the Octet Rule8.12 Resonance8.13 Molecular Structure: The VSEPR Model

Copyright © Cengage Learning. All rights reserved

3

A Chemical Bond• Forces that hold groups of atoms together and make them

function as a unit.• A bond will form if the energy of the aggregate is lower than

that of the separated atoms. Types of Chemical Bonds:• Ionic Bonding – electrons are transferred• Covalent Bonding – electrons are shared equally• Intermediate cases

Ionic Compound

Cation is always smaller than atom from which it is formed.Anion is always larger than atom from which it is formed.

Bond energy:

It is the energy required to break the bond.

Bond length:

It is the equilibrium distance where the energy of the system is minimum

How does a bonding force develop between two identical atoms ?

The Interaction of Two Hydrogen Atoms

When H atoms are brought close together, there are two unfavorable potential energy terms, proton-proton repulsion & electron-electron repulsion , and one favorable term, proton-electron attraction.

A bond will form (that is, the two H atoms will exist as a molecular unit)if the system can lower its total energy in the process .

Energy profile as a function of the distance between the nuclei of the hydrogen atoms.

The zero point of energy is :” the atoms at infinite separation . At very short distances , when the atoms are very close together Bond length is the distance at which the system has minimum energy .

The potential energy of each electron is lowered because of the increased attractive forces in this area

In the hydrogen molecule and in many other molecules in which electrons are shared by nuclei is called Covalent bonding.

Covalent Bond No electron transferElectrons are shared between two atoms,

positioned between the two nucleielectrons are shared equally between identical

atomsExample: H2, O2, etc.

Polar Covalent Bond• Unequal sharing of electrons between atoms in a

molecule.• Results in a charge separation in the bond (partial

positive and partial negative charge).• When a sample of hydrogen fluoride HF gas is

placed in an electric field , the molecules tend to orient themselves, with the fluoride end closest to the positive pole and the hydrogen end closest to the negative pole.

H +δ- F -δ

Polar covalent bond or polar bond is a covalent bond with greater electron density around one of the two atoms. The molecule is called “Dipolar”.

H F FHe- poor e- rich

The Effect of an Electric Field on Hydrogen Fluoride Molecules

(a) When no electric field is present, the molecules are randomly oriented.

(b) When the field is turned on, the molecules tend to line up with their negative ends toward the positive pole and their positive ends toward the negative pole.

Electronegativity• Electronegativity

– the ability of an atom in a molecule to attract shared electrons to itself

– I.e., …how much an atom “wants” electrons within a bond

– determined by Linus Pauling (1901 - 1995)

to determine how polar a bond will be, we introduce

•The general trend is electronegativity increase as we go right and up in the periodic table

The Relationship Between Electronegativity and Bond Type

•Electronegativity difference increase polarity increases

• Arrange (Order)the following bonds in order of increasing polarity:– H-H, O-H, Cl-H, S-H, and F-H– Solution:– The polarity of the bond increases as the difference in the electronegativity

increases.• H-H < S-H < Cl-H < O-H < F-H (2.1) (2.1) (2.5) (2.1) (3.0) (2.1) (3.5) (2.1) (4.0) (2.1)

– 0 0.4 0.9 1.4 1.9

Example 8.1::

Electronegativity difference

Covalent bond polar covalent bond Polarity increases

Bond Polarity and Dipole MomentsA molecule with a center of negative charge and a center of positive charge is said to be dipolar or has a dipole moment.

For example hydrogen fluoride

HF behaves in electric field as if it had two centers of charge, H positive and F negative.

• Dipole Moment– a polar molecule has a dipole moment– a polar molecule has a center of positive charge and a center of negative

charge– Ex: H-F

– The dipolar character of a molecule is represented by an arrow pointing to the negative charge center with the tail of the arrow indicating the positive center of charge.

• Place a polar molecule in an electric field– the molecule will line up so that its “negative” end will line

up with the positive pole and the “positive end” will line up with the negative pole

Representation of dipole moment

•Molecules with polar bonds and have net dipole moment:

Water molecule is polar molecule

The water molecule has a dipole moment .

The same type of behavior is observed for the NH3

The structure and charge distribution of the ammonia molecule

The dipole moment of the ammonia molecule oriented in an electric field

Polar molecule It has dipole moment

• Some molecules have polar bonds, but are nonpolar – these molecules have no net dipole moment– due to the overall geometry of the molecule, the bond

polarities cancel out, so the molecule has no net dipole moment

• Ex: A- Linear molecule: e.g: CO2 (O=C=O)

c. Tetrahedral molecules:

e.g: CCl4 , and CH4

Tetrahedral molecules with four identical bonds 109.5 degrees apart

Example 8.2:For each of the following molecules, show the direction of the bond polarities and indicate which ones have a dipole moment:•HCl•Cl2

•CH4 ( tetrahedral molecule)

•H2S (V-shaped molecule)Answer:•HCl

The electronegativity of chlorine is greater than that of hydrogen (3.02.1). Thus the chlorine will be partially

negative, and the hydrogen will be partially positive. The HCl molecule has a dipole moment.

•Cl2

The two chlorine atoms share the electrons equally. No bond polarity occurs, and the Cl2 molecule has no dipole moment

•CH4 ( tetrahedral molecule)

Carbon has a slightly higher electronegativity (2.5) than does hydrogen (2.1). this lead to small partial positive charges on the hydrogen atoms and small partial negative charge on the carbon. Bond polarities cancel. The molecule has no dipole moment.

•H2S (V-shaped molecule)

The H2S molecule has a dipole

moment.

Ions: Electron Configurations and sizes:

•Atoms in stable compounds usually have a noble gas electron configuration.

• In ionic compounds ;The nonmetals form anions, and the metals form cations to achieve a noble gas electron configuration ( Na+ - Cl-)

•When two nonmetals react to form a covalent bond, they share electrons to complete the valence electron configuration of both atoms. That is, both nonmetals have noble gas electron configuration.

•When metal react with nonmetal to form ionic compound, the valence electron configuration of the nonmetal achieves the electron configuration of the next noble gas atom and the valence orbital of the metal are emptied. In this way both ions achieve noble gas electron configuration.

Electron Configurations of Cations and AnionsOf Representative Elements

Na [Ne]3s1 Na+ [Ne]Ca [Ar]4s2 Ca2+ [Ar]Al [Ne]3s23p1 Al3+ [Ne]

Atoms lose electrons so that cation has a noble-gas outer electron configuration.

Atoms gain electrons so that anion has a noble-gas outer electron configuration.

H 1s1 H- 1s2 or [He]

F 1s22s22p5 F- 1s22s22p6 or [Ne]

O 1s22s22p4 O2- 1s22s22p6 or [Ne]

N 1s22s22p3 N3- 1s22s22p6 or [Ne]

Predicting formulas of ionic compounds

When we speak in this text of the stability of an ionic compound, we are referring to the solid state.

• Group I forms +1 ions• Group II forms +2 ions• Group III forms +3 ions• Group VI forms -2 ions• Group VII forms -1 ions

Example of Ionic Compounds MgO magnesium oxide is formed of Mg2+ and O2-

Mg :[Ne]4s2 Mg2+ :[Ne]

CaO formed from Ca2+ and O2-. Ca :[Ar]4s2 Ca2+ :[Ar] Al2O3 is formed of 2Al3+ and 3O2-.

Al:[Ne]3s23p1 Al3+ :[Ne]

O: 1s22s22p4 O2- 1s22s22p6 or [Ne]

• Size of Ions– Positive ion (Cation)

• formed by the loss of an electron from the parent atom• cation is smaller than the parent due to the loss of electron pair

repulsion and the increased nuclear charge felt by each electron– Negative ion (Anion)

• formed by the parent atom gaining an electron• anion is larger than the parent atom due to increased electron

pair repulsion

Isoelectronic ions Ions containing the same number of electrons.The number of electrons and the number or protons are affect on the size of ions. But in case of isoelectronic ions, ions have the same number of electrons.

ions with the same number of electronsNa+ , Mg+2, Al+3, F-, O-2, N-3

all contain 10 electronsall have the same electron configuration as Nein terms of size, N-3>O-2>F->Na+>Mg+2>Al+3

the ion with the greater number of protons in an isoelectronic series will be the smallest due to the greater nuclear charge pulling the electrons in closer

8O-2

10 electrons 1s2 2s2 2p6

9F-

10 electrons 1s2 2s2 2p6

11Na+

10 electrons 1s2 2s2 2p6

12Mg+2

10 electrons 1s2 2s2 2p6

13Al+3

10 electrons 1s2 2s2 2p6

The attraction force between the 10 electrons and the positive charge on the nucleus increase with increase the nuclear charge ‘Z’ increases. Therefore the size of ions decreases as the nuclear charge ‘Z’ increases. Also we can see of isoelectronic ions decreases with increasing atomic number O-2> F- > Na+ > Mg+2 > Al+3 size of isoelectronic ions

• Example :Arrange the ions Se-2, Br-, Rb+, and Sr+2 in order of decreasing size

• Answer: This is an isoelectric series of ions with the krypton electron configuration. Since these ions all have the same number of electrons,

• Z (ion) 36 36 36 36 (no. of electrons)• Their sizes will depend on the nuclear charge .• Z (atom) 34 (Se2-) 35 (Br-) 37(Rb+) 38 (Sr2+) (no. of

protons) Since the nuclear charge is greatest for (Sr2+) , it is the smallest of these ions.

• The (Se2-) ion is largest. (Z increases and size decreases)

• Ion Se2- > Br- > Rb+ > Sr2+

•

Example 8.4Choose the largest ion in each of the following groups:•Li+, Na+, K+, Rb+, Cs+ (Group 1A)•Ba2+, Cs+, I-, Te2- (isoelectronic series)Answer:

•Since size increases down a group Cs+ is largest.•This is an isoelectronic series of ions with xenon electron configuration (size decreases with increasing Z):

Ion Te2- > I- > Cs+ > Ba2+

Z (atom) 52 53 55 56 (no. of protons)

Z (ion) 54 54 54 5 (no. of electrons)

The Covalent Chemical Bond: A model:

The Localized Electron Bonding Model• The model assumes that A molecule is composed of atoms that are

bound together by sharing pairs of electrons using the atomic orbitals of the bound atoms.

• electron pairs are localized on a particular atom or in the space between two atoms.

Localized Electron Bonding Model

• Lone pairs– electrons localized on a particular atom

• Bonding pairs (or shared pairs)– electrons localized in the space between two atoms

• On to Lewis structures!

Lewis Structure

The Lewis structure of a molecule Shows how valence electrons are arranged among atoms in a molecule.

Octet rule, atoms tend to combine in such a way that they each have eight electrons in their valence shells, giving them the same electronic configuration as a noble gas.

 

Hydrogen forms stable molecules where it shares 2 electrons. When 2 hydrogen atoms ,each with 1 electron , combine to form the H2 molecule .

H . . H

H : HBy sharing electrons , each H in H2 , in effect, has 2 electrons ,that is , each H has a filled valence shell .

.. .. .. .. .. ..

: . : . : : : : :

.. .. .. .. .. ..

H+H H:H or H-H

or F F F F F F

• In covalent bond formation, atoms go as far as possible toward completing their octets (duplets) by sharing electron pairs.

• Each fluorine atom also has three pairs of electrons not involved in bonding. These are the lone pairs .

Steps for Writing Lewis Structures1. Sum the valence electrons from all the atoms.2. Use a pair of electrons to form a bond between each pair of bound

atoms.3. Atoms usually have noble gas configurations. Arrange the remaining

electrons to satisfy the octet rule (or duet rule for hydrogen).

•The central atom is usually written first in the formula [the central atom is usually the least electronegative atom, exceptions in H2O, and NH3 where O and N are the central

atoms]•Hydrogen is never the central atom

Examples:

1- H2O (water): 1H and 8OSum of valence electrons = 1 + 1 + 6 = 8

HOH ..

..

Lewis structure of water

H O+ + H OH H O HHor

8e-2e-2e-

single covalent bonds

2- CO2 (carbon dioxide): 6C and 8OSum of valence electrons = 4 + 6 + 6 = 16

Double bond – two atoms share two pairs of electrons

O C O

8e- 8e- 8e-

or O C Odouble bonds

double bonds

3- CN- (Cyanide ion): 6C and 7N

Sum of valence electrons = 4 + 5 + 1 = 10

Triple bond – two atoms share three pairs of electrons

c N

triple bond8e- 8e-

or C N

triple bond4. NO+ (Nitrogen oxide ion) 7N and 8O

Sum of valence electrons = 5 + 6 – 1 = 10

[NO]+

5. N2 (Nitrogen) 7N and 7N

Sum of valence electrons = 5 + 5 = 10

NN

6. CO2

O C O....

....

Exceptions to the Octet RuleThe Incomplete Octet (Be &B)

Some atoms are satisfied with less than an octet-Be (1s22s2) is stable with only four valence electrons

- Boron (1s2 2s2 2p1) also tends to form compoundswith less than eight electrons

Odd-Electron Molecules:- Some molecules have an odd number of electronscan't satisfy octet rule; usually N has the odd number

The Expanded Octet:– Atoms in and beyond the 3rd period canhave more than eight electrons when in acompound– Thus, when drawing Lewis electron-dotformulas, extra electrons go on the centralatom.Example: SF6 [6 + 6 (7) = 48 valence electrons ]

I-3: 3(7) + 1 = 22 valence electrons I[ ]II

-:

..:

....

......

. .

PCl5 : 5 + 5(7) = 40 electrons

P

Cl

Cl

ClCl

Cl

::

:

:

: ::

..

..

..

..

..

....:

Consider the Lewis structure for the nitrate ion NO-3:

-

][ N

O

OO

::

...... ..

....Based on Lewis structure it should be two types of N…O

bonds. Experiments show that NO-

3 has only one type of N…

O bond with length and strength between those expected for a single and double bonds.

. This can be explained on the basis of the resonance structures of NO-

3:

Resonance: Blending of Structures

NO3- All valid. We cannot find two bond

lengths (hypothetical N-O vs N=O

Physical evidence shows that NO3- has three equivalent bonds

The correct description of NO3- is not one of the three Lewis structures,

but an “average” of the three Lewis structures

.... ....

....

: :

O O

O

N[ ]- -

][ N

O

OO

::

.... ..

....

-

][ N

O

OO

::

......

....

.. ..

Example: Given the Lewis structure for ozone, O3

resonance structure: one of two or more Lewis structures representing a single molecule that cannot be described fully with only one Lewis structure

Formal ChargeAn 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.

1. For neutral molecules, sum of formal charges must equal zero.2. For ions, the sum of the formal charges must equal charge.

Consider the Lewis structure for POCl3. Assign the formal charge for each atom in the molecule.

P: 5 – (4+0) = +1O: 6 – (6+1) = –1Cl: 7 – (6+1) = 0

P

Cl

Cl O

Cl

(Valence-Shell Electron-Pair Repulsion model)VSEPR Model

This model is useful in predicting the geometries of molecules formed from nonmetals. The structures of molecules play a very important role in determining their chemical properties . The main postulate of this model is that “ The structure around a given atom is determined principally by minimizing electron-pair repulsions.”

The bonding and nonbonding pairs around a given atom will be positioned as far apart as possible.

Number of electron pairs Structure Example

2 Linear

3 Trigonal planar

4 Tetrahedral

5 Trigonal bipyrimidal

6 Octahedral

Cl Be Cl

180o

B

F

FF

120o

C

H

HH

H

109.5o

Cl

ClCl

Cl

Cl

P

90o

120o

F

F

F F

F

F

S

90o

Predicting a VSEPR Structure• 1. Draw Lewis structure for the molecule.• 2. Put the electron pairs as far apart as possible.• 3. Determine positions of atoms fro the way electron pairs are shared.• 4. Determine the name of molecular structure from positions of the atoms.

The Bond Angles In the CH4, NH3, and H2O Molecules

bonding-pair vs. bondingpair repulsion

lone-pair vs. lone pairrepulsion

lone-pair vs. bondingpair repulsion

< <

Example: Predict the molecular structure and bond angles for the following molecule or ion:

1.HCN2.PH3

3.CHCl3

4.NH4+

5.H2CO2

6.CO2

Answer:

a) HCN•Draw the Lewis structure H-CN:2- count the electron pairs•2 electron pairs as single bonds•2 effective electrons as triple bond3-determine the position of atoms H-CN:4- determine the name of molecular structure•Linear , 180o

b) PH3 trigonal pyramid 109.5o

C)CHCl3 tetrahedral, 109.5o

d) NH4+ tetrahedral, 109.5o

e) H2CO trigonal planar, 120 o

f)CO2 linear, 180 o