Chapter 13

Bonding:

General Concepts

Types of Chemical Bonds

Ionic bonding

Polar covalent bonding

Covalent bonding

Lennard-Jones 6-12 potential

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Ionic bonding

Ionic substances are formed when an atom that loses electrons relatively easily react with an atom that has a high affinity for electrons.

ex. metal-nonmetal compound

Covalent Bonding

Electron are shared by nuclei

Polar Covalent Bonding

A polar bond is a covalent bond in which there is a separation of charge between one end and the other , in other words in which one end is slightly positive and the other slightly negative.

Electronegativity

The ability of an atom in a molecule to attract shared electrons to itself.

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1)(

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B andA elementsbetween differenceativity electroneg2/11

BBEAAEBAE

molKJxx

AB

ABBA

Calculate Electronegativity H atom

∆HF=565-1/2(432+154)=272

|xH-xF|=0.102(272)1/2=1.68, xH-4.0=-1.7, xH=2.3 O atom

∆OF=190-1/2(146+154)=40

|xO-xF|=0.102(40)1/2=0.65, xO-4.0=-0.65, xO=3.4 C atom

∆CF=485-1/2(347+154)=234.5

|xC-xF|=0.102(234.5)1/2=1.6, xC-4.0=-1.6, xC=2.4

Bond Polarity and Dipole Moments

Dipole Moment

μ=QR

Q: center of charge of magnitude

R: distance

Dipole Moment of HF

1D=3.336×10-30 coulomb meterμ=(1.6×10-19 C)(9.17×10-11 m)=1.47×10-29

=4.4 D for fully ionicMeasured dipole moment=1.83 D1.83×3.336×10-30=δ(9.17×10-11)δ=6.66×10-20

Ionic character=1.83/4.4=41.6%

In practice no bond is totally ionic. There will always be a small amount of electron sharing.  

The compounds with more 50% ionic character are normally considered to be ionic solids.

Dipole Moment of Polyatomic Molecules

For dipole moment of polyatomic molecules, the dipole is the geometric sum of all bond dipole moment.

Achieving Noble Gas Electron Configurations (NGEC)

Two nonmetals react: They share electrons to achieve NGEC.

A nonmetal and a representative group metal react (ionic compound): The valence orbitals of the metal are emptied to achieve NGEC. The valence electron configuration of the nonmetal achieves NGEC.

Isoelectronic Ions

Ions containing the the same number of electrons

O2> F > Na+ > Mg2+ > Al3+

largest smallest

Formation of Binary Ionic Compounds

Lattice energy: The change in energy that takes place when separated gaseous ions are packed together to form an ionic solid.

M+(g)+X-(g) →MX(s)

Formation of an Ionic Solid

1. Sublimation of the solid metal

M(s) → M(g) [endothermic] 2. Ionization of the metal atoms

M(g) →M+(g) + e- [endothermic] 3. Dissociation of the nonmetal

1/2X2(g) → X(g) [endothermic]

Formation of an Ionic Solid(continued)

4. Formation of X ions in the gas phase:

X(g) + e- → X-(g) [exothermic] 5. Formation of the solid MX

M+(g) + X-(g) → MX(s) [quite exothermic]

Sublimation of Li

Ionization of Li

Dissociation of F2

Electron affinity of F

Formation of solid

Lithium-Fluoride structure

Lattice Energy Calculations

k: a proportionality constant that depends on the structure of the solid and the electron configuration of the ions

Q1 and Q2: charges on the ions

r: the shortest distance between the centers of cations and anions

)(Energy Lattice 21

r

QQk

Lattice Energies and the Strength of the Ionic Bond

The strength of the bond between the ions of opposite charge in an ionic compound depends on the charges on the ions and the distance between the centers of the ions when they pack to form a crystal.

An estimate of the strength of the bonds in an ionic compound can be obtained by measuring the lattice energy of the compound.

Lattice Energies for Alkali Metals Halides

The bond between ions of opposite charge is strongest when the ions are small.

The lattice energies for the alkali metal halides is therefore largest for LiF and smallest for CsI.

)(Energy Lattice 21

r

QQk

Lattice Energies of Alkali Metals Halides (kJ/mol)

F- Cl- Br- I-

Li+ 1036 853 807 757

Na+ 923 787 747 704

K+ 821 715 682 649

Rb+ 785 689 660 630

Cs+ 740 659 631 604

Lattice Energies for Salts of the OH- and O2- Ions

The ionic bond should also become stronger as the charge on the ions becomes larger.

The lattice energies for salts of the OH- and O2- ions increase rapidly as the charge on the ion becomes larger.

Lattice Energies of Salts of the OH- and O2- Ions (kJ/mol)

OH- O2-

Na+ 900 2481

Mg2+ 3006 3791

Al3+ 5627 15,916

Lattice Energies and Solubility

The lattice energy of a salt gives a rough indication of

the solubility of the salt in water because it reflects the energy needed to separate the positive and negative ions in a salt.

Sodium and potassium salts are soluble in water because they have relatively small lattice energies. Magnesium and aluminum salts are often much less soluble because it takes more energy to separate the positive and negative ions in these salts.

NaOH is very soluble in water (420 g/L), but Mg(OH)2 dissolves in water only to the extent of 0.009 g/L, and Al(OH)3 is essentially insoluble in water.

The Covalent Chemical Bond

Bond Energy of CH4

Experimental result : 1652 kJ/mol

C(g)+4H(g) →CH4(g) + 1652 kJ/mol

An average C-H bond energy per mole

of C-H bond: 1652/4=413 (kJ/mol)

Stepwise Decomposition of CH4

CH4(g) →CH3(g)+H(g) 435 kJ/mol

CH3(g) →CH2(g)+H(g) 453 kJ/mol

CH2(g) →CH(g)+H(g) 425 kJ/molCH(g) →C (g)+H(g) 339 kJ/mol

Bond Energies

Bond breaking requires energy (endothermic). Bond formation releases energy (exothermic).

H = D(bonds broken) D(bonds formed)

energy required energy released

Bond Energy of CH3Cl

C(g)+Cl(g)+3H(g) →CH3Cl(g)+1578kJ/mol

(C-Cl)+3(C-H)=1578

(C-Cl)+3(413)=1578

C-Cl=339 (kJ/mol)

Covalent Bond Energies and Chemical Reactions

H2+F2→2HF

ΔH=ΣD (bonds broken)-ΣD (bonds formed)

ΔH=DH-H+DF-F-2DH-F=1×432+1×154-2×565

=-544 kJ

CH4+2Cl2+2F2→CF2Cl2+2HF+2HCl

Reactants bonds broken:

CH4: 4×413=1652,

2Cl2: 2×239=478

2F2: 2×154=308 Total energy required: 2438kJ

Products bonds formed:

CF2Cl2: 2×485=970 (C-F) and 2×339=678 (C-Cl)

HF: 2×565=1130

HCl: 2×427=854 Total energy released: 3632 kJ

ΔH=2438-3632=-1194 kJ (-1126 kJ)

Localized Electron Model

A molecule is composed of atoms that are bound together by sharing pairs of electrons using the atomic orbitals of the bound atoms.

Description of valence electron arrangement (Lewis structure).

Prediction of geometry (VSEPR model). Description of atomic orbital types used to sha

re electrons or hold long pairs.

Lewis Structure

Shows how valence electrons are arranged among atoms in a molecule.

Reflects central idea that stability of a compound relates to noble gas electron configuration.

Resonance

Occurs when more than one valid Lewis structure can be written for a particular molecule.

These are resonance structures. The actual structure is an average of the resonance structures.

Comments About the Octet Rule

2nd row elements C, N, O, F obey the octet rule.

2nd row elements B and Be often have fewer than 8 electrons around themselves - they are very reactive.

3rd row and heavier elements CAN exceed the octet rule using empty valence d orbitals.

When writing Lewis structures, satisfy octets first, then place electrons around elements having available d orbitals.

Formal Charge

One method involves estimating the charge on each atom in the various possible Lewis structures and using the charges to select the most appropriate structure.

It allows chemists to determine the location of charge in a molecule as well as compare how good a Lewis structure might be.

Calculation of Formal Charge

Formal Charge=(number of valence electrons on a free atom)-(number of valence electrons assigned to the atom in the molecule)

Assumptions for Formal Charge

Lone pair electrons belong entirely to the atom in question.

Shared electrons are divided equally between the two sharing atoms.

(Valence electrons) assigned=(number of lone

pair electrons)+1/2(number of shared

electrons)

Consider the molecule H2CO2

00

0

-1

+1

Consider the molecule H2CO2

0

0

0

00

The two possible Lewis structures are

shown above. They are connected by a

double headed arrow and placed in brackets.

The non-zero formal charge on any atoms

in the molecule have been written near the

atom.

Valence Shell Electron Pair Repulsion (VSEPR Model)

It is used to predict the geometries of molecules formed from nonmetals.

Postulate: the structure around a given atom is determined principally by minimizing electron pair repulsion.

The bonding and nonbonding pairs should be positioned as far apart as possible.

Predicting a VSEPR Structure

Draw Lewis structure. Put pairs as far apart as possible. Determine positions of atoms from the way

electron pairs are shared. Determine the name of molecular structure from

positions of the atoms.

For non-metals compounds, four pairs of

electrons around a given atom prefer prior

to form a tetrahedral geometry to minimize

the electron repulsions.

Draw the Lewis structure Count the pairs of electrons and arrange them to

minimize repulsions Determine the positions of the atoms Name the molecular structure

Lone pairs require more space than bonding pair.

The bonding pairs are increasingly squeezed together as the number of lone pairs increases.

The bonding pair is shared between two nuclei; and the electrons can be close to either nucleus.

A lone pair is localized on only one nucleus, so both electrons are close to that nucleus only.

Lone pairs require more room than bonding pairs

square planar