PowerPoint to accompany Chapter 7 Basic Concepts of Chemical Bonding.

PowerPoint to accompany

Chapter 7

Basic Concepts of Chemical Bonding

Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia

Chemical Bonds

Three basic types of bonds: Ionic

Electrostatic attraction between ions

Covalent Sharing of electrons

Metallic Metal atoms bonded to several other

atoms

Figure 7.1


Ionic: Electrostatic attraction between ions

Mg+2O-2

K+2(Cr2O7)-2

Ni+2O-2


Covalent: Shared pairs of electrons between adjacent atoms

Br2 (l & g)C12H 22O 11 (s)

S8 (s)


Metallic: Shared conduction band electrons throughout crystal lattice

Mg0 (s)

Cu0 (s)

Au20 (s)


Lewis Symbols

The Lewis symbol of an element consists of the chemical symbol for the element plus a dot for each valence electron.

Sulfur, for example, has the electron configuration [Ne]3s23p4. Therefore it has the following Lewis symbol:

S• •

• •• •

which clearly depicts the six valence electrons.


Octet Rule

Atoms tend to gain, lose or share electrons until they are surrounded by eight valence electrons.

An octet of electrons consists of full s and p subshells in an atom.

Note: There are many exceptions to the octet rule, but it remains useful for introducing many important concepts of bonding.


Energetics of Ionic Bonding

As we have seen, it takes 495 kJ/mol to

remove the outer 3s1 electron from sodium.

We get 349 kJ/mol back by giving the 3p6

electron to chlorine.



These numbers don’t explain why the reaction of sodium metal and chlorine gas to form sodium chloride is so exothermic!

Figure 7.2



There must be a third piece to the puzzle.

What is as yet unaccounted for is the electrostatic attraction between the newly formed sodium cation and chloride anion.


Lattice Energy

This third piece of the puzzle is the lattice energy:

The energy required to completely separate a mole of a solid ionic compound into its gaseous ions.

The energy associated with electrostatic interactions is governed by Coulomb’s law:

Eel = Q1Q2d


Lattice Energy

Lattice energy increases with the charge on the ions.

It also increases with the decreasing size of ions.

Table 7.2



By accounting for all three energies (ionization energy, electron affinity, and lattice energy), we can get a good idea of the energetics involved in such a process.

In practice it is easier to measure such things empirically (heat of melting, heat of dissolution etc.) than to figure out why they are that way!



These phenomena also help explain the “octet rule”.

Metals, for instance, tend to stop losing electrons once they attain a noble gas configuration because energy would be expended that cannot be overcome by lattice energies.


Electron Configuration of Ions of the Representative Elements

Na 1s22s22p63s1 = [Ne]3s1

Na+ 1s22s22p63s1 = [Ne]

Cl 1s22s22p63s23p5 = [Ne] 3s23p5

Cl_

1s22s22p63s23p6 = [Ne] 3s23p6 = [Ar]


Electron Configuration of Ions of the Representative Elements

In polyatomic ions, two or more atoms are bound together by predominantly covalent bonds and the whole acts as a charged species when the ion forms an ionic compound with an ion of opposite charge.

E.g. NH4+ and NH4Cl

CO32- and Na2CO3


Electron Configuration of Ions of the Representative Elements In forming ions, transition metals lose the valence-shell

s electrons first, then as many d electrons as are required to reach the charge of the ion.

For example, Fe has the electron configuration [Ar]3d64s2. To form the Fe2+ ion, the two 4s electrons are lost giving the [Ar]3d6 electron configuration.

To form the Fe3+ ion, the two 4s electrons and one d electron are lost giving the [Ar]3d5 electron configuration.


Covalent Bonding

In these bonds, atoms share electrons.

There are several electrostatic interactions in these bonds:

Attractions between negative electrons and positive nuclei

Repulsions between -electrons Repulsions between +nuclei

Figure 7.4


Lewis Structures

Lewis structures are representations of molecules showing all electrons - bonding and nonbonding.

--- denotes 1 pair of electrons.


Lewis Structures: Multiple Bonds

When two electron pairs are shared, two lines are drawn, i.e. double bond.

O :: C :: O or O = C = O

When three electron pairs are shared, three lines are drawn, i.e. triple bond.

:N ::: N: or :N ≡ N:


Bond Polarity and Electronegativity

Nonpolar covalent bond the electrons are shared equally between

the two atoms, e.g. in Cl2 and N2

Polar covalent bond one of the atoms exerts a greater attraction

for the bonding electrons than the other, e.g. in H2O


Bond Polarity and Electronegativity

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

On the periodic chart, electronegativity increases as you go: from left to right across

a row from the bottom to the

top of a columnFigure 7.5


Dipole Moments When two atoms share electrons unequally, a

bond dipole results.

The dipole moment, , produced by two equal but opposite charges separated by a distance, r, is calculated:

= Qr

It is measured in debyes (D).

Figure 7.7


Polar Covalent Bonds

Although atoms often form compounds by sharing electrons, the electrons are not always shared equally.

Fluorine pulls harder on the electrons it shares with hydrogen than hydrogen does.

Therefore, the fluorine end of the molecule has more electron density than the hydrogen end.

Figure 7.6


Polar Covalent Bonds

The greater the difference in electronegativity, the more polar is the bond.

Figure 7.8

Table 7.3


Explaining Polar Molecules: HCl

O.N.:Cl 7-8 = -1H 1-0 = +1


Drawing Lewis Structures

1. Find the sum of valence electrons of all atoms in the polyatomic ion or molecule.

If it is an anion, add one electron for each negative charge.

If it is a cation, subtract one electron for each positive charge.

PCl3

5 + (3 x 7) = 26



2. Write the symbols for the atoms to show which atoms are attached to which, and connect them with a single bond, a dash (representing two electrons).



3. Complete the octets around all the atoms bonded to the central atom.



4. Place any leftover electrons on the central atom.

Remember we have 26 electrons



5. If there are not enough electrons to give the central atom an octet, try multiple bonds.


Drawing Lewis Structures - Formal Charge Assign formal charges:

For each atom, count the electrons in lone pairs and half the electrons it shares with other atoms.

Subtract that from the number of valence electrons for that atom: The difference is its formal charge.



The best Lewis structure is the one: with the fewest charges that puts a negative charge on the most

electronegative atom.


Resonance Structures

This is the Lewis structure we would draw for ozone, O3.



This is at odds with the true, observed structure of ozone, in which: both O—O bonds are

the same length, and both outer oxygens have

a charge of ½.

Figure 7.10



One Lewis structure cannot accurately depict a molecule such as ozone.

We use multiple structures called resonance structures, to describe the molecule.



Just as green is a synthesis of blue and yellow, ozone is a synthesis of these two resonance structures.

Figure 7.11


Resonance Structures In truth, the electrons that form the second C—O

bond in the double bonds below do not always sit between that C and that O, but rather can move among the two oxygens and the carbon.

They are not localised, but rather are delocalised.


Exceptions to the Octet Rule

There are three types of ions or molecules that do not follow the octet rule:1. Molecules and polyatomic ions containing

an odd number of electrons.

2. Molecules and polyatomic ions in which an atom has fewer than an octet of valence electrons.

3. Molecules and polyatomic ions in which an atom has more than an octet of valence electrons.


Exceptions to the Octet RuleOdd Number of Electrons

Though relatively rare and usually quite unstable and reactive, there are ions and molecules with an odd number of electrons, e.g. NO contains 5 + 6 = 11 valence electrons.

N = O...

..

.. N = O....

...and


Exceptions to the Octet RuleLess than an Octet of Valence Electrons Consider BF3

Giving boron a filled octet places a negative charge on the boron and a positive charge on fluorine.

This would not be an accurate picture of the distribution of electrons in BF3.


Exceptions to the Octet RuleLess than an Octet of Valence ElectronsTherefore, structures that put a double bond between boron and fluorine are much less important than the one that leaves boron with only 6 valence electrons.


Exceptions to the Octet RuleLess than an Octet of Valence Electrons

If filling the octet of the central atom results in a negative charge on the central atom and a positive charge on the more electronegative outer atom, don’t fill the octet of the central atom.


Exceptions to the Octet RuleMore than an Octet of Valence Electrons

The only way PCl5 can exist is if phosphorus has 10 electrons around it.

It is allowed to expand the octet of atoms on the 3rd row or below: Presumably d orbitals in these

atoms participate in bonding.


Exceptions to the Octet RuleMore than an Octet of Valence ElectronsEven though we can draw a Lewis structure for the phosphate ion that has only eight electrons around the central phosphorus, the better structure puts a double bond between the phosphorus and one of the oxygens.


Exceptions to the Octet RuleMore than an Octet of Valence Electrons This eliminates the charge on the phosphorus

and the charge on one of the oxygens, i.e. when the central atom is on the 3rd row or below and expanding its octet eliminates some formal charges, do so.


Strengths of Covalent Bonds Most simply, the strength of a bond is measured by

determining how much energy is required to break the bond.

This is the bond enthalpy.

The bond enthalpy for a Cl—Cl bond,

D(Cl—Cl), is measured to be 242 kJ/mol.


Average Bond Enthalpies

This table lists the average bond enthalpies for many different types of bonds.

Average bond enthalpies are positive, because bond breaking is an endothermic process.NOTE: These are average bond enthalpies, not

absolute bond enthalpies; the C—H bonds in methane, CH4, will be a bit different than the C—H bond in chloroform, CHCl3.

Table 7.4


Bond Enthalpies and Enthalpies of Reaction

Yet another way to estimate H for a reaction is to compare the bond enthalpies of bonds broken to the bond enthalpies of the new bonds formed, i.e.

Hrxn = (bond enthalpies of bonds broken) - (bond enthalpies of bonds

formed)



In this example, one

C—H bond and one

Cl—Cl bond are broken

and one C—Cl and one

H—Cl bond are formed.

Figure 7.12

CH4(g) + Cl2(g) CH3Cl(g) + HCl(g)



(bond enthalpies of bonds broken) - (bond enthalpies of bonds formed)

Hrxn = [D(C—H) + D(Cl—Cl)] [D(C—Cl) + D(H—Cl)]= [(413 kJ) + (242 kJ)] [(328 kJ) + (431 kJ)]= (655 kJ) (759 kJ)= 104 kJ


Bond Enthalpy and Bond Length

We can also measure an average bond length for different bond types.

As the number of bonds between two atoms increases, the bond length decreases.

Table 7.5


Nitroglycerin(1,2,3-Trinitroxypropane)

+Diatomaceous Earth

+Sodium Carbonate=

Dynamite

PowerPoint to accompany Chapter 7 Basic Concepts of Chemical Bonding.

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Transcript of PowerPoint to accompany Chapter 7 Basic Concepts of Chemical Bonding.