Chapter 9 Ionic and Covalent Bonding -...

Chapter 9Ionic and

Covalent Bonding

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A chemical bond is a strong attractive force that exists between certain atoms in a substance.

There are three types of chemical bonds: Ionic bondsCovalent bondsMetallic bonds


An ionic bond is a chemical bond formed by the electrostatic attraction between positive and negative ions.


An ionic bond forms when one or more electrons are transferred from the valence shell of one atom to the valence shell of another atom.

Na ([Ne]3s1) + Cl ([Ne]3s23p5) �Na+ ([Ne]) + Cl- ([Ne]3s23p6)

The atom that transferred the electron(s) becomes a cation.The atom that gained the electron(s) becomes an anion.


A Lewis electron-dot symbol is a notation in which the electrons in the valence shell of an atom or ion are represented by dots placed around the chemical symbol of the element.

Note: Dots are placed one to a side, until all four sides are occupied.


Table 9.1 illustrates the Lewis electron-dot symbols for second- and third-period atoms.

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Represent the transfer of electrons in forming calcium oxide, CaO, from atoms.

+ Ca2+Ca O + O2-][


Polyatomic ions are atoms held together by covalent bonds as a group and that, as a group, have gained or lost one or more electron.


A covalent bond is a chemical bond formed by sharing a pair of electrons.


A formula using dots to represent valence electrons is called a Lewis electron-dot formula.

An electron pair is represented by two dots.

A electron pair that is between two atoms is a bonding pair. It can also be represented by one line for each bonding pair.

Electron pairs that are not bonding are nonbonding, or lone pair electrons.


A coordinate covalent bond is formed when both electrons of the bond are donated by one atom.

The two electrons forming the bond with the hydrogen on the left were both donated by the nitrogen. Once shared, they are indistinguishable from the other N—H bonds.

H+ + N HH

HNHH

HH

+


In forming covalent bonds, atoms tend toward having a full eight electrons in their valence shell. This tendency is called the octet rule.

Hydrogen is an exception to the octet rule: it has two electrons in its valence shell (a duet).


A single bond is a covalent bond in which one pair of electrons is shared by two atoms.

A double bond is a covalent bond in which two pairs of electrons are shared by two atoms.

A triple bond is a covalent bond in which three pairs of electrons are shared by two atoms.

Double bonds form primarily with C, N, and O. Triple bonds form primarily with C and N.


A polar covalent bond (or polar bond) is a covalent bond in which the bonding electrons spend more time near one atom than near the other atom.

Electronegativity, X, is a measure of the ability of an atom in a molecule to draw bonding electrons to itself. Electronegativity is related to ionization energy and electron affinity.


Electronegativity increases from left to right and from bottom to top in the periodic table. F, O, N, and Cl have the highest electronegativity values.


The difference in electronegativity between the two atoms in a bond is a rough measure of bond polarity.

When the difference is very large, an ionic bond forms. When the difference is large, the bond is polar. When the difference is small, the bond is nonpolar.

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Using electronegativities, arrange the following bonds in order by increasing polarity: C—N, Na—F, O—H.

For Na—F, the difference is 4.0 (F) – 0.9 (Na) = 3.1.

For C—N, the difference is 3.0 (N) – 2.5 (C) = 0.5.

For O—H, the difference is 3.5 (O) – 2.1 (H) = 1.4.

C—N <Bond polarities: O—H < Na—F


Writing Lewis Electron-Dot Formulas1. Calculate the number of valence electrons.2. Write the skeleton structure of the molecule or

ion.3. Distribute electrons to the atoms surrounding

the central atom or atoms to satisfy the octet rule.

4. Distribute the remaining electrons as pairs to the central atom or atoms.

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Write the electron dot formulas for the following:

a. OF2

b. NF3

c. NH2OH, hydroxylamine


Count the valence electrons in OF2:

O 1(6)

F 2(7)

20 valence electrons

O is the central atom (it is less electronegative). Now, we distribute the remaining 16 electrons, beginning with the outer atoms. The last four electrons go on O.

OF F


Count the valence electrons in NF3:

N 1(5)

F 3(7)


N is the central atom (it is less electronegative). Now, we distribute the remaining 20 electrons, beginning with the outer atoms. The last two electrons go on N.

F FN

F


Count the electrons in NH2OH:

N 1(5)

H 3(1)

O 1(6)


N is the central atom. Now, we distribute the remaining six electrons, beginning with the outer atoms. The last two electrons go on N.

H

NH HO

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Write electron-dot formulas for the following: a. CO2

b. HCN


Count the electrons in CO2:

C 1(4)

O 2(6)


C is the central atom. Now, we distribute the remaining 12 electrons, beginning with the outer atoms.

Carbon does not have an octet, so two of the lone pairs shift to become a bonding pair, forming double bonds.

OO CO O


Count the electrons in HCN:H 1(1)C 1(4)N 1(5)

10 valence electrons.

C is the central atom. The remaining electrons go on N.

Carbon does not have an octet, so two of the lone pairs shift to become a bonding pair, forming a triple bond.

CH NN

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Phosphorus pentachloride exists in solid state as the ionic compound [PCl4]+[PCl6]-; it exists in the gas phase as the PCl5 molecule. Write the Lewis formula of the PCl4+ ion.


Count the valence electrons in PCl4+:

P 1(5)

Cl 4(7)

-1

32

P is the central atom. The remaining 24 nonbonding electrons are placed on Cl atoms. Add square brackets with the charge around the ion.

Cl

Cl

ClCl P

+


Delocalized bonding is a type of bonding in which a bonding pair of electrons is spread over a number of atoms rather than being localized between two atoms.


A single electron-dot diagram cannot properly describe delocalized bonding. Using the resonance description, the electron structure of a molecule or ion having delocalized bonding is given by writing all possible electron-dot formulas. They are connected with a double-headed arrow.

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Draw the resonance formulas of the acetate ion, CH3COO-.


CH3COO-

Valence electrons: 2(4) + 3(1) + 2(6) + 1 = 24

C is the central atom.

A double bond is needed between C—O.

There are two equivalent places for it, so two resonance structures are required.

C

H

H

H CO

O

-

C

H

H

H CO

O

-


Some molecules have electron-dot structures that do not satisfy the octet rule. Some have an odd number of electrons, such as NO. Other molecules either have too few or too many electrons around the central atom.


Elements that have too few electrons are in Groups IIA and IIIA. Be, B, and Al exhibit too few electrons around the central atom.

Al

FF

F

B

FF

F

BeF F


There are many more examples of central atoms with more than an octet. Because elements of the third period and beyond have a d subshell, they can expand their valence electron configurations. S, P, Cl (as a central atom), and other elements in period 3 are examples of atoms in this situation.

Elements in the second period, having only s and p subshells, are unable to do this.

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Give the Lewis formula of the IF5molecule.


Count the valence electrons in IF5:

I 1(7)

F 5(7)


I is the central atom. Thirty-two electrons remain; they first complete F octets. The remaining electrons go on I.

F

F

F

I

F

F


The formal charge on an atom in the Lewis formula is the hypothetical charge you obtain by assuming that bonding electrons are equally shared between bonded atoms and that the electrons of each lone pair belong completely to one atom.


Formal charge = valence electrons on free atom

– ½ (number of electrons in bonds)– (number of lone-pair electrons)

The sum of the formal charges on the atoms equals the charge on the formula.


Formal charges can help to determine the most likely electron-dot formula using three rules:1. Whenever you can write several Lewis formulas

for a molecule, choose the one having the lowest magnitudes of formal charges.

2. When two proposed Lewis formulas have the same magnitudes of formal charges, choose the one having the negative formal charge on the more electronegative atom.

3. When possible, choose Lewis formulas that do not have like charges on adjacent atoms.

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Compare the formal charges for the following electron-dot formulas of CO2.

CO O

For the left structure: For the right structure:

O: 6 – 2 – 4 = 0C: 4 – 4 – 0 = 0

The left structure is better.

CO O

Formal charge = group number – (number of bond pairs) – (number of nonbonding electrons)

C: 4 – 4 – 0 = 0O: 6 – 1 – 6 = –1O: 6 – 3 – 2 = +1


Bond length (or bond distance) is the distance between nuclei in a bond.

Bond order is, defined in terms of the Lewis formula, the number of pairs of electrons in a bond.

Bond length decreases as bond order increases.

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Consider the propylene molecule:

C C C

H

H

H

H

H

H

134 pm 150 pm

The shorter bond is the double bond; the longer bond is the single bond.

One of the carbon–carbon bonds has a length of 150 pm; the other 134 pm. Identify each bond with a bond length.

Chapter 9 Ionic and Covalent Bonding -...

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