Chapter 9Ionic and
Covalent Bonding
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Contents and Concepts
Ionic Bonds
Molten salts and aqueous solutions of salts are electrically conducting. This conductivity results from the motion of ions in the liquids. It suggests the possibility that ions exist in certain solids, held together by the attraction of ions of opposite charges.
1. Describing Ionic Bonds
2. Electron Configurations of Ions
3. Ionic Radii
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Covalent BondsNot all bonds can be ionic. Hydrogen, H2, is a clear example in which there is a strong bond between two like atoms. The bonding in the hydrogen molecule is covalent. A covalent bond forms between atoms by the sharing of a pair of electrons.
4. Describing Covalent Bonds 5. Polar Covalent Bonds; Electronegativity 6. Writing Lewis Electron-Dot Formulas 7. Delocalized Bonding: Resonance 8. Exceptions to the Octet Rule 9. Formal Charge and Lewis Formulas10. Bond Length and Bond Order11. Bond Energy
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Learning Objectives
Ionic Bonds
1. Describing Ionic Bonds
a. Define ionic bond.
b. Explain the Lewis electron-dot symbol of an atom.
c. Describe the energetics of ionic bonding.
d. Define lattice energy.
e. Describe the Born–Haber cycle to obtain a lattice energy from thermodynamic data.
f. Describe some general properties of ionic substances.
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2. Electron Configurations of Ions
a. State the three categories of monatomic ions of main-group elements.
b. Write the electron configuration and Lewis symbol for a main-group ion.
c. Note the polyatomic ions given earlier in Table 9.2.
d. Note the formation of +2 and +3 transition-metal ions.
e. Write the electron configurations of transition-metal ions.
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3. Ionic Radii
a. Define ionic radius.
b. Define isoelectronic ions.
c. Use periodic trends to obtain relative ionic radii.
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Covalent Bonds
4. Describing Covalent Bonds
a. Describe the formation of a covalent bond between two atoms.
b. Define Lewis electron-dot formula.
c. Define bonding pair and lone (nonbonding) pair of electrons.
d. Define coordinate covalent bond.
e. State the octet rule.
f. Define single, double, and triple bond.
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5. Polar Covalent Bonds; Electronegativity
a. Define polar covalent bond.
b. Define electronegativity.
c. State the general periodic trends in electronegativity.
d. Use electronegativity to obtain relative bond polarity.
6. Writing Lewis Electron-Dot Formulas
a. Write Lewis formulas having single bonds.
b. Write Lewis formulas having multiple bonds.
c. Write Lewis formulas for ionic species.
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7. Delocalized Bonding: Resonance
a. Define localized bonding.
b. Define resonance description.
c. Write resonance forms.
8. Exceptions to the Octet Rule
a. Write Lewis formulas (exceptions to the octet rule).
b. Note exceptions to the octet rule in Group IIA and Group IIIA.
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9. Formal Charge and Lewis Formulas
a. Define formal charge.
b. State the rules for obtaining the formal charge.
c. State two rules useful in writing Lewis formulas.
d. Use formal charges to determine the best Lewis formula.
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10. Bond Length and Bond Order
a. Define bond length (bond distance).
b. Define covalent radii.
c. Define bond order.
d. Explain how bond order and bond length are related.
11. Bond Energy
a. Define bond energy.
b. Estimate H from bond energies.
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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 bonds
Covalent bonds
Metallic bonds
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An ionic bond is a chemical bond formed by the electrostatic attraction between positive and negative ions.
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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.
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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.
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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 + O 2-][
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Let’s look next at the energy involved in forming ionic compounds.
The energy to remove an electron is the ionization energy.
The energy to add an electron is the electron affinity.
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The combination of ionization energy and electron affinity is still endothermic; the process requires energy.
However, when the two ions bond, more than enough energy is released, making the overall process exothermic.
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The lattice energy is the change in energy that occurs when an ionic solid is separated into gas-phase ions.
It is very difficult to measure lattice energy directly. It can be found, however, by using the energy changes for steps that give the same result.
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For example, to find the lattice energy for NaCl, we can use the following steps.
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The process of finding the lattice energy indirectly from other thermochemical reactions is called the Born–Haber cycle.
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Ionic substances are typically high-melting solids. There are two factors that affect the strength of the ionic bond. They are given by Coulomb’s law:
The higher the ionic charge, the stronger the force; the smaller the ion, the stronger the force.
221
r
QkQF
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Based on this relationship, we can predict the relative melting points of NaCl and MgO.
The charge on the ions of MgO is double the charge on the ions of NaCl. Because the charge is double, the force will be four times stronger.
The size of Na+ is larger than that of Mg2+; the size of Cl- is larger than that of O2-. Because the distance between Mg2+ and O2- is smaller than the distance between Na+ and Cl-, the force between Mg2+ and O2- will be greater.
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Based on the higher charge and the smaller distance for MgO, its melting point of MgO should be significantly higher than the melting point of NaCl.
The actual melting point of NaCl is 801°C; that for MgO is 2800°C.
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When we examine the electron configuration of main-group ions, we find that each element gains or loses electrons to attain a noble-gas configuration.
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Give the electron configuration and the Lewis symbol for the chloride ion, Cl-.
Cl -][
For chlorine, Cl, Z = 17, so the Cl- ion has 18 electrons. The electron configuration for Cl- is
1s2 2s2 2p6 3s2 3p6
The Lewis symbol for Cl- is
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Group IIIA to VA metals often exhibit two different ionic charges: one that is equal to the group number and one that is 2 less than the group number.
The higher charge is due to the loss of both the s subshell electrons and the p subshells electron(s). The lower charge is due to the loss of only the p subshell electron(s).
For example, in Group IVA, tin and lead each form both +4 and +2 ions. In Group VA, bismuth forms +5 and +3 ions.
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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.
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Transition metals form several ions.
The atoms generally lose the ns electrons before losing the (n – 1)d electrons.
As a result, one of the ions transition metals generally form is the +2 ion.
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Give the electron configurations of Mn and Mn2+.
Manganese, Z = 25, has 25 electrons;. Its electron configuration is
1s2 2s2 2p6 3s2 3p6 4s23d5
Mn2+ has 23 electrons. When ionized, Mn loses the 4s electrons first; the electron configuration for Mn2+ is
1s2 2s2 2p6 3s23p6 3d5
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a. Fe2+: [Ar]3d44s2
No. The 4s2 electrons would be lost before the 3d electrons.
b. N2-: [He]2s22p5
No. Nitrogen will gain three electrons to fill the shell, forming N3-.
c. Zn2+: [Ar]3d10
Yes! d. Na2+: [He]2s22p5
No. Sodium will lose only its one valence electron, forming Na+.
e. Ca2+: [Ne]3s23p6
Yes!
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Ionic radius is a measure of the size of the spherical region around the nucleus of an ion within which the electrons are most likely to be found. While ionic radius, like atomic radius, can be somewhat arbitrary, it can be measured in ionic compounds.
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A cation is always smaller than its neutral atom.
An anion is always larger than its neutral atom.
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The term isoelectronic refers to different species having the same number and configuration of electrons.
For example, Ne, Na+, and F- are isoelectronic.
Ionic radius for an isoelectronic series decreases with increasing atomic number.
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Using the periodic table only, arrange the following ions in order of increasing ionic radius: Br-, Se2-, Sr2+.
35Br
34Se
38Sr
These ions are isoelectronic, so their size decreases with increasing atomic number:
Sr2+ < Br- < Se2-
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A covalent bond is a chemical bond formed by sharing a pair of electrons.
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To consider how a covalent bond forms, we can monitor the energy of two isolated hydrogen atoms as they move closer together.
The energy decreases—first gradually, and then more steeply—to a minimum. As the atoms continue to move closer, it increases dramatically.
The distance between the atoms when energy is at a minimum is called the bond length.
This is illustrated on the following graph, from right to left.
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As the hydrogen atoms move closer together, the electron of each atom is attracted to both its own nucleus and the nucleus of the second atom. The electron probability distribution illustrates this relationship.
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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.
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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+
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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).
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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.
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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.
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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.
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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
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Writing Lewis Electron-Dot Formulas
1. 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
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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
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Count the valence electrons in NF3:
N 1(5)
F 3(7)
26 valence electrons
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
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Count the electrons in NH2OH:
N 1(5)
H 3(1)
O 1(6)
14 valence electrons
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
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Count the electrons in CO2:
C 1(4)
O 2(6)
16 valence electrons
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
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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.
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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
+
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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.
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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-.
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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
-
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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.
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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
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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 IF5 molecule.
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Count the valence electrons in IF5:
I 1(7)
F 5(7)
42 valence electrons
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
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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.
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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.
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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
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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.
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Bond energy is the average enthalpy change for breaking the A—B bond in a molecule in the gas phase. Bond energy is a measure of bond strength: the larger the bond energy, the stronger the bond.
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Bond energies can be used to estimate the enthalpy change, H, for a reaction. To do so, we imagine the reaction in two steps: breaking bonds and forming new bonds.
H = sum of the bond energies for bonds broken
– sum of the bond energies for bonds formed
When H is negative, heat is released. When H is positive, heat is absorbed.
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Estimate the enthalpy change for the following reaction, using bond energies:
Bonds Broken:
1 C=C 602 kJ
1 Cl—Cl 240 kJ
Absorbed 842 kJ
C C
H
H
H
H
+ Cl2 C C
H
H
Cl
H
Cl
H
Bonds Formed:
1 C—C 346 kJ
2 C—Cl 654 kJ
Released1000 kJH = 842 kJ – 1000 kJH = –158 kJ
