Download pdf - Chapter 8 Chemical Bonds - Angelo State University · PDF fileChapter 8: Chemical Bonds (+ VSEPR) 5 Covalent Bonds Share and Share Alike 6 Covalent Bonds and Molecules • Covalent

Chapter 8: Chemical Bonds (+ VSEPR)

Mr. Kevin A. Boudreaux Angelo State University

CHEM 1411 General Chemistry Chemistry: The Science in Context (Gilbert, 4th ed, 2015)

www.angelo.edu/faculty/kboudrea

Chapter Objectives:

• Understand the principal types of chemical bonds.

• Understand the properties of ionic and molecular compounds.

• Draw Lewis dot structures for molecular compounds, including resonance structures.

• (Chapter 9.1-9.3) Predict the shape of a molecule using VSEPR theory, and determine whether the molecule is polar.

Chapter 8 Chemical Bonds

(+VSEPR from Chapter 9)

2

Introduction

• Most of the substances that we encounter in daily life are not elemental substances, but compounds (and frequently, complex mixtures of compounds).

• Why do elements form compounds in the first place? Bonding lowers the potential energy between the charged particles that compose atoms.

• There are three major ways of modeling bonds between atoms, with varying degrees of complexity:

– Lewis theory (Lewis dot structures)

– Valence Bond theory

– Molecular Orbital theory


3

Types of Chemical Bonds

• There are three types of bonds between elements:

– Ionic bonds result from a transfer of electrons from one species (usually a metal) to another (usually a nonmetal or polyatomic ion).

– Covalent bonds result from a sharing of electrons by two or more atoms (usually nonmetals).

– Metallic bonds result from a “pooling” of valence electrons by two or more metals into a delocalized electron sea.

4

Lewis Dot Structures for Elements

• In a Lewis dot structure [G. N. Lewis] for an element, the valence electrons are written as dots surrounding the symbol for the element.

– Place one dot on each side first; the remaining dots are paired with one of the first set of dots.

– A maximum of two dots are placed on each side.

– The unpaired dots indicate where covalent bonds can form, or where electrons can be gained or lost.

– Atoms tend to form bonds to satisfy the octet rule, having eight electrons in their valence shells.

Li Be B C N O F Ne


5

Covalent Bonds

Share and Share Alike

6

Covalent Bonds and Molecules

• Covalent bonds form when two or more nonmetals share their electrons. The electrons are at their lowest potential energy when they are between the two nuclei that are being joined.

• Each atom in the bond “holds on” to the shared electrons, and the atoms are thus physically tied together.


7

The Formation of Diatomic Hydrogen

• As two isolated H atoms move closer together, the two positively-charged nuclei repel each other, and the two negatively-charged electrons repel each other, but each nucleus attracts both electrons.

• At some point, the attractions between the nuclei and the electrons are balanced against the repulsions between the nuclei and between the electrons.

– The shared electrons bind the two nuclei into an H2 molecule.

– The shared electrons act like they belong to both atoms in the bond.

Covalent Bonding and Potential Energy

• When two isolated H atoms approach each other, the potential energy is lowered — energy is released when an H2 molecule forms.

• Pushing the nuclei any closer causes the potential energy to rise.

8 Figure 8.1


Covalent Bonding and Potential Energy

• The optimum distance between nuclei where the attractive forces are maximized and the repulsive forces are minimized is called the bond length. (For H2, the bond length is 74 pm.)

• In H2, the highest probability of finding electrons is in the space between the H nuclei.

– The increased attractive forces in this area help to lower the potential energy of H2 relative to isolated H atoms.

9

Lewis Structures and the Octet Rule

• The sharing of electrons by nonmetals to form molecules is modeled by the use of Lewis structures, in which sticks (—) represent pairs of shared electrons, and dots (:) represent unshared electrons (lone pairs).

– Atoms share electrons in such a way as to satisfy the octet rule, which gives each atom a total of eight valence electrons.

– Hydrogen is an “exception” to the octet rule, since its 1s orbital can only hold two electrons.

• Once the Lewis structure has been drawn, the 3D shape of the molecule and the polarity of the molecule can be predicted using the VSEPR model.

10


11

Single Covalent Bonds

• The shared pairs of electrons are bonding pairs.

• The unshared pairs of electrons are lone pairs or nonbonding pairs.

Na+ Cl- ionic bond

H—H covalent bond

+ H HorH H H H

+ H ForH F H F

+ F ForF FF F

+ H OorO OH H+ H H H

12

Double and Triple Covalent Bonds

• Some atoms can satisfy the octet rule by sharing two pairs of electrons to form a double bond:

– Double bonds are shorter and stronger than single

bonds.

• Some atoms can share three pairs of electrons to form a triple bond:

– Triple bonds are even shorter and stronger than

double bonds.

O+ O OO OO

+ NN N N


13

Lewis Structures

Drawing the Line

14

How To Write Lewis Structures

1. Determine the total number of valence electrons in the molecule or ion.

• Add one electron for each unit of negative charge.

• Subtract one electron for each positive charge.

2. Write the correct skeletal structure.

• For molecules of the formula ABn, place the least electronegative in the center, and the remaining atoms on the outside. Draw single bonds between the outer atoms and the central atom.

• H is NEVER a central atom, since it can only form one bond.

• • • — — — • • •


15

How To Write Lewis Structures

3. Distribute the remaining valence electrons as lone pairs on the outer atoms first, making sure to satisfy the octet rule.

• Once all of the outer atoms have 8 electrons (or 2 for H), place any remaining electrons on the central atom.

4. Compare the number of valence electrons in the Lewis structure to the number in Step 1 to make sure you haven’t miscounted electrons.

5. Complete the octet on the central atom.

• This is done by sharing lone pairs from the outer atoms with the central atom.

• The formal charge can be used as a guideline for placing double bonds.

16

Formal Charges

• Formal charge is the difference between the number of valence electrons on the free atom and the number of valence electrons assigned to the atom in the molecule.

– The sum of the formal charges must equal the charge on the species.

– Smaller formal charges are better (more stable) than larger ones.

– Like charges on adjacent atoms are not desirable.

– When a formal charge cannot be avoided, negative formal charges should reside on more electronegative atoms.

Formal charge = valence e- – (½ bonding e-) – (lone pair e-)


17

Examples: Lewis Structures

1. Write Lewis structures for the following molecules.

a. CH4

b. NH3

c. NH4+

d. CCl4

e. H2O2

18



f. CCl2F2

g. H2S

h. C2H6

i. C2H4

j. N2O (atoms connected in the order N—N—O)


19



k. CO2

l. CaCl2

m COCl2

n. acetamide, C2H5NO

CC

N

O

H

H

H

HH

20



o. O3

Neither of the two O3 structures is correct by itself.


21

Resonance Structures

• When there is more than one valid Lewis structure for a molecule, the actual electronic structure is an average of the different possibilities, called a resonance hybrid.

• Resonance forms differ only in the placement of the valence electrons, not the positions of the atoms!

• Ozone does not have a “real” double bond and a “real” single bond; the actual bond lengths in ozone are identical, and are between those of an O—O bond and an O=O bond — a “one-and-a-half” bond.

When One Lewis Structure Ain’t Enough

22

Resonance Structures • This molecule is shown more correctly with two

Lewis structures, called resonance structures, with a two-headed arrow (T) between them:

• DO NOT USE THE STRAIGHT, DOUBLE-HEADED ARROW (T) FOR ANYTHING ELSE!

• Resonance structures are not real bonding depictions: O3 does not “change back and forth” between the two structures; the actual molecule is a hybrid of the two forms depicted.

O

OO O

OO

Bond order = 1 1/2


23

Delocalized Electrons and Charges

• In these resonance structures, one of the electron pairs (and hence the negative charge) is “spread out” or delocalized over the whole molecule.

– In contrast, the lone pairs on the oxygen in water are localized — i.e., they’re stuck in one place.

• Resonance delocalization stabilizes a molecule by spreading out charges, and often occurs when lone pairs (or positive charges) are located next to double bonds.

• (Resonance plays a big role in our understanding of structure and reactivity in organic chemistry.)

• A more accurate depiction of electron distribution is found in molecular orbital (MO) theory.

24

Resonance Structures in Benzene

• Another molecule in which resonance is important is benzene, C6H6, which has two resonance structures with alternating single and double bonds. The actual bond order of the carbon-carbon bonds in benzene is 1.5.

C

CC

C

CC

C

CC

C

CC

H

H

H

H

H

H H

H

H

H

H

H


25

A Resonance Analogy

• A mule is not sometimes a horse and sometimes a donkey; it’s always one thing (a mule), just like purple is not sometimes red and sometimes blue.

• A real person can be described as having characteristics of two or more fictional characters. The fictional characters don’t exist, but the real person does.

26

Examples: Lewis Structures With Resonance

2. Write Lewis structures for the following molecules, including resonance structures.

a. OCN-

b. CO32-

c. NO3-

d. CHO2-


27

Exceptions to the Octet Rule, Part 1

• Electron deficient species, such as beryllium (Be) and boron (B), can have fewer than eight electrons around them, but have zero formal charge.

• Free radicals (or just radicals, or odd-electron molecules) contain an odd number of valence electrons. Radicals always have an unpaired electron, and are paramagnetic. These species are usually unstable, and are extremely reactive.

• Expanded octets are found on atoms that have more than eight electrons around them. Nonmetals from period 3 or higher, such as sulfur and phosphorus, can get around the octet rule by shoving “extra” electrons into empty d orbitals. [more later]

28

Examples: Exceptions to the Octet Rule

3. Write Lewis structures for the following molecules, including resonance structures, if necessary. (BEWARE! The formula alone doesn’t tell you about lone pairs on the central atom!)

a. NO

b. NO2

c. BeCl2

d. BF3

e. SF6


29

Polar Covalent Bonds

30

Electronegativity

• In reality, fully ionic and covalent bonds represent the extremes of a spectrum of bonding types.

• Electronegativity is the ability of an atom in a molecule to attract shared electrons to itself. [Linus Pauling, 1939; Nobel Prize 1954, 1963]

• Electronegativity is a periodic property, and increases from bottom to top within a group and from left to right across a period (inversely related to atomic radius).


Electronegativity Values (Pauling Scale)

31 Figure 8.6

Electronegativity Values (The Pauling Scale)

• The atom in a covalent bond with the larger electro-negativity value has a partial negative charge (d-), because it pulls the shared electrons closer.

• The atom in a covalent bond with the smaller electro-negativity value has a partial positive charge (d+), because the shared electrons are partially pulled away.

32

H F


Polar and Nonpolar Covalent Bonds

• A bond in which the electronegativity difference between the two atoms is large (> 2.0) is an ionic bond, where the electron is transferred from one atom to the other:

• A bond in which there is a very small electro-negativity difference (< 0.4) between the bonded atoms is a nonpolar covalent bond:

33

Polar and Nonpolar Covalent Bonds

• A bond in which there is an intermediate electro-negativity difference between the bonded atoms (between 0.4 and 2.0) is a polar covalent bond:

– the electron cloud between the atoms is polarized; that is, it “leans” towards one side of the bond.

– The electrons are still shared, but unequally.

34


Electronegativity and Bond Polarity

• A polar covalent bond can be thought of as having partial covalent character and partial ionic character: the greater the DEN of the atoms, the larger the ionic character of the bond.

35

36



37


• In the period 3 chlorides below, as the difference in electronegativity (DEN) decreases, the bond becomes more covalent; we move from tightly-bound ionic solids (NaCl) to more weakly bound covalent liquids (SiCl4) to even more weakly bound gases (Cl2).

38

Dipole Moment

• Bond polarity is expressed numerically as a dipole moment, , which occurs when there is a separation between a positive and negative charge.

– The unit of dipole moment is the debye, D (1 D = 3.34×10-30 C·m)

38

• A thin stream of a polar solvent, such as water, is deflected by a static electric charge, while a nonpolar molecule, such as hexane, is not.


39

Examples: Nonpolar / Polar / Ionic Bonds

2. Determine whether the bond formed between each of the following pairs of atoms is a nonpolar covalent bond, a polar covalent bond, or an ionic bond.

a. Na and Cl

b. C and Cl

c. N and Cl

d. N and O

e. Sr and F

f. Cl and Cl

Vibrating Bonds and the Greenhouse Effect

• Because CO2 contains polar bonds, it can absorb and emit photons of infrared (IR) radiation. In some vibrational modes which change the electric fields associated with bond polarity, IR photons may be absorbed. This is why CO2 is a greenhouse gas.

• Because N2 and O2 are nonpolar molecules, there is no vibrational mode of stretching or bending which changes the electric fields associated with bond polarity; they are IR inactive.

40

Figure 8.10


41

VSEPR Theory

Getting in Shape

42

You’ve Just Crossed Over into . . .

• The Lewis structures we’ve seen so far only tell us how the atoms are connected, not what their arrangement in space is.

• The approximate molecular structure, the three-dimensional arrangement of the atoms in a molecule, can be predicted from the Lewis dot structure using the Valence-Shell Electron-Pair Repulsion (VSEPR) model,

• The 3D shape of a molecule is important in determining its overall polarity, chemical behavior, and is particular for biologically important molecules, which often have complex and specific 3D shapes.


The VSEPR Model

• In the VSEPR model, an atom’s bonding and nonbonding electron pairs are positioned as far apart as possible to minimize electron-pair repulsions.

– Each bond or lone pair, or unpaired electron counts as one electron group (the book uses the term steric number). Multiple bonds count as one electron group.

– There are five electron-group shapes (or electron-pair geometries) in the VSEPR model, based on the positions of all of the electron groups around an atom (bond and lone pairs).

– The molecular shape (or molecular geometry) is the three-dimensional shape of the molecule, based on the positions of the atoms only.

43

44

VSEPR Electron-Group Shapes

A Balloon Analogy for the VSEPR Model


45

Two Electron Groups — EGS = Linear

• Two electron groups: electron-group shape = linear

– 2 bonds, 0 lone pairs: molecular shape = linear, 180°

BeCl Cl

CO O

180°

linear

46

Three Electron Groups — EGS=Trigonal Planar

• Three electron groups: electron-group shape = trigonal planar

– 3 bonds, 0 lone pairs: molecular shape = trigonal planar, 120°

– 2 bonds, 1 lone pair: molecular shape = bent, < 120°

120°

trigonal planar

<120°

bent Electron clouds in lone pairs (b)

take up more room than those

in covalent bonds (a) ; this

causes the other atoms to be

more squashed together.)


Three Electron Groups — EGS=Trigonal Planar

47

BF F

F

C

O

H H

SO O

48

Four Electron Groups — EGS = Tetrahedral

• Four electron groups: electron-group shape = tetrahedral

– 4 bonds, 0 lone pairs: molecular shape = tetrahedral, 109.5°

– 3 bonds, 1 lone pair: molecular shape = trigonal pyramidal, <109.5°

– 2 bonds, 2 lone pairs: molecular shape = bent, <109.5°

<109.5°

trigonal pyramidal

<109.5°

bent

109.5°

tetrahedral

3D Models/Sulfur Dioxide.MOL

3D Models/Formaldehyde.MOL



49

CH H

H

H

50


H N H

H

OH H

3D Models/Methane.MOL

3D Models/Ammonia.MOL

3D Models/Ammonia.MOL


51

52

Molecular Shape and Polarity

• For diatomic molecule, the polarity of the molecule depends on the polarity of the only covalent bond: if the bond is polar the molecule must be polar.


• In molecules with more than one bond, the polarity of the bonds and the overall shape determine whether the molecule is polar.

– CO2 is nonpolar, because even though it has polar bonds, they point 180º from each other.

– Water is polar because it has polar O—H bonds

which point about 104.5º away from each other.

Molecular Shape and Polarity

53

54

Determining Molecular Polarity

• To determine whether a molecule is polar:

– Draw a Lewis structure for the molecule and predict its shape using VSEPR theory.

– Determine whether the molecule contains polar bonds. If the molecule contains polar bonds, superimpose a vector ( ), pointing towards the more electronegative atom of each bond.

– Add the vectors corresponding to the polar bonds in the molecule.

• If the vectors sum to zero, the molecule is nonpolar (zero dipole moment)

• If there is a net vector, the molecule is polar (dipole moment > 0)


55

Determining Molecular Polarity

CCl4 CHCl3

C

Cl

Cl

Cl

Cl C

Cl

Cl

Cl

H

nonpolar molecule polar molecule

56

Physical Properties of Polar Molecules

• The polarity of molecules has a large effect on their chemical and physical properties:

– Polar molecules attract one another more strongly than nonpolar molecules do, and generally have higher boiling points.

– Polar molecules interact with each other more than they interact with nonpolar molecules, which is why a mixture of oil (nonpolar) and water (polar) separates into two layers.


57

Examples: Shape and Polarity

1. Write Lewis structures for the following molecules and determine their shape. State whether or not the molecules will be polar.

a. BF3

b. NH3

c. SF2

d. CF4

e. CH2F2

58

Examples: Shape and Polarity

1. Write Lewis structures for the following molecules and determine their shape. State whether or not the molecules will be polar.

f. PF3

g. HCN

h. CH2O


59

Expanded Octets: Lewis Structures,

VSEPR, and Polarity

60

Exceptions to the Octet Rule, Part 2

• Expanded octets are found on atoms that have more than eight electrons around them. Nonmetals from period 3 or higher, such as sulfur and phosphorus, can get around the octet rule by shoving “extra” electrons into empty d orbitals.

– These atoms do not always violate the octet rule (e.g., the S in H2S follows the octet rule, but the S in SF4 and SF6 has an expanded octet).

– Period 2 elements cannot have expanded octets.

– In general, nonmetals from period 3 or higher have expanded octets when they are bonded to strongly electronegative elements (F, O, Cl), or when an expanded valence shell reduces the formal charges on the atoms.


Five Electron Groups — Trigonal Bipyramidal

• Five electron groups: electron-group shape = trigonal bipyramidal

– 5 bonds, 0 lone pairs: molecular shape = trigonal bipyramidal, 120° (equatorial), 90° (axial)

– 4 bonds, 1 lone pair: molecular shape = seesaw, <120° (equatorial), <90° (axial)

– 3 bonds, 2 lone pairs: molecular shape = T-shaped, <90°

– 2 bonds, 3 lone pairs: molecular shape = linear, 180°

61


62

trigonal bipyramidal

eqeq

eq

ax

axeq = equatorialax = axial

120°

90°

<120°

<90°

seesaw

<90°

T-shaped

180°

linear


63


P

Cl

Cl Cl

Cl Cl

SF F

F

F

64


F Br F

F

XeF F

3D Models/Phosphorus Pentachloride.MOL

3D Models/Sulfur Tetrafluoride.MOL

3D Models/Chlorine Trifluoride.MOL

3D Models/Triiodide Ion.MOL


Six Electron Groups — Octahedral

• Six electron groups: electron-group shape = octahedral

– 6 bonds, 0 lone pairs: molecular shape = octahedral, 90°

– 5 bonds, 1 lone pair: molecular shape = square pyramidal, <90°

– 4 bonds, 2 lone pairs: molecular shape = square planar, 90°

65

90° <90° 90°

octahedral square pyramidal square planar

Six Electron Groups — Octahedral

66

S

F

F

F

F F

F

BrF F

F

F F

Xe

F

F

F F

3D Models/Sulfur Hexafluoride.MOL

3D Models/Bromine Pentafluoride.MOL

3D Models/Xenon Tetrafluoride.MOL


67

68


69


1. Write Lewis structures for the following molecules, including resonance structures, if necessary. Predict the shape of the molecules, and whether or not the molecules are polar (BEWARE! The formula doesn’t tell you how many lone pairs on the central atom!)

a. SO3

b. SF4

c. SF6

d. SO42-

70


1. Write Lewis structures for the following molecules, including resonance structures, if necessary. Predict the shape of the molecules, and whether or not the molecules are polar (BEWARE! The formula doesn’t tell you how many lone pairs on the central atom!)

e. POCl3

f. I3-

g. XeF2

h. XeF4


71

Shapes of Larger Molecules

• The shapes of larger molecules are a composite of the shapes of the atoms within the molecule, each of which can be predicted using VSEPR theory.

C

H

H

CH H

H

H

C

H

H

CH O

H

H

H

72

3D Models/Ethane.MOL

3D Models/Ethanol.MOL

3D Models/Ethene.MOL


73

Bond Lengths and

Bond Energies

Bond Lengths

• Bond order is the number of bonds between two atoms:

– a single bond has a bond order of 1

– a double bond has a bond order of 2

– a triple bond has a bond order of 3

– As the bond order increases, the bond length decreases.

• Molecules in which where is resonance may have fractional bond orders if the double bond is “spread out” over more than one position:

74

b.o. = 2

b.o. = 1

b.o. = 1.5

Figure 8.18


75

Bond Lengths

• Multiple bonds are shorter and stronger than their single-bond counterparts.

• When larger atoms are joined together, the bond becomes longer.

• Longer bonds are weaker than shorter bonds.

Figure 8.19

76

Average Covalent Bond Lengths and Bond Energies


77

Bond Length & Bond Energy for Multiple Bonds

Bond Bond Length

(pm) Bond Energy

(kJ/mol)

C—C 154 347

C=C 134 614

CtC 120 839

C—O 143 358

C=O 123 745

CtO 113 1072

N—N 145 163

N=N 123 418

NtN 110 946

C—N 143 305

C=N 138 891

CtN 116 891

78

Bond Energies

• The bond energy (or bond dissociation energy) is the amount of energy required to break 1 mole of the covalent bond in the gas phase.

• For the Cl—Cl bond in Cl2, the bond energy is 243 kJ/mol:

Cl2(g) 2Cl(g); DH = +243 kJ

• For the H—Cl bond in HCl, the bond energy is 431 kJ/mol

HCl(g) H(g) + Cl(g); DH = +431 kJ

• The HCl bond is stronger than the Cl—Cl bond, because it takes more energy to break the bond.


79

Bond Energies

• Bond energies are positive (endothermic) because it takes energy to break a bond.

• When a bond forms, the bond energy is the amount of energy that is released (exothermic).

• Table 8.3 contains a list of average bond energies, with values obtained by averaging the energies of that type of bond in a number of different compounds.

Bond Energies and Enthalpy Changes

• Average bond energies can be used to estimate enthalpy changes for chemical reactions:

– A reaction is exothermic when weak bonds break and strong bonds form.

– A reaction is endothermic when strong bonds break and weak bonds form.

80

DDD formed bonds s' broken bonds s'rxn HHHpositive values negative values

or

DDD formed bonds s' broken bonds s'rxn HHH


81

Bond Energies and Enthalpy Changes

Example: Use the table of average bond energies to estimate the change in enthalpy for the combustion of methane (CH4).

CH4 + 2O2 CO2 + 2H2O

Bonds Broken

4 × C—H 4(+413 kJ)

2 × O=O 2(+495 kJ)

Bonds Formed

2× C=O 2(-799 kJ)

4× H—O 4(-463 kJ)

DH = (1652 + 990) +

(-1598 + -1852) kJ

= (2642) + (-3450) kJ

= -808 kJ

82

Examples: Bond Length and Bond Strength

1. Using the periodic table, rank the bonds in each set in order of increasing bond length and increasing bond strength.

a. S—F, S—Br, S—Cl

b. C=O, C—O, CtO


83

Examples: Estimating Enthalpy Changes

2. Hydrogen gas can be made by the reaction of methane gas and steam.

CH4(g) + 2H2O(g) 4H2(g) + CO2(g)

Use the bond energies in Table 8.3 to estimate DH for this reaction.

Answer: +162 kJ

84

Properties of Molecular Compounds

• Ionic bonds are nondirectional, and hold together an entire array of ions in a crystal lattice. Covalent bonds are directional, and hold specific atoms together in a molecule.

• Molecular compounds are generally gases, liquids, or low-melting solids. The covalent bonds within molecules are very strong, but the attractive forces between the separate molecules are fairly weak.


85

Ionic Bonds

Ionic Bonds

• Ionic bonds form when one atoms transfers one or more electrons to another atom, producing ions.

• Ionic compounds are compounds that are held together by ionic bonds between positively-charged cations and negatively-charged anions.

– The ionic bond is the strong attraction between the cations and the anions.

– The cation and anion are not physically joined (i.e., they do not form a molecule).

• Ionic compounds generally result when a metal combines with a nonmetal:

– Metal + Nonmetal ionic compound

– Metal + Polyatomic ion ionic compound 86


87

Formation of Ionic Solids

• When an element which gives up electrons easily (i.e., that has a small ionization energy) comes in contact with an element that accepts an electron easily (i.e., that has a large negative electron affinity), an electron may be transferred, yielding a cation and an anion.

Na [He] 2s2 2p6 3s1

Cl [Ne] 3s2 3p5

Na+ [He] 2s2 2p6 = [Ne]

Cl- [Ne] 3s2 3p6 = [Ar]

Na + Cl Na + Cl

• In both cases, we have formed an ion where the new valence shell is full, having eight electrons.

Formation of Sodium Chloride

88

Sodium metal, Na Chlorine gas, Cl2 Sodium Chloride, NaCl


89

The Octet Rule

• The Na+ and Cl- ion’s electron configurations are the same as that of the nearest noble gas (the ions are said to be isoelectronic with the nearest noble gas). Atoms “prefer” to have a filled outermost shell because this is more electronically stable.

• This can be generalized into the octet rule: Main-group elements tend to undergo reactions that leave them with eight outer-shell electrons.

– That is, main-group elements react so that they attain a noble gas e- configuration with filled s and p sublevels in their valence electron shell.

– There are many exceptions to the octet rule, but it is useful for making predictions about some chemical bonds.

Lewis Structures of Ionic Compounds

1. Use Lewis structures to predict the formula of the ionic compound formed by the following elements.

a. Ca and Cl

b. Ca and S

c. Na and S

90


91

Energy Changes in the Formation of NaCl

• The energy changes in the formation of NaCl from Na(s) and Cl2(g) are shown on the next slide:

– Step 1: sublimation of Na(s) to Na(g): Na(s) → Na(g) DH = +108 kJ/mol

– Step 2: ionization of Na(g) into Na+(g) ions: Na(g) → Na+(g) + e- DH = +496 kJ/mol

– Step 2: dissociation of Cl2(g) into Cl(g) atoms: ½Cl2(g) → Cl(g) DH = +122 kJ/mol

– Step 4: addition of e- to Cl(g) to form Cl-(g) ions: Cl(g) + e- → Cl-(g) DH = -349 kJ/mol

• The total so far:

Na(s) + ½Cl2(g) → Na+Cl-(g); DH = +377 kJ/mol

Born-Haber Cycle for the Formation of NaCl

92


93

Energy Changes in the Formation of NaCl

• But we don’t find NaCl in the gas phase (unless it’s really hot outside!):

– Step 5: formation of solid NaCl crystals from isolated Na+ and Cl- ions in the gas phase: Na+Cl-(g) → Na+Cl-(s) DH = -788 kJ/mol

• Revised total:

Na(s) + ½Cl2(g) → Na+Cl-(s); DH = -411 kJ/mol

– Since this is a release of energy, this is a much more favorable energetic process.

– The energy change in step 5 drives the overall process of the formation of NaCl.

– The energy changes are summarized in a diagram called a Born-Haber cycle (previous slide).

94

Lattice Energy

• The energy change in step 5 is called the lattice energy (DHlattice), the energy change associated with the formation of a crystal lattice from isolated ions in the gas phase. It is the result of the electrostatic interactions between ions, and thus it is a measure of the strength of the ionic bonds in an ionic crystal.

• Ionic solids exist only because the lattice energy drives the energetically unfavorable electron transfer. The energy required for elements to gain or lose electrons is supplied by the electrostatic attraction between the ions they form.


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Trends in Lattice Energy

• The force F that results from the interaction of electric charges is described by Coulomb’s law:

where k is a constant, z1 and z2 are the charges on the ions, and d is the distance between their centers.

• The lattice energy is the force F times distance:

– As the size of the ions increases, d becomes larger, and the lattice energy decreases.

– As the magnitude of the cation and anion charge increases, the lattice energy also increases (NaF = -923 kJ/mol, MgO = -3916 kJ/mol).

2

21 d

zzkF

d

zzkFdH 21

lattice D

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Trends in Lattice Energy

DHlattice (kJ/mol)

LiCl -834 NaCl -787 KCl -701 CsCl -657

DHlattice (kJ/mol)

NaF -923 CaO -3414


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Examples: Lattice Energies

2. Which compound in each of the following pairs of ionic substances has the most negative lattice energy?

a. NaCl or KCl

b. LiF or LiCl

c. NaCl or Na2O

d. MgO or BaS

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Ionic Solids and Crystals

• The ionic bond is the strong attraction between the cations and the anions. Unlike molecules, the cation and anion are not physically joined together.

• Thus, there is no molecule of NaCl; ionic compounds instead form ionic solids, which contain equal amounts of positive and negative charge surrounding each other in a regular array called a crystal.


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Properties of Ionic Compounds

• The oppositely charged ions are attracted to each other by electrostatic forces forming an ionic bond.

– The substance that forms is an ionic solid.

– The solid consists of a three-dimensional array called a crystal, which consists of cations surrounded in some fashion by anions

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• A typical ionic compound, such as rock salt (NaCl) is hard (doesn’t dent), rigid (doesn’t bend), and brittle (cracks without deforming).

• The attractive forces in ionic compounds hold the ions in specific positions; moving the ions out of position requires overcoming these forces, so the sample resists denting and bending.


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• Ionic compounds have very high melting points (NaCl melts at 801ºC, MgO melts at 2852ºC), and extremely high boiling points, because it takes a lot of heat energy to overcome the electrostatic interactions between cations and anions.

MP (ºC) BP (ºC)

CsBr 636 1300

NaI 661 1304

MgCl2 714 1412

KBr 734 1435

CaCl2 782 >1600

NaCl 801 1413

LiF 845 1676

KF 858 1505

MgO 2852 3600

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Solid ionic compounds do not conduct electricity, because the ions are not free to move.

When melted, or dissolved in water, the ions are free to move, allowing electricity to be conducted through the solution.


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Metallic Bonding: The Electron Sea Model

• In a solid metal, each metal atom donates one or more of its valence electrons to form an electron sea that surrounds the metal atoms.

– Metals conduct electricity because the electrons in the electron sea are free to move around.

– Metals conduct heat because the electrons help to disperse thermal energy throughout the metal

– Metals are malleable and ductile because there are no localized bonds between the metal atoms, allowing the metal to be deformed easily.

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