9-1

Models of Chemical Bonding

9-2

Models of Chemical Bonding

9.1 Atomic Properties and Chemical Bonds

9.2 The Ionic Bonding Model

9.3 The Covalent Bonding Model

9.4 Between the Extremes: Electronegativity and Bond Polarity

9.5 An Introduction to Metallic Bonding

9-3

Chemical Bonds

Chemical Bonds– The attractive forces that hold atoms or ions together to

form molecules or crystals

Octet rule – atoms tend to gain, lose, or share valence electrons to

get an octet.– Everything wants to be like a noble gas.– exceptions

• near He obey duet rule • Transition metals • n = 3 and above

9-4

A general comparison of metals and nonmetals

Figure 9.1

9-5

Types of Chemical Bonding

1. Metal with nonmetal:

electron transfer leads to ionic bonding

2. Nonmetal with nonmetal:

electron sharing leads to covalent bonding

3. Metal with metal:

electron pooling leads to metallic bonding

Typically:

9-6

Figure 9.2 The three models of chemical bonding

9-7

Lewis Electron-Dot Symbols

For main group elements -

Example:

Nitrogen, N, is in Group 5A and therefore has 5 valence electrons.

N:.

..

:

N .. ..N :.

. :N ...

The A group number gives the number of valence electrons.

Place one dot per valence electron around the four sides of the element symbol. Do not pair dots until all four sides have an electron.

- A method for depicting valence electrons and interactions of atoms

9-8

Figure 9.3

Lewis electron-dot symbols for elements in Periods 2 and 3

Nonmetals - The number of unpaired dots indicates the number of electrons it gains, or the number of covalent bonds it usually forms.Metals – The total number of dots is the maximum number of electrons it may lose when forming a cation.

9-9

Sec 9.2 Ionic Bonding

• In ionic bonding, electrons are gained or lost, the resulting bonds are based on electrostatic attraction.

• ex Na has 1 valence e, Cl has 7

• If Na could only get rid of 1, if Cl could only gain 1….

• When sodium metal is placed in Cl2 gas, they react by

transferring 1 e- from Na to Cl to form Na+ and Cl-. Now each has an octet. Because both now have a charge they are attracted to each other to form NaCl.

9-10

SAMPLE PROBLEM 9.1 Depicting Ion Formation

PLAN:

SOLUTION:

PROBLEM: Use partial orbital diagrams and Lewis symbols to depict the formation of Na+ and O2- ions from the atoms, and determine the formula of the compound.

Draw orbital diagrams for the atoms and then move electrons to make filled outer levels. It can be seen that 2 sodiums are needed for each oxygen.

3s 3p

Na

3s 3p

Na2s 2p

O2s 2p

O2-

2 Na+

:Na

Na+ O

.:

..

.2Na+ + O 2-

:: ::

9-11

Electron configurations

Li 1s22s1

Orbital diagrams

Lewis electron-dot symbols

+ F 1s22s22p5 Li+ 1s2 + F- 1s22s22p6

Three ways to represent the formation of Li+ and F- through electron transfer.

Figure 9.4

Li

1s 2s 2p

F

1s 2s 2p

+

Li+

1s 2s 2p

F-

1s 2s 2p+

.+ F: ::Li . Li+ + F -::

::

9-12

Energy in Ionic Bonding• Li(g) Li+

(g) + e- IE1 = 520 kJ• F(g) + e- F-

(g) EA = -328 kJ• So the process would appear to be endothermic• Li(g) + F(g) Li+

(g) + F-(g) E = 192kJ

• Overall the process is very exothermic, this is because of the lattice energy.– the enthalpy change of gaseous ions coalescing into a crystalline

solid.– Indicates the strength of the two ions attraction– Influences melting point, hardness, and solubility

• Ionic solids exist only because the lattice energy drives the unfavorable electron transfer.

9-13

Calculating lattice energy• Lattice energy cannot be directly measured, so it is

found by using Hess’s Law.• The enthalpy change for an overall reaction is the

sum of the enthalpy changes of the reactions which make it up.

• Lattice energies are calculated by using a Born-Haber Cycle– A series of chosen steps from elements to ionic

compounds for which all the enthalpies are known.– The steps are hypothetical and not the actual steps of

the process

9-14

Figure 9.6 The Born-Haber cycle for lithium fluoride

9-15

Periodic Trends in Lattice Energy

Coulomb’s Law

charge A X charge B

electrostatic force distance2

energy = force X distance therefore

charge A X charge B

electrostatic energy distance

cation charge X anion charge

electrostatic energy cation radius + anion radius

H0lattice

9-16

Trends in lattice energy

• Effect of ion size.– Increasing the size of the ions decreases lattice energy,

therefore attraction between cations and anions decreases down in a group

• Effect of ionic charge.– Increasing the charge of the ions increases the lattice

energy.

9-17

Figure 9.7

Trends in lattice energy

9-18

Properties of ionic compounds

• Ionic compounds are hard, rigid, and brittle– This is a result of ions being held in specific positions

in a crystal. So a crystal retains it’s shape until enough energy is applied to shift positions and crack the crystal.

9-19

Figure 9.8

Electrostatic forces and the reason ionic compounds crack.

9-20

Properties of ionic compounds

• Do not conduct electricity in the solid state– Ions in fixed positions

• Do conduct when melted or dissolved– Ions can move independently

9-21

Figure 9.9 Electrical Conductance and Ion Mobility

Solid ionic compound

Molten ionic compound

Ionic compound dissolved in water

9-22

Properties of ionic compounds

• High melting and boiling points (all solid at RT)– Enough energy must be supplied to free ions from the

attractions of the surrounding ions

• Ionic compounds vaporize as ion-pairs even though no “molecules” exist in the crystal

9-23

Table 9.1 Melting and Boiling Points of Some Ionic Compounds

Compound mp (0C) bp (0C)

CsBr

661

1300

NaI

MgCl2

KBr

CaCl2

NaCl

LiF

KF

MgO

636

714

734

782

801

845

858

2852

1304

1412

1435

>1600

1413

1676

1505

3600

9-24

Figure 9.10

Vaporizing an ionic compound.

9-25

Sec 9.3 Covalent Bonding• Elements can also form octets by sharing e- between them,

the bonds that result are called covalent bonds.

• usually occurs in nonmetal/nonmetal compounds.

• More compounds are covalent than ionic.

• A single Cl atom has 7 valence electrons, in a sample of pure Cl one atom cannot steal an electron from another, so they share to form Cl2.

• Molecule – a compound formed by 2 or more atoms joined by covalent bonds that behaves as a single particle.

9-26

Covalent Bonding• atoms share electrons by overlapping orbitals so electrons

can exist in the orbitals of both atoms at once.

• These shared or bonding pairs of electrons are represented by lines in structures.

• Other valence electrons that are not involved in bonding are called unshared or lone pairs.

• Every pair of electrons shared between atoms is a bond– 1 pair – single bond

– 2 pairs – double bond , stronger

– 3 pairs – triple bond, strongest

– aka bond order

– Lewis structures of Cl2, O2, N2

9-27

Figure 9.11

Covalent bond formation in H2.

9-28

Figure 9.12

The attractive and repulsive forces in covalent bonding.

9-29

Bond Energy

• Bond energy or Bond Strength - the energy required to overcome the attraction of covalently bonded atoms.– It is defined as energy required to break bonds in 1

mole of gaseous atoms.

– Bond energy depends on the specific elements involved.

– It can vary from molecule to molecule so table values are averages.

9-30

9-31

9-32

Figure 9.13

Internuclear distance(bond length)

Covalent radius

Internuclear distance(bond length)

Covalent radius

Internuclear distance(bond length)

Covalent radius

Internuclear distance(bond length)

Covalent radius

Bond length and covalent radius.

9-33

• Bond length, bond energy, and bond order are closely related– Higher bond order is shorter, and stronger for a given

set of atoms

– With a constant bond order, longer bonds are usually weaker.

9-34

9-35

SAMPLE PROBLEM 9.2 Comparing Bond Length and Bond Strength

PROBLEM:

PLAN:

SOLUTION:

Using the periodic table, but not Tables 9.2 and 9.3, rank the bonds in each set in order of decreasing bond length and bond strength:

(a) S - F, S - Br, S - Cl (b) C = O, C - O, C O

(a) The bond order is one for all and sulfur is bonded to halogens; bond length should increase and bond strength should decrease with increasing atomic radius. (b) The same two atoms are bonded but the bond order changes; bond length decreases as bond order increases while bond strength increases as bond order increases.

(a) Atomic size increases going down a group.

Bond length: S - Br > S - Cl > S - F

Bond strength: S - F > S - Cl > S - Br

(b) Using bond orders we get

Bond length: C - O > C = O > C O

Bond strength: C O > C = O > C - O

9-36

Properties of covalent cmpds• The physical properties of molecular compounds, are not

related to the strength of their covalent bonds.– Most covalent compounds have low m.p. and b.p. because the

strong covalent bonding is typically isolated within molecules. The attractions between separate molecules, called intermolecular forces, are what must be overcome to melt or boil these covalent substances.

• The physical properties of network covalent solids, are related to the strength of their covalent bonds.– In these substances there are no individual molecules, the covalent

bonding extends in 3-D throughout the substance.

– Ex Quartz (SiO2) very hard, mp 1550°C

– Diamond (C) hardest known substance, mp 3550°C

9-37

Figure 9.14 Strong covalent bonding forces within molecules

Weak intermolecular forces between molecules

Strong forces within molecules and weak forces between them.

9-38

Figure 9.15 Covalent bonds of network covalent solids.

9-39

Properties of covalent cmpds

• Most covalent substance are poor electrical conductors, when solid, liquid, or dissolved.

9-40

Sec. 9.4 Between the extremes• Most real bonds fall somewhere between the ideal of ionic

or covalent bonding theory.

• The type of bond atoms form depends on electronegativity– some atoms attract e- more strongly than others, we say these are

more electronegative

– electronegativity increases going right and up the table

• Covalent bonds in which e- are not shared equally because of electronegativity differences are called polar covalent bonds

9-41

Figure 9.16 The Pauling electronegativity (EN) scale.

9-42

Figure 9.17 Electronegativity and atomic size.

9-43

Representation of Polar Bonds

9-44

SAMPLE PROBLEM 9.3 Determining Bond Polarity from EN Values

PROBLEM:

PLAN:

SOLUTION:

(a) Use a polar arrow to indicate the polarity of each bond: N-H, F-N, I-Cl.

(b) Rank the following bonds in order of increasing polarity: H-N, H-O, H-C.

(a) Use Figure 9.16(button at right) to find EN values; the arrow should point toward the negative end.

(b) Polarity increases across a period.

(a) The EN of N = 3.0, H = 2.1; F = 4.0; I = 2.5, Cl = 3.0

N - H F - N I - Cl

(b) The order of increasing EN is C < N < O; all have an EN larger than that of H.

H-C < H-N < H-O

9-45

Nonpolar, Polar, or Ionic• In general if the electronegativity difference between two

bonded atoms is:

• 0, usually between identical nonmetal atoms, called nonpolar covalent

• < .4 , mostly covalent

• .4 to 1.7, 2 different nonmetals called polar covalent

• > 1.7 , usually nonmetals and reactive metals, is mostly ionic

• Note: there is no perfect ionic bond.

9-46

Figure 9.18EN

3.0

2.0

0.0

Boundary ranges for classifying ionic character of chemical bonds.

9-47

Figure 9.19

Percent ionic character of electronegativity difference (EN).

9-48

Figure 9.20

Li F

9-49

Figure 9.21

Properties of the Period 3 chlorides.

9-50

Sec. 9.5 Metallic Bonding

• Solid Metals and metal alloys have metallic bonding

• Electron Sea Model – All the metal atoms contribute their valence electrons

to a delocalized pool of electrons. The metal cations are held together by attraction to the delocalized electrons.

9-51

Properties of Metals

• Most are solid at RT, with moderate to high mp, and very high b.p.– m.p. are not very high because the metallic bonds don’t

have to be broken to become liquid– b.p. are very high because the cation and it’s electrons

must be separated from the others

• m.p. are higher for metals with more valence electrons – cation charges are higher resulting in greater cation-

electron sea attractions

9-52

Table 9.5 Melting and Boiling Points of Some Metals

Element mp(0C) bp(0C)

Lithium (Li) 180 1347

Tin (Sn) 232 2623

Aluminum (Al) 660 2467

Barium (Ba) 727 1850

Silver (Ag) 961 2155

Copper (Cu) 1083 2570

Uranium (U) 1130 3930

9-53

Figure 9.23

Melting points of the Group 1A(1) and Group 2A(2) elements.

9-54

Properties of Metals

• Metals are good conductors of electricity when solid, or liquid.– The delocalized electrons are able to move under an

electric field

• Metals are good conductors of heat.– The delocalized electrons disperse heat more quickly

• Metals are malleable and ductile, not brittle– The cations are able to slide past each other and still

retain their attraction to the electron sea.

9-55

Figure 9.24

metal is deformedThe reason metals deform.

9-56

Infrared Spectroscopy

Tools of the Laboratory

Figure B9.1

Some vibrational modes in general diatomic and triatomic molecules.

9-57

Infrared Spectroscopy

Tools of the Laboratory

Figure B9.1

Some vibrational modes in general diatomic and triatomic molecules.

9-58

Infrared SpectroscopyTools of the Laboratory

Figure B9.1

Some vibrational modes in general diatomic and triatomic molecules.

9-59

Tools of the Laboratory

Figure B9.2

The infrared (IR) spectrum of acrylonitrile.