Models of Chemical Bonding

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Models of Chemical Bonding
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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
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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

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Figure 9.1
A general comparison of metals and nonmetals
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Types of Chemical Bonding
Typically
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
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Figure 9.2
The three models of chemical bonding
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Lewis Electron-Dot Symbols
- A method for depicting valence electrons and
interactions of atoms
For main group elements -
Example
Nitrogen, N, is in Group 5A and therefore has 5
valence electrons.
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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.
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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.

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SAMPLE PROBLEM 9.1
Depicting Ion Formation
PLAN
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.
SOLUTION
2 Na
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Three ways to represent the formation of Li and
F- through electron transfer.
Figure 9.4
Electron configurations
Li 1s22s1

F 1s22s22p5
Li 1s2
Orbital diagrams
Lewis electron-dot symbols
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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.

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Calculating lattice energy

• Lattice energy cannot be directly measured, so it
  is found by using Hesss 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

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Figure 9.6
The Born-Haber cycle for lithium fluoride
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Periodic Trends in Lattice Energy
Coulombs Law
charge A X charge B
electrostatic force a
distance2
energy force X distance therefore
charge A X charge B
electrostatic energy a
distance
cation charge X anion charge
a DH0lattice
electrostatic energy a
cation radius anion radius
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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.

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Figure 9.7
Trends in lattice energy
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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
  its shape until enough energy is applied to
  shift positions and crack the crystal.

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Figure 9.8
Electrostatic forces
and the reason ionic compounds crack.
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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

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Figure 9.9
Electrical Conductance and Ion Mobility
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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

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Table 9.1 Melting and Boiling Points of Some
Ionic Compounds
Compound
mp (0C)
bp (0C)
CsBr
1300
636
661
NaI
1304
MgCl2
714
1412
KBr
734
1435
CaCl2
782
gt1600
NaCl
801
1413
LiF
845
1676
KF
858
1505
MgO
2852
3600
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Figure 9.10
Vaporizing an ionic compound.
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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.

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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

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Figure 9.11
Covalent bond formation in H2.
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Figure 9.12
The attractive and repulsive forces in covalent
bonding.
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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.

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Figure 9.13
Bond length and covalent radius.
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• 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.

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SAMPLE PROBLEM 9.2
Comparing Bond Length and Bond Strength
PROBLEM
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
PLAN
(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.
SOLUTION
(a) Atomic size increases going down a group.
(b) Using bond orders we get
Bond length S - Br gt S - Cl gt S - F
Bond strength S - F gt S - Cl gt S - Br
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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 1550C
• Diamond (C) hardest known substance, mp 3550C

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Strong forces within molecules and weak forces
between them.
Figure 9.14
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Figure 9.15
Covalent bonds of network covalent solids.
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Properties of covalent cmpds

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

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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

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Figure 9.16
The Pauling electronegativity (EN) scale.
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Figure 9.17
Electronegativity and atomic size.
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Representation of Polar Bonds
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SAMPLE PROBLEM 9.3
Determining Bond Polarity from EN Values
PROBLEM
(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.
PLAN
(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.
SOLUTION
(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 lt N lt O
all have an EN larger than that of H.
H-C lt H-N lt H-O
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Nonpolar, Polar, or Ionic

• In general if the electronegativity difference
  between two bonded atoms is
• 0, usually between identical nonmetal atoms,
  called nonpolar covalent
• lt .4 , mostly covalent
• .4 to 1.7, 2 different nonmetals called polar
  covalent
• gt 1.7 , usually nonmetals and reactive metals, is
  mostly ionic
• Note there is no perfect ionic bond.

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Figure 9.18
Boundary ranges for classifying ionic character
of chemical bonds.
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Figure 9.19
Percent ionic character of electronegativity
difference (DEN).
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Figure 9.20
Li
F
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Figure 9.21
Properties of the Period 3 chlorides.
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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.

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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
  dont have to be broken to become liquid
• b.p. are very high because the cation and its
  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

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Table 9.5 Melting and Boiling Points of Some
Metals
Element
mp(0C)
bp(0C)
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Figure 9.23
Melting points of the Group 1A(1) and Group 2A(2)
elements.
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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.

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Figure 9.24
The reason metals deform.
metal is deformed
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Tools of the Laboratory
Infrared Spectroscopy
Figure B9.1
Some vibrational modes in general diatomic and
triatomic molecules.
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Tools of the Laboratory
Figure B9.1
Infrared Spectroscopy
Some vibrational modes in general diatomic and
triatomic molecules.
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Tools of the Laboratory
Infrared Spectroscopy
Figure B9.1
Some vibrational modes in general diatomic and
triatomic molecules.
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Tools of the Laboratory
Figure B9.2
The infrared (IR) spectrum of acrylonitrile.