Bonding: General Concept
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Types of Chemical Bonds
The Formation of Ions and Their Electron Configurations
Ionic Size and Charges, and the Relative Strength of Ionic Bonds
Energy of Formation of Ionic Compounds – The Born-Haber Cycle
Covalent Bond: Electronegativity, and Bond Polarity
Lewis Structures and the Octet Rule
Exceptions to the Octet Rule
Resonance Lewis Structures
Bond Energies
The Calculation of Enthalpy of Reaction from Bond Energy
The VSEPR Model and Molecular Shapes
Ionic bonds
• electrostatic attractions between cations and
anions;
• Bonds formed between metals and nonmetals
• Reactions that produce ionic bonds involve the
transfer of one, two, or three electrons from a
metal atom to nonmetal atom
Covalent Bonds
• One, two or three pairs of electrons shared
between two atoms
• Bonds between two nonmetals or between
semimetal and nonmetal atoms
• Bonds are formed when two atoms share electron
pairs
The Formation of Ions
• Ions are formed when metals react with
nonmetals, in which the metal atoms donate their
valence electrons to the nonmetals;
• Atoms of the representative metals lose their
valence electrons to become cations that have the
electron configurations of noble gases;
• The nonmetal atoms gain a number of electrons to
become anions that also have the electron
configuration of the noble gases;
Formation of Cations
• From the alkali metals (1A):
– M  M+ + e• From the alkaline Earth metals (2A):
– M  M2 + + 2e• From Group 3A metals: M  M3+ + 3e- ;
Formation of Anions
• From the halogen family (VIIA):
– X + e-  X • From the oxygen family (VIA):
– X + 2e-  X2• From N and P (in Group VA):
– X + 3e-  X3-
Common Ions of the Representative Elements
• Ions with the electron configuration of He = 1s2 – Li+ and H• Ions with the electron configuration of Ne = 1s2 2s2 2p6
– Na+, Mg2+, Al3+, F-, O2-, and N3• Ions with the electron configuration of Ar = 1s2 2s2 2p6 3s2 3p6
– K+, Ca2+, Sc3+, Cl-, S2-, and P3• Ions with the electron configuration of Kr =
1s22s22p63s23p64s23d104p6
– Rb+, Sr2+, Y3+, Br-, and Se2-;
• Ions with the electron configuration of Xe =
1s22s22p63s23p64s23d104p65s24d105p6
– Cs+, Ba2+, La3+, I-, and Te2-;
Cations From the Transition Metals
• Transition metal atom loses variable number of electrons
• Cations derived from transition metals have variable charges
• Cations of transition metals do not acquire the electron
configurations of noble gases
• Examples:
– Cr  Cr2+ + 2e-;
Cr2+: [Ar] 3d4
– Cr  Cr3+ + 3e-;
Cr3+: [Ar] 3d3
– Fe  Fe2+ + 2e-;
Fe2+: [Ar] 3d6
– Fe  Fe3+ + 3e-;
Fe3+: [Ar] 3d5
• (Note that these cations do not have the 4s electrons in their
electron configurations.)
Charge Density and The Strength of Ionic Bonds
Relative sizes of isoelectronic ions:
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Al3+ < Mg2+ < Na+ < Ne < F- < O2- < N3-;
Sc3+ < Ca2+ < K+ < Ar < Cl- < S2- < P3-;
Trend of ionic radii within a group:
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Li+ < Na+ < K+ < Rb+ < Cs+;
F- < Cl- < Br- < I-;
Ionic Bond Strength:
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Strength of Ionic bonds is related to charge density of the ions;
Greater charge and smaller ions lead to stronger ionic bond;
Ionic Bond Strength and Lattice Energy
• The strength of ionic bonds is associated with the magnitude of
the lattice energy
• Lattice energy - energy released when gaseous ions combine to
form a mole of solid ionic compound:
– M+(g) + X-(g)  MX(s);
UL = lattice energy
• Example: Na+(g) + Cl-(g)  NaCl(s);
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Li+(g) + F- (g)  LiF(s);
UL = -787 kJ/mol
UL = -1047 kJ/mol
• Lattice energy = k(q1q2/r);
– where q1 and q2 are charge magnitude on ions, r is the
internuclear distance, and k is the proportionality constant.
– Lattice energy increases with charge magnitude but decreases
with ionic size
Lattice Energies of Some Ionic Compounds
• Lattice Energy, UL(kJ/mol)
– The energy required to separate a mole of ionic solids into the
gaseous/vapor ions;
– MX(s)  M+(g) + X-(g)
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Mn+/XnLi+
Na+
K+
Mg2+
Ca2+
F1047
923
821
2957
2628
Cl853
787
715
2526
2247
Br807
747
682
2440
2089
I757
704
649
2327
2059
O22942
2608
2311
3919
3570
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The Born-Haber Cycle for NaCl
• Na+(g) + Cl(g) _______________
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-349 kJ
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+496 kJ _______ Na (g) + Cl (g)
• Na(g) + Cl(g)___________
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+121 kJ
• Na(g) + ½Cl2(g)________
? kJ
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+108 kJ
• Na(s) + ½Cl2(g)________
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-411 kJ
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NaCl(s)_________________
Chemical Processes in the Formation of NaCl
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Na(s)  Na(g);
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½Cl2(g)  Cl(g);
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Na(g)  Na+(g) + e-;
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Cl(g) + e-  Cl-(g);
• Na+(g) + Cl-(g)  NaCl(s);
• Na(s) + ½Cl2(g)  NaCl(s);
DHs = +108 kJ
½BE = +121 kJ
IE = +496 kJ
EA = -349 kJ
UL = ? kJ
DHf = -411 kJ
– UL = DHf – (DHs + ½BE + IE + EA)
(DHs = Enthalpy of sublimation; IE = Ionization energy;
BE = Bond energy; EA = Electtron affinity; UL = Lattice energy;
DHf = Enthalpy of formation)
The Born-Haber Cycle for LiF
• Li+(g) + F(g) _______________
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-328 kJ
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+520 kJ _______Li (g) + F (g)
• Li(g) + F(g)___________
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+77 kJ
• Li(g) + ½F2(g)________
? kJ
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+161 kJ
• Li(s) + ½F2(g)________
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-617 kJ
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LiF(s)_________________
Chemical Processes in the Formation of LiF
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Li(s)  Li(g);
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½F2(g)  F(g);
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Li(g)  Li+(g) + e-;
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F(g) + e-  F-(g);
• Li+(g) + F-(g)  LiF(s);
• Li(s) + ½F2(g)  LiF(s);
DHs = +161 kJ
½BE = +77 kJ
IE = +520 kJ
EA = -328 kJ
UL = ?
DHf = -617 kJ
– UL = DHf – (DHs + ½BE + IE + EA)
(DHs = Enthalpy of sublimation; IE = Ionization energy;
BE = Bond energy; EA = Electtron affinity; UL = Lattice energy;
DHf = Enthalpy of formation)
The Born-Haber Cycle for MgO
• Mg2+(g) + O2-(g) _____________
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+737 kJ
• Mg2+(g) + O(g)________
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+2180 kJ
• Mg(g) + O(g)_________
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+247 kJ
• Mg(g) + ½O2(g)________
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? kJ
+150 kJ
• Mg(s) + ½O2(g)________
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-602 kJ
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MgO(s)_________________
Chemical Processes in the Formation of MgO
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Mg(s)  Mg(g);
DHs = +150 kJ
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½O2(g)  O(g);
½BE = +247 kJ
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Mg(g)  Mg2+(g) + 2e-;
IE = +2180 kJ
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O(g) + 2e-  O2-(g);
EA = +737 kJ
• Mg2+(g) + O2-(g)  MgO(s);
UL = ? kJ
• Mg(s) + ½O2(g)  MgO(s);
DHf = -602 kJ
– UL = DHf – (DHs + ½BE + IE + EA)
(DHs = Enthalpy of sublimation; IE = Ionization energy;
BE = Bond energy; EA = Electtron affinity; UL = Lattice energy;
DHf = Enthalpy of formation)
Covalent Bonds
• Covalent bonds
– A bond between two nonmetal atoms or between a semimetal and a nonmetal atoms;
– Bonded atoms may share one, two, or three pairs of
electrons.
• Nonpolar covalent bonds are formed between identical atoms
or if bonded atoms have the same electronegativity.
• Polar covalent bonds are formed when bonded atoms have
different electronegativity;
• Polar covalent bonds are covalent bonds with partial ionic character
Potential Energy Diagram for Covalent
Bond Formation
Potential energy of H-atoms during the
formation of H2 molecule
Lewis Model for the Formation of Covalent
Bonds and Covalent Molecules
General trends:
• Electronegativity increases from left to right along a period
• For the representative elements (s and p block) the
electronegativity decreases as you go down a group
• The transition metal group is not as predictable as far as
electronegativity.
Electronegativity
• Electronegativity is the relative ability of a bonded atom to
attract shared electrons closer to itself.
– Electronegativity increases going across a period and
decreases going down a group.
– Most electronegative elements – at top right corner of PT
– Least electronegative elements – at bottom left corner of PT
– Fluorine (F) is most electronegative with EN value of 4.0
– Francium (Fr) is least electronegative with a value of 0.7
• The polarity of a covalent bond depends on the
electronegativity difference (DEN) of the two bonded
atoms.
Electronegativity and Bond Polarity
Compound
Electronegativity
Difference
Type of Bond
F2
HF
LiF
4.0 - 4.0 = 0
4.0 - 2.1 = 1.9
4.0 - 1.0 = 3.0
Nonpolar
covalent
Polar covalent
Ionic (noncovalent)
•In F2 the electrons are shared equally between the atoms, the
bond is nonpolar covalent
•In HF the fluorine atom has greater electronegativity than the
hydrogen atom.
•The sharing of electrons in HF is unequal: the fluorine atom attracts
electron density away from the hydrogen (the bond is thus a polar
covalent bond)
Electronegativity and bond polarity
The H-F bond can thus be represented as:
•The 'd+' and 'd-' symbols indicate partial positive and negative charges.
•The arrow indicates the "pull" of electrons off the hydrogen and towards
the more electronegative atom.
•In lithium fluoride the much greater relative electronegativity of the
fluorine atom completely strips the electron from the lithium and the result
is an ionic bond (no sharing of the electron)
Predicting Bond Type From Electronegativity
A general rule of thumb for predicting the type of bond based
upon electronegativity differences (DEN):
• If DEN between the two atoms is 0-0.5, the bond is non-polar
covalent;
• If DEN between the two atoms is greater than 0.5, but less
than 1.5, the bond is polar covalent
• If DEN between the two atoms is 1.5, or greater, the bond is
ionic
Bond Length
The bond length is defined as the distance between the nuclei of the
two atoms involved in the bond. In general, the larger the atoms
involved in a bond, the longer the bond length, and the more bonds
between two atoms, the shorter the bond length.
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Bond Energy
Bond energy is the energy required to break the bond(s)
between two atoms. In general, the shorter the bond, the higher
the bond energy.
Lewis Symbols for Atoms and The Formation of
Covalent Molecules
Lewis Symbols for O, F, and Na
How to draw a Lewis structure of a simple
molecular compound or polyatomic ion?
A Lewis structure can be drawn for a molecule or ion by following three steps:
1. Calculate the number of valence electrons (including charges, if any)
2. Write skeleton structure
– Lowest EN (electronegative) atom, largest atom, and/or atom forming most
bonds is usually the central atom
– Hydrogen and Fluorine cannot be the central atom.
– Connect all atoms with a single bond – using a line or two dots.
– In oxoacids, such as sulfuric acid (H2SO4), the oxygen bonds to central atom
and the H to oxygen.
– Compounds are usually compact and symmetrical structures
3. Count how many electrons have been used
– Distribute remaining electrons to terminal atoms filling them until each has
eight electrons (octet rule), unless it is hydrogen atom.
4. Central atom octet is filled last
– Any remaining electrons become lone pairs on central atom.
– If central atoms do not have an octet, move from terminal atoms one pair at a
time to form double and triple bonds.
Covalent Bonds and Lewis Structures Some Molecules
Lewis Structures of HF, H2O, NH3, & CH4
Lewis Structures of CH4, NH3 and H2O
Lewis Structures of CO2, HCN, and C2H2
Lewis Structures of Other Covalent Molecules
Assigning Formal Charges in Lewis Dot Structures?
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Formal Charge
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To determine the formal charge of an atom from a Lewis dot
structure we need to assign each electron to an atom in the
structure. To do this we use the following rules:
1. All nonbonding electrons (unshared electrons) are assigned
to the atom on which they are found.
2. Each atom in a bond is assigned ½ of the total number of
electrons in the bond (i.e. for a single bond each atom is
assigned 1 electron, for a double bond each atom is
assigned 2 electrons, etc.)
3. For each atom the number of electrons assigned in the
above steps is subtracted from the number of valence
electrons in the atom.
Choosing the correct or best Lewis structures based
on formal charges
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If two or more Lewis dot structures can be drawn which
satisfy the octet rule, the most stable one will be the
structure where:
1.
2.
The formal charges are as small as possible.
Any negative charges are located on the more electronegative
atoms.
Assigning Formal Charges
Which Lewis structures of CO2 & N2O are correct?
Resonance Lewis Dot Structures for CO32-
Resonance Lewis Structures of PO43-
Assigning Appropriate Formal Charges
Lewis Structures, Molecular Shapes & Polarity
The Shapes of Methane and Ammonia Molecules
The Shape of Water Molecules
Structures and Shapes of Formaldehyde and Ethylene
Bond Length and Bond Energies
• Bond length (pm) and bond energy (kJ/mol)
• Bond Length Energy Bond Length Energy
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H─H
C─C
N─N
O─O
F─F
Cl─Cl
Br─Br
I─I
C─C
C─N
C─O
O─O
O=O
C─Br
C─I
74
154
145
148
142
199
228
267
154
147
143
148
121
194
214
436
348
170
145
158
243
193
151
348
308
360
145
498
288
216
H─C
H─N
H─O
H─F
H─Cl
H─Br
H─I
109
101
96
92
127
141
161
413
391
366
568
432
366
298
C─C
C=C
C≡C
C─S
C─F
C─Cl
N─N
N≡N
154
134
120
182
135
177
145
110
348
614
839
272
488
330
170
945
Bond Breaking and Bond Formation in the
Reaction to form H2O
Using Bond Energies to Calculate Enthalpy
of Reactions in Gaseous State
• Chemical reactions in the gaseous state only involve:
– the breaking of covalent bonds in reactants and
– the formation of covalent bonds in products.
• Bond breaking requires energy input - an endothermic
process),
• while bond formation releases energy – an
exothermic process;
 DHreaction = S(Energy of bond breaking) + S(Energy of
bond formation)
Calculation of Reaction Enthalpy Using Bond
Energies
• Use bond energies to estimate the DH for the reaction:
CH3OH(g) + 2 O2(g)  CO2(g) + 2H2O(g);
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S(energy of bond breaking) (in kJ)
= 3 x BE(C─H) + BE(C─O) + BE(O─H) + 2 x BE(O═O)
= (3 x 413) + 358 + 467 + (2 x 495) = 3054 kJ
S(energy of bond breaking) (in kJ)
= 2 x -BE(C═O)* + 4 x -BE(O─H)
= (2 x -799) + (4 x -495) = -3578 kJ
DHreaction = 3054 + (-3578) = -524 kJ
Molecular Shapes of BeI2, HCl, IF2-, ClF3, and NO3-
The VSEPR Shapes
Linear and Trigonal Planar Electron-Pair Geometry
The Tetrahedral Electron-Pair Geometry
Trigonal Bipyramidal Electron-Pair Geometry
The Octahedral Electron-Pairs Geometry