John A. Schreifels

1
Chapter 9

• Ionic and Covalent Bonding

2
Overview

• Ionic Bonds
• Describing Ionic Bonds
• Electron Configuration of Ions
• Ionic Radii
• Covalent Bonds
• Describing Covalent Bonds
• Polar Covalent Bonds Electronegativity
• Writing Lewis Electron-Dot Formulae
• Bond Length and Order
• Bond Energy

3
IONIC BONDING

• Ionic bonds are electrically neutral groups that
  are held together by the attraction arising from
  the opposite charges of a cation and an anion.
• Substances that have ionic bonds in a solid form
  a salt having high melting point and high
  crystallinity.
• Bonding thought of as the result of the
  combination of neutral atoms with transfer of one
  or more electrons from one atom to the other.

4
LEWIS SYMBOLS AND THE OCTET RULE

• It was observed that the electron configuration
  of many substances after ion formation was that
  of an inert gas ? octet rule.
• Octet rule Main-group elements gain, lose, or
  share in chemical bonding so that they attain a
  valence octet (eight electrons in an atoms
  valence shell).
• E.g. The electron configuration of each reactant
  in the formation of KCl gives
• K is that of Ar
• Cl? is also that of Ar.
• The other electrons in the atom are not as
  important in determining the reactivity of that
  substance.
• The octet rule is particularly important in
  compounds involving nonmetals.

5
Energy in Ionic Bonding

• When potassium and chlorine atoms approach each
  other we have
• K(g)? K(g) e? Ei 418 kJ
• Cl(g) e?? Cl?(g) Eea ?349 kJ
• K(g)Cl(g)? K(g) Cl?(g) ?E 69 kJ
• Positive ?E energy absorbed ? energetically not
  allowed.
• Driving force must be the formation of the
  crystalline solid.
• K(g) Cl?(g) ? KCl(s)

6
Formation of Crystalline Lattice

• Energy of crystallization estimated from
  Coulombs Law by
• assuming ions are spheres.
• Use ionic radii to determine charge separation.
• rK 133x10?12 m rCl? 181x10?12 m
• d 133x10?12 m 181x10?12 m 314x10?12 m
• z1 z2 1.602x10?12 C(oulombs) actually one is
  the negative of the other.
• k 8.99x109 J?m/C2

• This is related to the negative of the lattice
  energy, as discussed later.

7
BORN-HABER CYCLE AND LATTICE ENERGIES

• Overall energetics for the formation of
  crystalline solids can be determined from a
  Born-Haber cycle that accounts for all of the
  steps towards the formation of solid salts from
  the elements. For the formation of KCl from its
  elements we have
• Net energy change of ?434 kJ/mol indicates
  energetically favored.
• Energy for the fifth step is the negative of the
• lattice energy energy required to break ionic
  bonds and sublime (always positive).
• E.g. Determine the lattice energy of BaCl2 if
  the heat of sublimation of Ba is 150.9 kJ/mol and
  the 1st and 2nd ionization energies are 502 and
  966 kJ/mol, respectively. The heat for the
  synthesis of BaCl2(s) from its elements is
  ?806.06 kJ/mol.

8
Energy Level Diagram of Born Haber Cycle
9
LATTICE ENERGIES AND PERIODICITY

• Lattice energy can also be determined from
  Coulombs law
• Directly proportional to charge on each ion.
• Inversely proportional to size of compound (sum
  of ionic radii).
• Table (right) presents the lattice energies for
  alkali and alkaline earth ionic compounds. The
  lattice energies
• decrease for compounds of a particular cation
  with atomic number of the anion.
• decrease for compounds of a particular anion with
  atomic number of the cation.

10
Ionic Radii

• Ionic radius a measure of the size of a
  spherical region around the nucleus of an atom
  where electrons are most likely to reside.
• Cation loses electrons from its valence shell.
  Electrons from other valence shells are closer to
  the nucleus.
• Cation also has more protons than electrons which
  adds to the pull on the remaining electrons and
  decreases the radius.

• Anion has more electrons than protons the pull
  of the nucleus is less per valence electron.
  Also, the electron electron repulsion is
  greater. These lead to larger radius for an
  anion.

11
Ionic Radii - Trends

• Ionic radii increase down any column because of
  the addition of electron shells.
• In general, across any period the cations
  decrease in radius. When you reach the anions,
  there is an abrupt increase in radius, and then
  the radius again decreases.

12
Ionic Radii - Isoelectronic Ions

• Isoelectronic substances have the same number of
  electrons and electron configuration.

All have 18 electrons

• Largest radius ion with smallest number of
  protons.
• Smallest radius ion (atom) with largest number
  of protons.

13
The Covalent Bond

• Repulsive forces of the electrons offset by the
  attractive forces between the electrons and the
  two nuclei.
• Most stable bond energy and bond distance
  characterizes bonds between two atoms.
• Strengths of Covalent Bonds
• Bonds form because their formation produces lower
  energy state than when atoms are separated.
• Breaking bonds increases the overall energy of
  the system. Energy for breaking bonds has a
  positive sign (negative means that energy is
  given off).
• H - H (g) ? 2H(g) DH 436 kJ.
• Ionic vs. Covalent Bonds
• Ionic compounds have high melting and boiling
  points and tend to be crystalline
• Covalently bound compounds tend to have lower
  melting points since the attractive forces
  between the molecules are relatively weak.

14
Lewis Structures

• Lewis structure valence electrons represented by
  dots and are placed where they would be in any
  bonding that might exist.
• Lewis structures of second row elements
• H2 BH3
• CH4 NH3
• H2O HF
• Each has 8 electrons around the central atom
  thus we can predict the number of bonds that will
  form from the position in the periodic table.
• E.g. The structure of chlorine is

• Bonding electrons shared electrons.
• Non-bonding or lone pair unshared electrons

15
Lewis Structures(contd)

• Octet can be filled by donation of electrons from
  each atom or one atom can supply both electrons.
• E.g. H NH3 ? . "co-ordinate covalent bond".
  
• E.g.2
• Multiple bonds may form as a result when the two
  atoms forming the bond do not have enough
  electrons.
• OO
• N?N
• Multiple bonds are shorter and stronger than
  single bonds because of the extra electrons
  holding the two atoms together.

16
Polar Bonds Electronegativity

• Electronegativity is a measure of the atoms
  ability to gain or lose electrons. It is
  directly related to its ionization tendency and
  its ability to form the inert gas configuration.
  Obtained by
• where Ei ionization energy and Eea the
  electron affiinity.
• E.g. Li has a very low ionization energy and
  electron affinity, while Cl has a both a high
  ionization energy and high electron affinity.
  Electronegativity will be high for Cl and low for
  Li.
• Fluorine has the highest electronegativity of
  4.0.
• Electronegativities (see Fig. 9.15)
• increase from bottom to top of periodic table and
  
• increase to a maximum towards the top right.
• Combination of elements with intermediate
  electronegativities forms bonds that are
  intermediate between covalent and ionic.
• can provide an insight as to the type of bond
  that would be expected.
• Ionic bonds formed when ?? ? 2
• covalent bonds forms when ?? ? 1.
• Polar covalent forms when 1 ? ?? ? 2, the bonding
  is "intermediate" between the two.

17
Polar Bonds Electronegativity2

• E.g.1 Determine the polarity of the N H in
  NH3.
• E.g. 2 Predict the type of bond formed in CCl4.
• The magnitude of ?? indicates if electrons are
  polarized around one element in preference to the
  other.
• Polar bond polar. With intermediate ??, a small
  charge on the atom due to that bond develops. ?
  and ?? designates which is the positive and
  negative side respectively.
• E.g.3 Determine the relative polarities of HF,
  HCl, HBr and HI.

18
Lewis Structures of Polyatomic Molecules

• Procedure for more complicated molecules
• Determine the total number of valence electrons
  from each atom.
• Distributed atoms around the central atom (least
  electronegative. Hydrogen atoms are usually
  attached to any oxygen.
• Satisfy the octet of the atoms bonded to the
  central atom.
• Satisfy the octet of the central atom by
  distributing the remaining electrons as electron
  pairs around it. (multiple bonds may be
  necessary)
• E.g. Determine the Lewis structure of H2SO4.
• E.g. Draw the Lewis dot structures of NCl3, CSe2,
  and CO.

19
FORMAL CHARGES

• Formal Charge (of an atom in a Lewis formula) the
  hypothetical charge obtained by assuming that
  bonding electrons are equally shared between the
  two atoms involved in the bond. Lone pair
  electrons belong only to the atom to which they
  are bound.
• E.g. determine the formal charge on all
  elements PCl3, PCl5, and HNO3.
• formal charge (FC) allows the prediction of the
  more likely resonance structure.
• To determine the more likely resonance structure
• FC should be as close to zero as possible.
• Negative charge should reside on the most
  electronegative and positive charge on the least
  electronegative element.
• E.g. draw the resonance structures of H2SO4
  determine the formal charge on each element and
  decide which is the most likely structure.

20
Lewis Structures and Resonance

• Quantum theory indicates that any position
  possible for an electron.
• Equivalent electron positions often possible
• E.g. SO2 OS-O and O-SO.
• Each structure equally likely.
• the true form of the molecule is a hybrid of
  these and is called resonance and the hybrid form
  is called a resonance hybrid.

21
Exceptions to the Octet Rule

• Although many molecules obey the octet rule,
  there are exceptions where the central atom has
  more than eight electrons.
• Generally, if a nonmetal is in the third period
  or greater it can accommodate as many as twelve
  electrons, if it is the central atom.
• These elements have unfilled d subshells that
  can be used for bonding.
• E.g determine the Lewis dot structure of XeF4,
  ICl3, and SF4

22
Bond Dissociation Enthalpies

• Bond dissociation energy, D the energy required
  to break one mole of a type of bond in an
  isolated molecule in the gas phase.
• Useful for estimation of heat of unknown
  reactions.
• Average bond energies listed in tables (e.g. C
  H bond) rest pf structure not very important
• HO-H bond in H2O and CH3O-H bond are 492 and 435
  kJ/mol.
• Hesss law can be used with bond dissociation
  energies to estimate the enthalpy change of a
  reaction. The breaking in a C H bond would be C
  H(g) ? C(g) H(g) ?H D 410 kJ.
• Sign always positive since energy must be
  supplied to break bond.

23
Using Bond Dissociation Enthalpies

• E.g. Estimate the heat of formation of H2O(g)
  from bond dissociation energies. Thus determine
  
• H2(g) ½ O2(g) ? H2O(g) ?
• From the book (Table 9.5)
•   H H (g) ? 2H(g) ?H D1 436 kJ 
• ½ OO ? O(g) ?H D2 494/2 247 kJ
• 2H(g) O(g) ? H O H (g) ?H ?2D3 ?2459
  kJ  
• H2(g) ½ O2(g) ? H2O(g) ?235 kJ
• Actual ?241.8 kJ
• Can be determined by suming all the energies for
  the bonds broken and subtract from if the sum of
  the energies for the bonds formed.
• E.g. 2 Estimate the energy change for the
  chlorination of ethylene
• CH2CH2(g) Cl2(g)? CH2ClCH2Cl

24
Using Bond Dissociation Enthalpies

• It may be necessary to include a phase change
  since many reactions or reactants are not in the
  gas phase.
• E.g. Determine the heat of formation of CCl4(l).
  
• Solution The reaction is
• C(gr) 2Cl2(g) ? CCl4(l) ?
• Write the reactions and sum energies

25
Electronegativities
Return to slide 16
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