CH 8: Bonding

1
CH 8 Bonding

• General Concepts

2
Chapter Outline Part I

• Types of chemical bonds (8.1)
• Electronegativity and bond polarity (8.2/3
• Ions (8.4)
• Energy changes when a binary ionic compound forms
  (8.5)
• Ionic character of covalent bonds (8.6)

3
Introduction to Bonding

• Chemical bond force that holds atoms together
  so that they function as a unit.
• Consider 2 classes of bonds
• Ionic bonding
• Covalent bonding

4
Bond Types

• Ionic bonds attractive forces among oppositely
  charged ions
• Forms when a metal loses electron(s) to a
  nonmetal.
• Bond strength can be calculated using Coulombs
  law

5
Ionic Bonds

• Strength of the attraction between the ions can
  be calculated using Coulombs law.
• E (2.31 x 10-19 J nm) (Q1Q2/r)
• Q1 and Q2 are the charges on the ions.
• r distance between ion centers in nm

6
Using Coulombs Law

• E (2.31 x 10-19 J nm) (Q1Q2/r)
• Sign on E???
• The more negative E, the stronger the attractive
  force between the ions.

7
Using Coulombs Law

• E (2.31 x 10-19 J nm) (Q1Q2/r)
• Magnitude of E.
• E is more negative when

8
Covalent Bonds

• Covalent bond bonded atoms share pairs of
  valence electrons
• Covalent bonding results in formation of a
  molecule.
• Covalent bonding occurs between nonmetals.

9
Types of Covalent Bonds

• Pure covalent bond electrons are shared by
  like nonmetals
• E.g. diatomic molecules
• Results in equal sharing of the electrons
• Aka nonpolar covalent bond

10
Types of Covalent Bonds

• Polar covalent bond unequal sharing of
  electrons by the bonded atoms
• bond between different nonmetals each with its
  own ability to attract the shared electrons

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Polar Covalent Bonds

• Showing bond polarity
• Consider the HF molecule.
• See board and/or page 346.
• Experimental determination of bond polarity, page
  344

12
Bond Polarity

• To predict bond polarityconsider the
  electronegativity (EN) of the bonded atoms.
• EN the ability of an atom in a molecule to
  attract shared electrons.

13
EN Values

• The higher the EN the greater the atoms ability
  to attract shared electrons.
• EN values and the periodic table
• EN ________ down a group.
• EN ________ across a period.
• See back of the periodic table.

14
EN and Bond Polarity

• As the difference in EN between bonded atoms
  increases so does the polarity of the bond.
• Can also say that the ionic character of the bond
  is increasing.
• See table 8.1 on page 346.

15
Bond Polarity and Dipoles

• Polar molecules have a preferred orientations
  when placed in an electric field.
• Said to have a dipole moment.
• Dipole moment molecule has a center of positive
  charge and a center of negative charge

16
Bond Polarity and Dipoles

• Not all molecule with polar bonds have dipole
  moments!
• Bond polarities cancel each other in molecules
  with symmetrical dipoles.
• Molecules with equal, opposing dipoles.
• See 8.2 on page 348
• Dog walking example!

17
Compound Formation

• Atoms gain, lose, or share enough electrons to
  achieve the same stable electron configuration as
  a noble gas
• Nonmetals share electrons
• Form molecules with covalent bonds
• Representative metals lose electrons to nonmetals
  in ionic compounds
• Ions are isoelectronic to noble gases

18
8.4 Ion Formation

• Binary ionic compounds
• The metal loses electron(s) to a nonmetal
• Focus on represenative metals
• The atoms lose/gain enough electrons to obtain a
  noble gas electron configuration.

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Cations

• Group IA metals form ions with a _____ charge.
• Na atom
• Na ion
• Isoelectronic to ____________________

20
Anions

• Group VIA elements form ions with a ______
  charge.
• Sulfur atom
• Sulfur ion (called _________________)
• Isoelectronic to ____________________

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Ionic Compound

• Consider the compound formed between sodium and
  sulfur.
• Each sodium atom loses 1 electron.
• Each sulfur atom needs 2 electrons.
• Formula for compound

22
Ion Size

• Cations are smaller than their parent atom.
• Atoms lose their valence shell when the ion
  forms.
• Na 1s22s22p63s1 ____ protons
• Na 1s22s22p6 ____ protons

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Ion Size

• Anions are larger than their parent atom.
• Atoms add electrons to their valence shell when
  the ion forms proton remains the same.
• F 1s22s22p5 ____ protons
• F1- 1s22s22p6 ____ protons

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Ion Size

• The diagram on page 352 should make sense.
• Isoelectronic ions decrease in size as the number
  of protons increases.
• Example ions with 10 electrons
• 10 e O2- F1- Na1 Mg2 Al3
• p 8 9 11 12 13

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Isoelectronic Ions

• 10 e O2- F1- Na1 Mg2 Al3
• p 8 9 11 12
  13
• Radius 140 136 95 65 50
• picometers

26
8.5 Energy in Binary Ionic Compounds

• Lattice energy change in energy when separated
  gaseous ions form an ionic solid.
• M(g) X-(g) ? MX(s)
  LE lt 0

27
Lattice Energy

• LE k (Q1Q2)/r
• K is the proportionality constant
• Q1 and Q2 are the charges on the ions
• r is the ionic radius

28
Lattice Energy

• LE becomes more exothermic as the ion charges
  increase and the ion radius decreases.
• Small highly charged ions have more exothermic LE
• See board for examples.

29
Formation of ionic compounds.

• Consider energy changes associated with formation
  of a binary ionic compound.
• 5 step process, page 354/355
• Most common series of steps is shown on the next
  slide.

30
Formation of ionic compounds.

• Sublime the metal.
• Ionize the gaseous metal atoms.
• 1st ionization energy.
• Dissociate the nonmetal (if diatomic).
• Bond energy
• Ionize the gaseous nonmetal atoms.
• Electron affinity
• Form the solid from the gaseous ions
• LE

31
Formation of NaF

• See page 355

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8.6 Partial Ionic Character

• When atoms with different EN bond the result is
  either a polar covalent or an ionic bond.
• Theres evidence that some level of electron
  sharing occurs in all bonds.
• Even in what we consider as ionic bonds.

33
8.6 Partial Ionic Character

• Classify a bond as ionic if it conducts
  electricity when melted.
• Essentially all compounds with metals meet this
  criteria.
• These compounds generally have more than 50
  ionic character.

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8.7 Models

• Read Fundamental Properties of Models on page
  350.

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8.8 Covalent Bond Energies

• Strength of a given bond depends upon the
  compound.
• Not all C-H bonds are of the same energy!
• See page 350.
• Bond energies given in tables are averages based
  on experimental data.

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Bond Energies

• Consider the bond energies on page 352.
• Compare the bond energies and bond length
  associated with single, double, and triple bonds
  between a given pair of atoms.

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8.8 Using Bond Energies

• The DH for a reaction can be estimated from bond
  energies.
• DH energy needed to break bonds of reactants
  
• energy released when product bonds
  form

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8.8 Using Bond Energies

• Estimate the DH for 54a on page 384.
• Expected answer - 158 kJ

39
8.9 LE Bonding Model

• Localized electron bonding model
• Assumes a molecule is made of atoms bound
  together by sharing pairs of electrons using the
  orbitals of the bonding atoms.

40
8.9 LE Bonding Model

• Localized electron bonding model
• Shared electrons are pictures to be localized in
  the space between the atoms
• Called bonding pairs
• Non-bonding valence electrons are pictured to be
  localized on the parent atom.
• Called lone pairs
• Consider HCl

41
8.10 Lewis Structures

• Lewis structures show the arrangement of the
  valence electrons in molecules (and ions).
• Representative atoms will have the same number of
  valence electrons as one of the noble gases
• 2 electrons to be like H
• 8 electrons to be like all other noble gases

42
Lewis Structures

• Lewis structures illustrate LE bonding model.
• Show the bonding electrons and the lone pairs.
• Lewis structures can be used to predict the 3D
  geometry of a molecule.
• Requires application of VSEPR Theory
• More to come on this..

43
1st Goal To Write Lewis Structures

• Sum the valence electrons.
• Use a pair of electrons to form a bond between
  each of the bonded atoms.
• Put the atom that needs the most electrons in the
  center when the molecule contains more than 2
  atoms.
• Arrange the remaining electrons to satisfy the
  duet rule for H and the octet rule for elements
  in the 2nd row of elements.

44
Writing Lewis Structures

• Practice!
• H2O
• O2
• HCN
• NO31-
• PH3

45
8.12 Resonance

• More than one valid Lewis structure can often be
  drawn for molecules with multiple bonds (double,
  triple..)
• Consider NO21-
• 2 valid Lewis structures can be drawn.

46
Resonance Structures

• Lewis structure just drawn indicate 2 types of
  bonds in NO21- -- single bond and a double bond
• However.the data shows that both bonds in NO21-
  are of the same energy and bond length
• Both bonds are stronger and shorter than a single
  bond, but not as strong or short as a double bond!

47
Exceptions to the Octet Rule

• Less than an octet.
• Be and B
• More than an octet
• 3rd period elements and up
• Odd number of electrons

48
Exceptions to the Octet Rule

• Less than an octet.
• Be - satisfied/stable with 4 electrons
• B - satisfied/stable with 6 electrons

49
Exceptions to the Octet Rule

• 2. More than an octet
• Atoms in the 3rd period and up can use their
  unfilled d orbitals to accommodate more than 8
  electrons
• Commonly see 10 electrons and 12 electrons around
  the central atom.
• Up refers to periods 4, 5,6,

50
More than an octet

• ICl3
• PF5

51
Exceptions to the Octet Rule

• Odd number of electrons
• A small number of molecules have an odd number of
  electrons
• Called free radicals
• Molecules are highly reactive/unstable
• steal an electron from other molecules
• Example NO

52
VSEPR Theory

• Valence Shell Electron Pair Repulsion Theory
• Structure around a given atom is determined by
  minimizing e-pair repulsions
• Atoms arrange themselves in 3D space in a manner
  that minimizes electron pair repulsive forces

53
VSEPR Theory

• Predictions based on VSEPR theory agree closely
  with experimental data.
• CO2
• BF3
• SO2

54
Still to Come

• VSEPR animation
• Applying LE and VSEPR Theory
• From Lewis structure to electronic vs. molecular
  geometry and bond angles
• From molecular geometry and bond polarity to
  molecular polarity