Bonding

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Chapter 8

• Bonding

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Chemical Bonds, Lewis Symbols, and the Octet Rule

• Chemical bond attractive force holding two or
  more atoms together.
• Covalent bond results from sharing electrons
  between the atoms. Usually found between
  nonmetals.
• Ionic bond results from the transfer of electrons
  from a metal to a nonmetal.
• Metallic bond attractive force holding pure
  metals together.

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Chemical Bonds, Lewis Symbols, and the Octet Rule

• Lewis Symbols
• As a pictorial understanding of where the
  electrons are in an atom, we represent the
  electrons as dots around the symbol for the
  element.
• The number of electrons available for bonding are
  indicated by unpaired dots.
• These symbols are called Lewis symbols.
• We generally place the electrons one four sides
  of a square around the element symbol.

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Lewis Dot Structure
Lewis wrote in a memorandum dated March 28, 1902
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Chemical Bonds, Lewis Symbols, and the Octet Rule
Lewis Symbols
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Chemical Bonds, Lewis Symbols, and the Octet Rule

• The Octet Rule
• All noble gases except He has an s2p6
  configuration.
• Octet rule atoms tend to gain, lose, or share
  electrons until they are surrounded by 8 valence
  electrons (4 electron pairs).
• Caution there are many exceptions to the octet
  rule.

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Ionic Bonding
Consider the reaction between sodium and
chlorine Na(s) ½Cl2(g) ? NaCl(s) DHºf -410.9
kJ
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Ionic Bonding

• An atom with a low ionization energy reacts with
  an atom with high electron affinity.
• The electron moves.
• Opposite charges hold the atoms together.

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Ionic Bonding

• Energetics of Ionic Bond Formation
• Lattice energy the energy required to completely
  separate an ionic solid into its gaseous ions.
• Example NaCl(s) ? Na(g) Cl-(g)
• Lattice energy (El) depends on the charges on the
  ions and the sizes of the ions
• k is a constant (8.99 x 10 9 Jm/C2), Q1 and Q2
  are the charges on the ions, and d is the
  distance between ions.

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Ionic Bonding

• Energetics of Ionic Bond Formation
• Lattice energy increases as
• The charges on the ions increase
• The distance between the ions decreases.

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El drops as negative ion gets bigger!!
El large due to 2 and -2 Charges gets smaller
as ions get larger.
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Ionic Bonding

• Transition Metal Ions
• Lattice energies compensate for the loss of up to
  three electrons.
• In general, electrons are removed from orbitals
  in order of decreasing n (i.e. electrons are
  removed from 4s before the 3d).
• Polyatomic Ions
• Polyatomic ions are formed when there is an
  overall charge on a compound containing covalent
  bonds. E.g. SO42-, NO3-.

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Formation of Ionic Bonds

• Formation of ionic crystal lattice is determined
  by
• Ionization energies of the atoms involved
  (increase with electrons removed)
• Lattice energy of the final crystal
• Rarely do we see ionic compounds with greater
  than 3 (indicating 3 e- removed)

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NaCl
This is the formation of an ionic bond.
-

Na
Cl
electron transfer
and the formation of ions
Cl2
This is the formation of a covalent bond.
sharing of a pair of electrons
and the formation of molecules
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Covalent Bonding

• When two similar atoms bond, none of them wants
  to lose or gain an electron to form an octet.
• When similar atoms bond, they share pairs of
  electrons to each obtain an octet.
• Each pair of shared electrons constitutes one
  chemical bond.
• Example H H ? H2 has electrons on a line
  connecting the two H nuclei.

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Covalent Bonding
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Energy
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Internuclear Distance
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Energy
0
Internuclear Distance
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Energy
0
Internuclear Distance
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Energy
0
Internuclear Distance
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Energy
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Bond Length
Internuclear Distance
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Energy
Bond Energy
0
Internuclear Distance
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Bond Polarity and Electronegativity

• In a covalent bond, electrons are shared.
• Sharing of electrons to form a covalent bond does
  not imply equal sharing of those electrons.
• There are some covalent bonds in which the
  electrons are located closer to one atom than the
  other.
• Unequal sharing of electrons results in polar
  bonds.
• So there is actually a continuum between purely
  covalent bonds (as in H2) and a purely ionic bond
  (KF)
• How can we quantify the distribution of
  electrons? With electronegativity

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Bond Polarity and Electronegativity

• Electronegativity
• Electronegativity The ability of one atoms in a
  molecule to attract electrons to itself.
• Linus Pauling set electronegativities on a scale
  from 0.7 (Cs) to 4.0 (F). Hydrogen was
  arbitrarily set at 2.1
• Electronegativity Trends
• Increases across a period and
• Decreases down a group.

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Bond Polarity and Electronegativity
Electronegativity
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Ionic
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Ionic Character
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Polar Covalent
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Covalent
2.0
1.0
3.0
Electronegativity difference
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Bond Polarity and Electronegativity

• Dipole Moments
• Consider HF
• The difference in electronegativity leads to a
  polar bond.
• There is more electron density on F than on H.
• Since there are two different ends of the
  molecule, we call HF a dipole.
• Dipole moment, m, is the magnitude of the dipole
• where Q is the magnitude of the charges (in
  Coulombs)
• Dipole moments are measured in debyes, D.

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How It is drawn
or
0.41 0.41- H F
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Computer simulation of electron density
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Calculation of dipole moment

• Measured value? How?
• By Capacitance differences. See how molecules
  affect an electric field.
• Calculate dipole moment given Q and r
• Calculate portion of electron transferred given
  dipole moment and r

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Covalent Bonding

• Lewis Structures
• Covalent bonds can be represented by the Lewis
  symbols of the elements
• In Lewis structures, each pair of electrons in a
  bond is represented by a single line

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Covalent Bonding

• Multiple Bonds
• It is possible for more than one pair of
  electrons to be shared between two atoms
  (multiple bonds)
• One shared pair of electrons single bond (e.g.
  H2)
• Two shared pairs of electrons double bond (e.g.
  O2)
• Three shared pairs of electrons triple bond
  (e.g. N2).

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Drawing Lewis Structures

• Draw a skeleton structure. A skeleton structure
  is a rough map showing the arrangement of atoms
  within the molecule. In general, you need to
  determine the skeleton experimentally, but here
  are a few guidelines for predicting skeleton
  structures from molecular formulas.

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Drawing Lewis Structures

• Draw a Skeleton Structure (Contd)
• Central atoms are usually
• the largest atoms, or
• the least electronegative atom.
• H and the halogens are usually outside atoms.
• Don't put more than four atoms around a central
  atom unless the central atom is third period or
  lower.

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Drawing Lewis Structures

• Count total valence electrons.
• Add the number of electrons in the valence shells
  of all atoms in the molecule.
• If the molecule is charged, add an electron for
  each negative charge and subtract an electron for
  each positive charge.
• Noble gas compounds are very uncommon (except on
  general chemistry tests!) Should you encounter
  one, each noble gas atom has 8 valence electrons.

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Drawing Lewis Structures(Contd)

• Connect the structure.
• Draw a bond between the central atom and each
  outside atom.
• Each bond uses 2 valence electrons.
• Place electrons on outside atoms.
• Use remaining electrons to satisfy the octets for
  each of the outside atoms.
• If you run out of electrons at this point, the
  skeleton structure was wrong. Go back to step 1.

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Drawing Lewis Structures (Contd)

• Place all remaining electrons on the central
  atom.
• If there are more than 8 electrons on the central
  atom, and the central atom is not third period or
  lower, you counted the number of valence
  electrons incorrectly. Go back to step 2. (C, N,
  O, and F are NEVER surrounded by more than 8
  electrons!)
• If the octet on the central atom is not complete,
  try sharing lone pairs of outside atoms to form
  double or triple bonds.

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Drawing Lewis Structures (Contd)

• Place all remaining electrons on the central
  atom. (Contd)
• If you can't get an octet on the central atom, at
  this point, check to see whether the total number
  of valence electrons for this molecule is odd.
  It's impossible to give octets to all atoms in an
  odd electron molecules. Get as close to an octet
  as possible by forming multiple bonds.

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Drawing Lewis Structures (Contd) Formal Charges

• Useful if trying to decide between different
  possible structures.
• Defined as charge atom would have if all atoms
  had the same electronegativity.
• In other words, they shared the electrons equally.

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Drawing Lewis Structures (Contd) Formal Charges

• How to Calculate
• All of the unshared (nonbonding) electrons are
  assigned to the atom on which they are found.
• Half of the bonding electrons are assigned to
  each atom in the bond.
• Formal Charge valence e- in the isolated atom
  - e- assigned to it in Lewis structure.
• The sum of the formal charges in a neutral atom
  must equal zero.

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Drawing Lewis Structures (Contd) Formal Charges

• When several Lewis structures are possible, the
  most stable structure will be one in which
• The atoms bear formal charges closest to zero.
• Any negative formal charge resides on the more
  electronegative atom (usually this is oxygen)

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Apply to CO2
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Drawing Lewis Structures

• Resonance Structures
• Some molecules are not well described by Lewis
  Structures.
• Typically, structures with multiple bonds can
  have similar structures with the multiple bonds
  between different pairs of atoms
• Resonance structures are a consequence of valence
  bond theory that deals in localized electrons (ie
  in three atom species the electrons are localized
  between a pair of electrons.
• MO theory (to be covered later) avoids resonance
  difficulties

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Drawing Lewis Structures
Resonance Structures
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Drawing Lewis Structures

• Resonance Structures
• Example experimentally, ozone has two identical
  bonds whereas the Lewis Structure requires one
  single (longer) and one double bond (shorter).

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Drawing Lewis Structures

• Resonance Structures
• Example in ozone the extreme possibilities have
  one double and one single bond. The resonance
  structure has two identical bonds of intermediate
  character.
• Common examples O3, NO3-, SO42-, NO2, and
  benzene.

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Drawing Lewis Structures

• Resonance in Benzene
• Benzene consists of 6 carbon atoms in a hexagon.
  Each C atom is attached to two other C atoms and
  one hydrogen atom.
• There are alternating double and single bonds
  between the C atoms.
• Experimentally, the C-C bonds in benzene are all
  the same length.
• Experimentally, benzene is planar.

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Drawing Lewis Structures

• Resonance in Benzene
• We write resonance structures for benzene in
  which there are single bonds between each pair of
  C atoms and the 6 additional electrons are
  delocalized over the entire ring
• Benzene belongs to a category of organic
  molecules called aromatic compounds (due to their
  odor).

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Exceptions to the Octet Rule

• There are three classes of exceptions to the
  octet rule
• Molecules with an odd number of electrons
• Molecules in which one atom has less than an
  octet
• Molecules in which one atom has more than an
  octet.
• Odd Number of Electrons
• Few examples. Generally molecules such as ClO2,
  NO, and NO2 have an odd number of electrons.

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Exceptions to the Octet Rule

• Less than an Octet
• Relatively rare.
• Molecules with less than an octet are typical for
  compounds of Groups 1A, 2A, and 3A.
• Most typical example is BF3.
• Formal charges indicate that the Lewis structure
  with an incomplete octet is more important than
  the ones with double bonds.

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Exceptions to the Octet Rule

• More than an Octet
• This is the largest class of exceptions.
• Atoms from the 3rd period onwards can accommodate
  more than an octet.
• Beyond the third period, empty d-orbitals are
  available to participate in bonding and accept
  the extra electron density.