Chapter 7: Covalent Bonding

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Chapter 7 Covalent Bonding

• Chapter Outline
• 7.1 Lewis Dot Structures
• 7.2 Molecular Geometry
• 7.3 Polarity of Molecules
• 7.4 Atomic Orbitals Hybridization

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Covalent Bond in Hydrogen

• A covalent bond consists of a pair of electrons
  shared between two atoms. Can be represented as
• HH H-H
• This does not imply that the electrons are fixed
  in position between the 2 nuclei.
• A more accurate picture is to look at the
  electron density (cloud) between the two atoms.

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Bond Distance (Length)

• Graph 7.2 Energy of two hydrogen atoms as a
  function of distance
• There is an ideal distance between the hydrogen
  nuclei at which the molecule is in its most
  stable configuration.
• This distance balances the attractive forces
  between the electrons and the positively charged
  nucleus and the repulsive forces between the two
  positively charged nuclei.

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Lewis Structures The Octet Rule

• The covalent bond was first theorized by the
  American chemistry G.N. Lewis in 1916.
• He pointed out the apparent stability of the
  electron configuration of noble gases (as well as
  the noble-gas like structures of of many
  monatomic ions).
• Atoms, by sharing electrons to form a bond, can
  achieve a noble gas configuration.

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Bonding Diagram for Hydrogen
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Bonding Diagram of HF
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Lewis Dot Structures

• These structures, in which all the valence
  electrons are represented by dots, are called
  Lewis dot structures.
• In writing Lewis dot structures, be sure to only
  show the valence electrons.
• The number of valence electrons is equal to the
  last digit of the group number in the periodic
  table.

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Two Types of Valence Electrons

• In Lewis dot structures of molecules or
  polyatomic ions, valence electrons tend to occur
  in pairs
• i) A pair of electrons shared between two atoms
  is a covalent bond.
• ii) An unshared pair of electrons, residing
  entirely on one atom, is show as a pair of dots
  on that atom (i.e. HF). These electrons are
  often called lone pairs.

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Examples of Lewis Dot Structures

• Draw on transparencies

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Types of Bonds

• A single pair of electrons shared between two
  atoms is called a single bond.
• Two pairs of electros shared between the same
  atoms is called a double bond.
• Three pairs of electrons shared between the same
  atoms is called a triple bond.

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Octet Rule

• Atoms involved in a covalent bond tend to have
  noble gas configurations. This is referred to as
  the octet rule.
• Nonmetals, except hydrogen, achieve a noble gas
  configuration by sharing in an octet of electrons.

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Rules for Writing Lewis Dot Structures

• I. Count the number of valence electrons.
  Molecules sum up the valence electrons of the
  atoms present.
• Polyatomic anion one electron added for each
  negative charge.
• Polyatomic cation one electron subtracted for
  each positive charge.

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Rules (Contd)

• II. Draw a skeleton structure for the species,
  joining atoms by single bonds. Most molecules
  and ions have a central atom bonded to two or
  more atoms. The central atom is usually the one
  written first in the chemical formula.

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Rules (Contd)

• III. Determine the number of valence electrons
  still available.
• IV. Determine the number of valence electrons
  required to fill out an octet for each atom.
• a) if number of electrons available is equal to
  the number required, distribute them as lone
  pairs.

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Rules (Contd)

• b) If the number of electrons available is less
  than the number required, the skeleton structure
  must be modified by introducing double and/or
  triple bonds.
• Multiple bonds are usually limited to C, N, O,
  and S.

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Resonance Forms

• There are times when Lewis structure do not
  accurately describe the properties of the
  molecule or ion it represents.
• e.g. SO2 sulfur dioxide
• Each sulfur-oxygen bond is intermediate between a
  single and double bond. The two forms are
  referred to as resonance forms.

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Properties of Resonance Structures

• 1. Resonance forms do not imply different
  molecules there is only one type of molecule and
  its structure is intermediate between the various
  resonance forms.
• 2. Resonance is expected when you can write two
  or more Lewis structures that are reasonable.
• 3. Resonance structures differ only in electron
  distribution not in the arrangement of atoms.

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Formal Charge

• The formal charge of any atom in a Lewis
  structure can be calculated.
• The formal charge is simply the difference
  between the number of valence electrons in the
  free atom and the number assigned in the Lewis
  structure.
• Assigned electrons include lone pairs and one
  half of the bonding electrons shared by the atom.
• Formal Charge (Cf) of valence electrons -
  of unshared electrons ½ of bonding electrons

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Calculation of Formal Charge

• Do Examples on transparency.

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Exceptions to the Octet Rule Electron-Deficient
Molecules

• Species which do not follow the octet rule are
  molecules or ions which have an odd number of
  valence electrons, such as NO (nitrogen oxide) or
  NO2 (nitrogen dioxide).
• For these odd-electron species, no Lewis
  structures, in which the octet rule is obeyed,
  can be drawn.
• Other example include molecular oxygen, O2.

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Electron Deficient Atoms

• Group 2 elements typically only form 2 bonds to
  other atoms.
• Group 13 typically only form 3 bonds to other
  atoms.

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Exceptions to the Octet Rule Expanded Octets

• Two examples of molecules with a central atom
  surrounded by more than 8 electrons are PCl5 and
  SF6.
• These compounds typically involved the central
  atoms surrounded by halogen atoms or oxygen. The
  central atom is always a nonmetal found in group
  15 or 16 (S, P, As, Pb, Se, Te).
• These atoms are able to have expanded octets
  because they have empty d orbitals.

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Molecular Geometry

• Molecules containing two atoms such as Cl2, H2 or
  HCl are always linear.
• e.g. H-H Cl-Cl H-Cl
• With molecules containing 3 atoms, the angle
  between bonds (bond angle) must be considered.
  Consider the molecule YX2.

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Geometry (Contd)
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Valence-Shell Electron Pair Repulsion (VSEPR)

• First proposed by Sidgwick and Powell in 1940.
• Geometry is predicted on the principle of
  electron pair repulsion.
• The valence electron pairs surrounding an atom
  repel one another. Consequently, the orbitals
  containing those electron pairs are oriented to
  be as far apart as possible.

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Ideal Geometries for Systems in Which the Central
Atom Has 2-6 Electron Pairs

• Consider a central atom A, which is surrounded by
  2-6 electron pairs, all of which are involved in
  bonds to terminal atoms, X.
• e.g. AX2, AX3, AX4, AX5, AX6
• (refer to figure 7.4)

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Molecular Geometries

• AX2 linear
• AX3 triangular planar
• AX4 tetrahedron
• AX5 triangular bipyramid
• AX6 octahedron

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Effects of Unshared Electrons on Molecular
Geometry

• The VSEPR model can also predict the geometries
  of molecules with unpaired electrons.
• 1. Electron-pair geometry is roughly the same
  as that observed for molecules with only single
  bonds. Bond angles typically smaller.
• 2. Molecular geometry is very different when
  unpaired electrons are present. When describing
  geometry, we refer only to the positions of the
  bonded atoms.

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Examples

• Ammonia, NH3 (show on transparency)
• Water, H2O (show on transparency)

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Multiple Bonds

• Multiple bonds, for the sake of molecular
  geometry, are treated like single bonds.
• Extra electrons in double and triple bonds have
  to exist in the same region of space as single
  bonds, thus no change in geometry.
• Molecular geometry is determined solely by the
  number of terminal atoms, X, bonded to the
  central atom, irrespective of whether the bonds
  are single, double or triple and the number of
  unpaired electrons around the central atom.

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Polarity of Molecules

• Covalent compounds can exist as tow different
  types of compounds
• i) polar unsymmetrical distribution of electron
  density between atoms. There exists a positive
  and negative pole and therefore a dipole moment.
• ii) nonpolar symmetrical distribution of
  electron density. No dipole moment.

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Polar and Nonpolar Molecules

• Polar molecules have a partial positive charge
  and a partial negative charge (poles) at
  different points within the molecule.
• Polar molecules orient themselves in an electric
  field.
• Nonpolar molecules have no partial charges and
  are thus unaffected by an external electric field

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Example of Dipole Moments

• Show on transparency

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Criteria for Determining Polarity

• Polarity is dependent on bond polarity and
  molecular geometry.
• If the polar A-X bonds in a molecule AXmEn are
  arranged symmetrically around the central atom A,
  the molecule is nonpolar.

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Hybridization

• Linus Pauling developed a theoretical model of
  the covalent bond, called the atomic orbital or
  valence bond orbital model. He won the Nobel
  Prize in 1954.
• According to the model, a covalent bond contains
  two electrons of opposite spin within a orbital.
• The number of covalent bonds an atom can form is
  dependent on the number of unpaired electrons it
  has.

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Hybridization (Contd)

• Hybrid Orbitals sp, sp2, sp3, sp3d, sp3d2
• Consider the examples of BeF2 and CCl4
• Hybrid orbitals result from the mixing of two
  different orbitals (e.g. s and p).

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Hybrid Orbitals

• sp orbital (BeF2)
• 1 s atomic orbital 1 p atomic orbital two sp
  hybrid orbitals
• The number of hybrid orbitals formed is equal to
  the number of atomic orbitals mixed.
• Shared as well as unshared electrons are found in
  hybridized orbitals.

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Multiple Bonds

• Multiple bonds have no impact on molecular
  geometry. The electrons that make up multiple
  bonds (pi) are not found in hybridized orbitals.
• Do examples Ethylene and ethyne.