Covalent Bonding

1
Covalent Bonding

• Sometimes two atoms that both need to gain
  valence electrons to become stable have similar
  attraction for electrons.
• Sharing of electrons is another way that these
  atoms can acquire the electron configuration of
  noble gases.
• The octet rule states that atoms lose, gain or
  share electrons to achieve a stable configuration
  of eight valence electrons or an octet.

2
What Is a Covalent Bond?

• A covalent bond is a chemical bond that results
  from the sharing of valence electrons.
• In a covalent bond, the shared electrons are
  considered to be part of the complete outer
  energy level of both atoms involved.
• Covalent bonding occurs generally between atoms
  that are in close proximity to one another on the
  periodic table.
• Most covalent bonds occur in nonmetals.
• A molecule is formed when two or more atoms bond
  covalently.

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Single Covalent Bonds

• When a single set pair of electrons are shared, a
  single covalent bond is formed.
• A single covalent bond can be shown as a pair of
  dots between the symbols of the adjoining atoms,
  or by a single line connecting the symbols, which
  is called the Lewis structure.
• For example when 2 atoms of hydrogen are
  covalently bonded to one another H2 would be
  shown as.
• H H HH or HH.

4
Lewis Structures

• Lewis structures use electron dot diagrams to
  show how electrons are arranged in molecules.
• For example
• H
• ?
  ?
• H? H? H? H? ?C? ? HCH
• ? ?
• H

5
The Sigma Bond

• Single covalent bonds are also called sigma bonds
  (s).
• A sigma bond results if the atomic orbitals
  overlap end to end, concentration the electrons
  in a bonding orbital between the two atoms.
• A bonding orbital is a localized region where
  bonding electrons will most likely be found.

6
Multiple Covalent Bonds

• Atoms can gain noble gas configuration by sharing
  more than one pair of electrons.
• Double and triple covalent bonds are examples of
  multiple bonds.
• A double bond is when two pairs of electrons are
  shared.
• A triple bond is when three pairs of electrons
  are shared between two atoms.

7
The pi Bond

• A pi bond,(p), is formed when parallel orbitals
  overlap to share electrons.
• A multiple bond consists of one sigma bond and at
  least one pi bond.
• A double covalent bond has one sigma bond and one
  pi bond.
• A triple covalent bond consists of one sigma bond
  and two pi bonds.
• A pi bond always accompanies a sigma bond when
  forming double and triple bonds.

8
Strength of Covalent Bonds

• The strength of a covalent depends on how much
  distance separates bonded nuclei.
• The distance between two bonding nuclei at the
  position of maximum attraction is called bond
  length.
• The shorter the bond length, the stronger the
  bond.
• Triple bonds have the shortest bond length, then
  double bonds and single bonds have the longest
  bond length.

9
Strength of Covalent Bonds

• The amount of energy required to break a specific
  covalent bond is called bond dissociation energy.
  
• Bond dissociation energy is always a positive
  value, because the breaking of bonds always
  requires the addition of energy.
• Bond dissociation energy indicates the strength
  of a chemical bond.
• The shorter the bond length the more energy it
  will require to break the bond that holds the
  molecule together.

10
Strength of Covalent Bonds

• In chemical reactions bonds in reactant molecules
  are broken and new bonds are formed as product
  molecules are formed.
• Endothermic reactions occur when a greater amount
  of energy is required to break the existing bon
  bonds in the reactants than is released when the
  new bonds form in the product molecule.
• Exothermic reactions occur when more energy is
  released forming new bonds than is required to
  break bonds in the initial reactants.

11
Naming Molecules

• Binary molecular compounds are composed of two
  different nonmetals and do not contain metals, or
  ions.
• These compounds have common names used by the
  general public plus they have scientific names
  given to them by scientists.

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Rules for naming binary compounds.

• The first element in the formula is always named
  first, using the entire element name.
• The second element in the formula is named using
  the root of the element and adding the suffix
  ide.
• Prefixes are used to indicate the number of atoms
  of each type that are present in the compound.

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Prefixes in covalent compounds

• of atoms Prefix
• 6 hexa-
• 7 hepta-
• 8 octa-
• 9 nona-
• 10 deca

• of atoms Prefix
• 1 mono-
• 2 di-
• 3 tri-
• 4 tetra-
• 5 penta-

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Common names of some molecular compounds.

• Formula Common Name Molecular Compound Name
• H2O Water dihydrogen monoxide
• NH3 ammonia nitrogen trihydride
• N2H4 hydrazine dinitrogen tetrahydride
• N2O nitrous oxide dinitrogen monoxide
• NO nitric oxide nitrogen monoxide

15
Naming Acids

• Water solutions of some molecules are acidic and
  are named as acids.
• If the compound produces hydrogen ions in
  solution, it is an acid.
• Two common type of acids exist, binary acids, and
  oxyacids.

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Naming Binary Acids

• A binary acid contains hydrogen and one other
  element.
• When naming a binary acid, use the prefix hydro
  to name the hydrogen part of the compound.
• The rest of the name consists of a form of the
  root of the second element plus the suffix ic,
  followed by the word acid.
• For example HBr would be called hydrobromic acid.
  

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Naming Binary Acids Continued

• The term binary indicates exactly two elements,
  but there are a few acids that contain more than
  two elements that are named according to the
  rules for naming binary compounds.
• If no oxygen is present in the formula for the
  acidic compound, the acid is named in the same
  way as a binary acid, except that the root of the
  second part of the name is the root of the
  polyatomic ion that the acid contains.
• For example, HCN, which is composed of hydrogen
  and the cyanide ion is called hydrocyanic acid.

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Naming Oxyacids

• Another set of rules is used to name an acid that
  contains an oxyanion.
• An oxyanion is a polyatomic ion that contains
  oxygen.
• Any acid that contains hydrogen and an oxyanion
  is referred to as an oxyacid.

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Naming Oxyacids Continued

• Since the name of an oxyacid depends on the
  oxyanion present in the acid, you must first
  identify the anion present.
• The name of an oxyacid consists of a form of the
  root of the anion, a suffix, and the word acid.
• If the anion suffix is ate, it is replaced with
  the suffix ic.
• When the anion suffix is ite, it is replaced
  with ous.

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Naming Oxyacids Continued

• The oxyacid HNO3 has nitrate (NO3-) for the
  oxyanion. So it is named nitric acid.
• The anion for HNO2 has nitrite ion (NO2-). So it
  is named nitrous acid.
• Notice that the hydrogen in an oxyacid is not
  part of the name.

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Writing Formulas from Names

• The name of any binary molecule allows you to
  write the correct formula with ease.
• Subscripts are determined from the prefixes used
  in the name because the name indicates the exact
  number of each atom present in the molecule.
• The formula for an acid can be derived from the
  name as well.

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Molecular Structures

• The structural formula uses letter symbols and
  bonds to show relative positions of atoms.
• This is one of the most useful molecular models.
• The structural formula can be predicted for many
  molecules by drawing the Lewis structure.
• When drawing Lewis structures it is a good idea
  to follow a regular procedure.

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Steps for determining Lewis structures.

• Predict the location of certain atoms.
• Hydrogen is always a terminal, or end, atom.
  Because it can share only one pair of electrons,
  hydrogen can be connected to only one other atom.
  
• The atom with the least attraction for shared
  electrons in the molecule is the central atom.
  This atom is usually the one closer to the left
  of the periodic table. The central atom is
  located in the center of the molecule, and all
  other atoms become terminal atoms.
• Find the total number of valence electrons in the
  atoms in the molecule.

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Steps for determining Lewis structures continued.

• Determine the number of bonding pairs by dividing
  the number of electrons available for bonding by
  two.
• Place one bonding pair (single bond) between the
  central atom and each of the terminal atoms.

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Steps for determining Lewis structures continued.

• Subtract the number of pairs you used in step 4
  from the number of bonding pairs you determined
  in step 3. The remaining electron pairs include
  lone pairs as well as pairs used in double and
  triple bonds. Place lone pairs around each
  terminal atom bonded to the central atom to
  satisfy the octet rule. Any remaining pairs are
  assigned to the central atom.

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Steps for determining Lewis structures continued.

• 6. If the central atom is not surrounded by
  four electron pairs, it does not have an octet.
  You must convert one or two of the lone pairs on
  the terminal atoms to a double bond or a triple
  bond between the terminal atom as well as with
  the central atom. Remember that, in general,
  carbon, nitrogen, oxygen, and sulfur can form
  double or triple bonds with the same element or
  with another element.

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Drawing Lewis Structures for polyatomic ions.

• The main difference between drawing the Lewis
  structure for covalent compounds and polyatomic
  ions is finding the total number of electrons
  available for bonding.
• To find the total number of electrons available
  for binding, first find the number available in
  the atoms present in the ion.
• Then subtract the ion charge if the ion is
  positive, and add the ion charge if the ion is
  negative.

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Resonance Structures

• It is possible to have more than one correct
  Lewis structure when a molecule or polyatomic ion
  has both a double covalent bond and a single
  covalent bond.
• Resonance is a condition that occurs when more
  than one valid Lewis structure can be written for
  a molecule or an ion.
• The two or more correct Lewis structures that
  represent a single molecule are often called
  resonance structures.

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Resonance Structures Continued

• Resonance structures only differ in the position
  of the electron pairs, never the atom positions.
• Each actual molecule or ion that undergoes
  resonance behaves as it has only one structure.
• Meaning that no matter which structure it is in
  it will have the same properties as the other
  structures.

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Exceptions to the Octet Rule.

• Some molecules and ions do not obey the octet
  rule.
• Three reasons exist for these exceptions.
• First, a small group of molecules has an odd
  number of valence electrons and cannot form an
  octet around each atom.
• For example NO2 has five valence electrons from
  nitrogen and 12 from oxygen, totaling 17, which
  cannot form an exact number of electron pairs.

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Exceptions to the Octet Rule Cont.

• Second, some compounds form with fewer than eight
  electrons present around the atom.
• This group is relatively rare.
• BH3 is an example.
• Boron forms three covalent bonds with other
  nonmetallic atoms.
• So a total of 6 valence electrons can be shared.

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Exceptions to the Octet Rule Cont.

• When one atom donates a pair of electrons to be
  shared with an atom or ion that needs two
  electrons to become stable, a coordinate covalent
  bond forms.

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Exceptions to the Octet Rule Cont

• The third group of compounds that does not follow
  the octet rule has central atoms that contain
  more than eight valence electrons.
• This arrangement is called an expanded octet.
• When you draw the Lewis structure for these
  compounds, extra lone pairs are added to the
  central atom or more than four bonding atoms are
  present in the molecule.
• See the example problem on page 257 for an
  example.

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Molecular Shape

• The shape of a molecule determines whether or not
  molecules can get close enough to react.
• The model used to determine the molecular shape
  is referred to as the valence shell electron
  repulsion model. (VSPER model).
• The VSPER model is based on an arrangement that
  minimizes the repulsion of shared and unshared
  electrons around the central atom.

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VSPER Model Continued

• The repulsion among electron pairs in a molecule
  result in atoms existing at fixed angles to each
  other.
• The angle formed between any two terminal atoms
  and the central atom is a bond angle.

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VSPER RULES

• Draw the Lewis structure for the molecule or ion.
  
• Count the total number of regions of high
  electron density (bonding and unshared electron
  pairs) around the central atom.
• Double and triple bonds count as ONE REGION OF
  HIGH ELECTRON DENSITY.
• An unpaired electron counts as ONE REGION OF HIGH
  ELECTRON DENSITY.
• For molecules or ions that have resonance
  structures, you may use any one of the resonance
  structures.

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VSPER RULES Continued.

• Identify the most stable arrangement of the
  regions of high electron density as ONE of the
  following
• linear
• trigonal planar
• tetrahedral
• trigonal bipyramidal
• octahedral
• See page 260 for pictures of these arrangements

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Listed below are the best arrangements of the
regions of high electron repulsion.
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VSPER RULES Continued.

• Determine the positions of the atoms based on the
  types of electron pairs present (i.E., Bonding
  pairs vs. Unshared pairs). For trigonal
  bipyramidal and octahedral arrangements, there
  can sometimes be more than one possible
  arrangement of the bonding and unshared pairs
• Trigonal bipyramidal - place any unshared pairs
  in the plane of the triangle.
• Octahedral - if you have two unshared pairs,
  place them on opposite sides of the central atom.
  
• Identify the molecular structure based on the
  positions of the ATOMS (NOT on the regions of
  high electron density).

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