Chemical Bonding I: The Covalent Bond

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Chemical Bonding I The Covalent Bond

• Lewis dot symbols
• The covalent bond
• Electronegativity
• Writing Lewis structures
• Formal charge and Lewis structure
• The concept of resonance
• Exceptions to the octet rule
• Strength of the covalent bond

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Lewis dot symbols

• Lewis dot symbol
• By American chemist Gilbert Lewis
• Consists of the symbol of an element and one dot
  for each valence electron in an atom of the
  element.

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• S Ne3s23p4
• The dots (representing electrons) are placed on
  the four sides of the atomic symbol (top, bottom,
  left, right)
• Each side can accommodate up to 2 electrons
• The number of valence electrons in the table
  below is the same as the column number of the
  element in the periodic table (for representative
  elements only)

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The Covalent Bond

• The covalent bond
• results from the sharing of electrons between two
  atoms.
• typically involves one nonmetallic element with
  another

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• The diatomic hydrogen molecule (H2) is the
  simplest model of a covalent bond, and is
  represented in Lewis structures
• The shared pair of electrons provides each
  hydrogen atom with two electrons in its valence
  shell (the 1s) orbital.

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• Covalent bonding between many-electron atoms
  involves only covalent electrons.
• Nonbonding electrons (lone pairs)
• The valence electrons that are not involved in
  covalent bond formation.

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• The structures we use to represent H2and Cl2 is
  called Lewis structures.
• Lewis Structures
• Is a representations of covalent bonding using
  Lewis dot symbols in which shared electron pairs
  are shown either as lines or as pairs of dots
  between two atoms, and lone pairs are shown as
  pairs of dots on individual atoms.

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• Octet Rule
• An atom other than hydrogen tends to form bonds
  until it is surrounded by eight valence
  electrons.

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• Single bond
• Two atoms held together by one electron pair are
  said to be joined by single bond.
• Multiple bonds
• In many molecules atoms attain complete octets by
  sharing more than one pair of electrons between
  them
• Two electron pairs shared a double bond
• Three electron pairs shared a triple bond

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• Bond length
• The distance between the nuclei of two bonded
  atoms in a molecule.
• Multiple bonds are shorter than single covalent
  bonds.

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Bond Polarity and Electronegativity

• The electron pairs shared between two atoms are
  not necessarily shared equally
• Two extreme examples
• In Cl2 the shared electron pairs is shared
  equally
• In NaCl the 3s electron is stripped from the Na
  atom and is incorporated into the electronic
  structure of the Cl atom - and the compound is
  most accurately described as consisting of
  individual Na and Cl- ions
• For most covalent substances, their bond
  character falls between these two extremes

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• Bond polarity is a useful concept for describing
  the sharing of electrons between atoms
• A nonpolar covalent bond is one in which the
  electrons are shared equally between two atoms
• A polar covalent bond is one in which one atom
  has a greater attraction for the electrons than
  the other atom. If this relative attraction is
  great enough, then the bond is an ionic bond

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• Electronegativity
• A quantity termed 'electronegativity' is used to
  determine whether a given bond will be nonpolar
  covalent, polar covalent, or ionic.
• Electronegativity is defined as the ability of an
  atom in a particular molecule to attract
  electrons to itself

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Drawing Lewis Structures

• The general procedure
• 1. Sum the valence electrons from all atoms
• Use the periodic table for reference
• Add an electron for each indicated negative
  charge, subtract an electron for each indicated
  positive charge
• 2. Write the symbols for the atoms to show which
  atoms are attached to which, and connect them
  with a single bond
• You may need some additional evidence to decide
  bonding interactions
• If a central atom has various groups bonded to
  it, it is usually listed first CO32-, SF4
• Often atoms are written in the order of their
  connections HCN

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• 3.Complete the octets of the atoms bonded to the
  central atom (H only has two)
• 4. Place any leftover electrons on the central
  atom (even if it results in more than an octet
• 5. If there are not enough electrons to give the
  central atom an octet, try multiple bonds (use
  one or more of the unshared pairs of electrons on
  the atoms bonded to the central atom to form
  double or triple bonds

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• Draw the Lewis structure of phosphorous
  trichloride (PCl3)
• 1. We will have 5(P) plus 21 (37, for Cl), or 26
  total valence electrons
• 2. The general symbol, starting with only single
  bonds, would be

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• 3. Completing the octets of the Cl atoms bonded
  to the central P atom
• 4. This gives us a total of (18 electrons) plus
  the 6 in the three single bonds, or 24 electrons
  total. Thus we have 2 extra valence electrons
  which are not accounted for. We will place them
  on the central element

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• 5. The central atom now has an octect, and there
  is no need to invoke any double or triple bonds
  to achieve an octet for the central atom. We are
  finished.

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Formal Charge and Lewis Structure

• Formal charge
• The difference between the valence electrons in
  an isolated atom and the number of electrons
  assigned to that atom in a Lewis structure is
  called formal charge
• Two examples
• O3 ( on textbook, P 258)
• CO2

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• Example Carbon Dioxide (CO2)
• Carbon has 4 valence electrons. Each oxygen has 6
  valence electrons, therefore our Lewis structure
  of CO2 will have 16 electrons.
• Both structures fulfill the octet rule.
• Which one is true?
• Use formal charge to determine

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• Which structure is correct? In general, when
  several Lewis structures can be drawn the most
  stable structure is the one in which
• The formal charges are the smallest
• Any negative charge is found on the most
  electronegative atom

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• In the above case, the second structure is the
  one with the smallest formal charges (i.e. 0 on
  all the atoms).
• Furthermore, in the first possible Lewis
  structure the carbon has a formal charge of 0 and
  one of the oxygens it is bonded to has a formal
  charge of 1.
• Oxygen is more electronegative than Carbon, so
  this situation would seem unlikely

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The Concept of Resonance

• Resonance structure
• Is one of two or more Lewis structures for a
  single molecule that cannot be described fully
  with only one Lewis structure.
• Resonance
• Means the use of two or more Lewis structures to
  represent a particular molecule.

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• This indicates that the ozone molecule is
  described by an average of the two Lewis
  structures (i.e. the resonance forms)

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• The important points to remember about resonance
  forms are
• The molecule is not rapidly oscillating between
  different discrete forms
• There is only one form of the ozone molecule, and
  the bond lengths between the oxygens are
  intermediate between characteristic single and
  double bond lengths between a pair of oxygens
• We draw two Lewis structures (in this case)
  because a single structure is insufficient to
  describe the real structure

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Exceptions to the Octet Rule

• There are three general ways in which the octet
  rule breaks down
• 1. Molecules with an odd number of electrons(The
  Odd-Electron Molecules)
• 2. Molecules in which an atom has less than an
  octet (The Incomplete Octet)
• 3. Molecules in which an atom has more than an
  octet( The Expanded Octet)

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• Less than an octet (most often encountered with
  elements of Boron and Beryllium)
• It has 6 electrons
• the structure of BF3, with single bonds, and 6
  valence electrons around the central boron is the
  most likely structure

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• Odd number of electrons
• Total electrons 6511
• There are currently 5 valence electrons around
  the nitrogen.
• We appear unable to get an octet around each atom
  

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• More than an octet (most common example of
  exceptions to the octet rule)
• The orbital diagram for the valence shell of
  phosphorous is
• Third period elements occasionally exceed the
  octet rule by using their empty d orbitals to
  accommodate additional electrons

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Strengths of Covalent Bonds

• Bond-dissociation energy (i.e. "bond energy")
• Bond energy is the enthalpy change (DH, heat
  input) required to break a bond (in 1 mole of a
  gaseous substance)