Achievement Standard 2.4

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Achievement Standard 2.4

• COLVALENT BONDING
• LEWIS STRUCTURES AND
• SHAPES OF MOLECULES

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LEWIS STRUCTURES

• Developed from observation that noble gases have
  a stable electron configuration.
• Atoms that share electrons form an electron-pair
  bond that helps to attain a stable noble-gas
  structure.
• Octet Rule The principle that bonded atoms
  (except H) tend to have a share in eight valance
  electrons.

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LEWIS STRUCTURES

• In writing Lewis structures, only the valance
  electrons are used.
• There are two kinds of electron pairs
• Shared electrons form covalent bonds (indicated
  by a line)
• Unshared pairs of electrons (indicated by two
  dots)

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WRITING LEWIS STRUCTURES

• Count the total number of valance electrons
  available (A) by first adding the group numbers
  of atoms present. (For group numbers with two
  digits, use last digit only!!)
• For polyatomic anion also add a value equal to
  charge to number of electrons.
• For polyatomic cation also subtract a value
  equal to charge from number of electrons.

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WRITING LEWIS STRUCTURES

• Calculate the total number of electrons needed
  (N) for each atom to have its own noble gas
  structure (two for H and eight for all atoms C
  and beyond).
• Subtract the number in step 1 from the number in
  step 2. This represents the number of shared, or
  bonding electrons (S)
• S N - A

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WRITING LEWIS STRUCTURES

• To predict arrangement of atoms in molecules and
  ions, use the following as a guide
• Many molecules and ions will have a central atom
  and two or more terminal atoms bonded to it.
• Hydrogen is almost always a terminal atom oxygen
  and the halogens are often terminal atoms.
• The central atom is often the first atom
  presented in a formula (e.g., S in SO2 or C in
  CH4)

Note exceptions
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WRITING LEWIS STRUCTURES

• From step 3, assign two bonding electrons (shared
  pairs) to each connection between atoms in the
  molecule or ion. The remaining electrons
  represent a pool of electrons you use for any
  lone pairs and/or to make multiple bonds, if
  necessary.

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WRITING LEWIS STRUCTURES

• Place lone pairs about each terminal atom (except
  H) to satisfy the octet rule.
• If the central atom is not yet surrounded by four
  electron pairs, convert one or more terminal atom
  lone pairs to make multiple bonds between central
  and terminal atoms. In general, C, N, O and S
  have a tendency to form multiple bonds (double or
  triple).

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EXCEPTIONS TO OCTET RULE

• Electron-deficient molecules
• Odd electron species (NO or NO2)
• Central atom with only two or three bonds (BeF2
  or BF3)
• Expanded octets where central atom is
  surrounded by more than 4 pairs of valance
  electrons (PCl5 or SF6 or XeF4)
  Note this only occurs in Level 3 examples

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SHAPES OF MOLECULES

• Molecular geometry refers to the three
  dimensional shape of molecules.
• Shapes can be predicted based on electron
  repulsion.
• VSEPR or valence shell electron pair repulsion
  theory helps to predict molecular geometry
• In essence, VSEPR says that the electron pairs
  surrounding an atom repel one another.
  Therefore, the electron pairs are oriented to be
  as far apart as possible.

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VSEPR THEORY

• Consider a central atom (A) bonded with terminal
  atoms (X), with no unshared electron pairs. The
  common species with 2 to 6 electron pairs and
  their geometries would be
• AX2 180o linear
• AX3 120o triangular planer
• AX4 109.5o tetrahedron
• And for Level 3 shapes
• AX5 90o, 120o, 180o triangular bipyramid
• AX6 90o, 180o octahedron

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VSEPR THEORY

• Now consider what happens when the central atom
  (A) has both terminal atoms (X) and unshared
  pairs of electrons (E)
• AX2 180o linear
• AX3 120o triangular planar
• AX2E 120o bent (angular)
• AX4 109.5o tetrahedral
• AX2E2 109.5o bent
• And the Level 3 example
• AX3E 109.5o triangular pyramid

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VESPR THEORY

• VSEPR theory also applies to
• Species with double and triple bonds multiple
  bonds are treated like single bonds in terms of
  shape. (e.g. carbon dioxide, dinitrogen oxide)
• Species with no single central atom. (e.g.
  Acetylene, ethylene)

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POLARITY BONDS

• Nonpolar A symmetrical distribution of
  electrons between atoms. Nonpolar bonds form
  whenever two atoms joined are identical. (e.g.
  H2)
• Polar An unsymmetrical distribution of
  electrons between atoms. Polar bonds form
  whenever two atoms joined are different. (e.g.
  HCl)

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POLARITY MOLECULES

• A polar molecule is one that contains positive
  and negative poles (a dipole).
• A dipole for a simple molecule such as HF can be
  illustrated as H F
• A nonpolar molecule is one that contains no
  positive and negative poles.

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POLARITY PREDICTION

• For simple molecules with two atoms, determining
  polarity is easy
• All diatomic molecules are nonpolar.
• Molecules with two different atoms are polar.

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POLARITY PREDICTION

• To determine polarity for molecules with more
  than two atoms you need to know bond polarity and
  molecular shape. (BeF2, H2O and CCl4).
• If the polar AX bonds in a molecule are arranged
  symmetrically in 3D (has 2 or more planes of
  symmetry) around the central atom A, the molecule
  is nonpolar.
• Polarity can also be determined by drawing the
  molecule in the shape of a , using vectors for
  bonds. If the vectors add to zero, then the
  molecule is nonpolar.