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Covalent Bonding

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					      Chapter 11
   Chemical Bonding
Forces that hold atoms together
      The Nature of Bonding
• There are several major types of bonds.
  Ionic, covalent and metallic bonds are the
  three most common types of bonds.
• Covalent bonds – electrons are shared
  between atoms.
• Ionic bonds – electrons are transferred
  between atoms, creating cations and anions.
• Metallic bonds – two or more metals bonded
  together.
The Nature of Covalent Bonding
• There are two different types of covalent
  bonds, polar covalent and nonpolar
  covalent.
  – polar covalent – electrons are not shared
    equally between the two bonded atoms. The
    electrons are pulled toward the more
    electronegative of the elements.
  – nonpolar covalent – electrons are shared
    equally between the two bonded atoms.
9_12
                                         Electronegativities

                                                       H
                                                       2.1
IA     IIA                                                                      IIIA   IVA   VA    VIA   VIIA

Li     Be                                                                        B      C     N    O      F
1.0    1.5                                                                      2.0    2.5   3.0   3.5   4.0

Na     Mg                                                                       Al     Si     P     S    Cl
                                                      VIIIB
0.9    1.2    IIIB     IVB   VB    VIB   VIIB                       IB    IIB
                                                                                1.5    1.8   2.1   2.5   3.0

K      Ca     Sc       Ti     V    Cr    Mn     Fe     Co     Ni    Cu    Zn    Ga     Ge    As    Se    Br
0.8    1.0    1.3      1.5   1.6   1.6   1.5    1.8    1.8    1.8   1.9   1.6   1.6    1.8   2.0   2.4   2.8

Rb     Sr      Y       Zr    Nb    Mo    Tc     Ru     Rh     Pd    Ag    Cd    In     Sn    Sb    Te     I
0.8    1.0    1.2      1.4   1.6   1.8   1.9    2.2    2.2    2.2   1.9   1.7   1.7    1.8   1.9   2.1   2.5

Cs     Ba    La–Lu     Hf    Ta    W     Re     Os     Ir     Pt    Au    Hg    Tl     Pb    Bi    Po    At
0.7    0.9   1.1–1.2   1.3   1.5   1.7   1.9    2.2    2.2    2.2   2.4   1.9   1.8    1.8   1.9   2.0   2.2

Fr     Ra    Ac–No
0.7    0.9   1.1–1.7
   The formation of a bond
between two hydrogen atoms.




  Source: Andrey K. Geim/High Field Magnet Laboratory/University of Nijmegen
Probability representations of the electron sharing
in HF. (a) What the probability map would look like
 if the two electrons in the H–F bond were shared
equally. (b) The actual situation, where the shared
 pair spends more time close to the fluorine atom
             than to the hydrogen atom.
The Nature of Covalent Bonding
• Ionic bonds are formed when there is an
  electronegativity difference (DEN) greater
  than 2.0.
• Polar covalent bonds form when there is a
  DEN between 0.5 and 1.7.
• Nonpolar covalent bonds form when there
  is a DEN between 0 and 0.49.
The Nature of Covalent Bonding
• If the DEN is between 1.7 and 2.0, an ionic
  bond will form if a metal is one of the
  elements, and a polar covalent bond will
  form if only nonmetals or metalloids are
  present.
The Nature of Covalent Bonding
• What type of bond is formed between the
  following elements?

• N and O               K and F

• Mg and Cl             P and F

• C and H
             Bond Polarity
• Covalent bonding between unlike atoms
  results in unequal sharing of the
  electrons
  – One end of the bond has larger electron
    density than the other
• The result is bond polarity  H • F 
                                   •
  – The end with the larger electron density
    gets a partial negative charge
  – The end that is electron deficient gets a
    partial positive charge
           The three possible types of bonds:
(a) a covalent bond formed between identical atoms;
    (b) a polar covalent bond, with both ionic and
               covalent components; and
      (c) an ionic bond, with no electron sharing.
             Dipole Moment
• Bond polarity results in an unequal electron
  distribution, resulting in areas of partial
  positive and partial negative charge

• Any molecule that has a center of positive
  charge and a center of negative charge in
  different points is said to have a dipole
  moment (two different poles of charge).
            Dipole Moment
• If a molecule has more than one polar
  covalent bond, the areas of partial negative
  and positive charge for each bond will
  partially add to or cancel out each other
• The end result will be a molecule with one
  center of positive charge and one center of
  negative charge
• The dipole moment effects the attractive
  forces between molecules and therefore the
  physical properties of the substance
(a) The charge
distribution in
the water
molecule.
(b) The water
molecule
behaves as if it
had a positive
end and a
negative end, as
indicated by the
arrow.
(a) Polar water molecules are strongly attracted to
 positive ions by their negative ends. (b) They are
  also strongly attracted to negative ions by their
                   positive ends.
Polar water molecules are strongly
      attracted to each other.
Electron Configuration in Ionic Bonding

• Metals tend to lose their valence electrons,
  leaving a complete octet in their next-
  lowest energy level.
• Sodium – (1 valence electron) loses 1
  electron and becomes Na+1.
• Na ([Ne]3s1)  1e- + Na+1([Ne])
• Calcium – (2 valence electrons) loses 2
  electrons and becomes Ca+2.
• Ca ([Ar]4s2)  2e- + Ca+2([Ar])
Electron Configuration in Ionic Bonding

• Nonmetals tend to gain or share valence
  electrons to complete an octet in their
  highest energy level.
• Oxygen – (6 valence electrons) gains two
  electrons to become O-2 .
• O ([He]2s22p4) + 2e-  O-2 ([He] 2s22p6)
• Phosphorus – (5 valence electrons) gains
  three electrons to become P-3.
• P ([Ne]3s23p3) + 3e-  P-3 ([Ne] 3s23p6)
   Formation and Properties of Ionic
             Compounds
• Ionic bonds – forces of attraction that bind
  cations and anions together.
• Ionic compound – consists of electrically
  neutral group of ions joined by
  electrostatic forces.

• Example: Sodium chloride
   Formation and Properties of Ionic
             Compounds
• At room temperature, most ionic
  compounds are crystalline solids, where
  ions are arranged in various 3-D patterns.
• Because of the large attractive forces of
  the ions to each other the compounds
  become very stable and have high melting
  points.
Sodium
Chloride
Crystals
The structure of lithium fluoride.
Electron Configuration in Ionic Bonding
• Scientists have learned that all of the elements
  within each group behave similarly because they
  have the same number of valence electrons.

• Valence electrons - # of electrons in the highest
  occupied energy level of an atom.

• The number of valence electrons is related to
  the group numbers on the periodic table.
Electron Configuration in Ionic Bonding

•   Group 1 elements = 1 valence electron.
•   Group 2 elements = 2 valence electrons.
•   Groups 3-12 elements = 2 valence electrons.
•   Group 13 elements = 3 valence electrons.
•   Group 14 elements = 4 valence electrons.
•   Group 15 elements = 5 valence electrons.
•   Group 16 elements = 6 valence electrons.
•   Group 17 elements = 7 valence electrons.
•   Group 18 elements = 8 valence electrons.
 Determining Valence Electrons for
      an Ion or a Compound
• 1. Multiply the number of valence electrons by
  the number of moles of each element.
• 2. Add up all the electrons for each of the
  elements.
• 3. If there is a charge and it is negative, add
  that number of electrons to the total.
• 4. If there is a charge and it is positive, subtract
  that number of electrons from the total.
• Total # of electrons should always be an even
  number!
Determining Valence Electrons Examples
• Determine the number of valence electrons in
  each of the following compounds and ions:
• NH4+1



• CH2ClBr



• PO4-3
Electron Configuration in Ionic Bonding

• Valence electrons are the only electrons
  involved in bonding, and are the only ones
  written when drawing electron dot structures.

• In forming compounds, atoms tend to achieve
  the electron configuration of a noble gas, having
  8 valence electrons which as known as having a
  stable octet (octet for 8 valence electrons).
Lewis Symbols of Atoms and Ions
• Also known as electron dot symbols
• Use symbol of element to represent nucleus and inner
  electrons
• Use dots around the symbol to represent valence electrons
   – put one electron on each side first, then pair
• Elements in the same group have the same Lewis symbol
   – Because they have the same number of valence
     electrons
• Cations have Lewis symbols without valence electrons
• Anions have Lewis symbols with 8 valence electrons
                        •     ••        ••      ••   ••
Li•    Be•        •B• •C•         •N•       •O:   :F:      :Ne:
        •          •          •    •         •         •    ••
                                       ••      •• -1
            Li•        Li+1        :F:       [:F:]
                                       •       ••
The Nature of Covalent Bonding
• Structural formula – chemical formulas
  that show the arrangement of atoms in
  molecules and polyatomic ions.

• Octet rule – atoms gain or lose electrons
  to acquire the stable electron configuration
  of a noble gas, usually having 8 valence
  electrons.
            Lewis Structures
• You can represent the formation of the
  covalent bond in H2 as follows:


     H   . + .H                      H H:
    – This uses the Lewis dot symbols for the hydrogen
      atom and represents the covalent bond by a pair
      of dots.
            Lewis Structures
• The shared electrons in H2 spend part of
  the time in the region around each atom.


                    H H:
    – In this sense, each atom in H2 has a helium
      configuration.
            Lewis Structures
• The formation of a bond between H and Cl
  to give an HCl molecule can be
  represented in a similar way.

        . + .Cl:
                : :



                                            : :
    H                                 H Cl: :
    – Thus, hydrogen has two valence electrons about
      it (as in He) and Cl has eight valence electrons
      about it (as in Ar).
            Lewis Structures
• Formulas such as these are referred to as
  Lewis electron-dot formulas or Lewis
  structures.           bonding pair



                         : :
                       : :
                   H Cl                  lone pair


    – An electron pair is either a bonding pair (shared
      between two atoms) or a lone pair (an electron
      pair that is not shared).
The Nature of Covalent Bonding
• Exceptions to the octet rule:
  – H needs 2 electrons to be stable
  – Be needs 4 electrons to be stable
  – B needs 6 electrons to be stable
The Nature of Covalent Bonding
• Steps for Drawing Lewis-dot structures
1. Determine the number of valence
   electrons in the molecule.
  - When drawing determining valence electrons
     for an ion, add electrons if it an anion, and
     subtract electrons if it is a cation.
2. The first element in the compound will be
   the central atom. Exception: hydrogen
   will never be the central atom.
The Nature of Covalent Bonding
Steps for Drawing Lewis-dot Structures
3. Use one pair of electrons to bond each
   outer or terminal atom to the central
   atom.
4. Make all outer or terminal atoms stable
   using the valence electrons.
5. Put any remaining electrons around the
   central atom as lone pairs.
The Nature of Covalent Bonding
• Draw the Lewis structure for:
• NH3

• PO43-

• CHFClBr

• PF5-2
The Nature of Covalent Bonding
• Single covalent bond – a bond in which
  two atoms share a pair of electrons.

• Double covalent bond – a bond in which
  two atoms share two pairs of electrons.

• Triple covalent bond – a bond in which two
  atoms share three pairs of electrons.
The Nature of Covalent Bonding
• If you have used up all of the valence electrons
  and you still need two more electrons to make
  the central atom stable, you must have one
  double bond.
• If you still need four more electrons to make the
  central atom stable, you must have either one
  triple bond or two double bonds.
• Double and triple bonds exist most commonly
  between C, N, O, and S atoms.
The Nature of Covalent Bonding
• Draw Lewis structures for:
• NOCl

• CO2

• N2

• SiO3-2
The Nature of Covalent Bonding
• Resonance structures – molecules or ions
  that can have two or more different Lewis
  structures. They must contain a double
  bond to have any resonance structures.
• Resonance structures don’t truly have a
  single bonds or a double bond, but a
  hybrid mixture of bonds where the extra
  bond is spread equally among the other
  single bonds.
The Nature of Covalent Bonding
• Draw Lewis structures for:
• NOCl

• CO2

• N2

• SiO3-2
The Nature of Covalent Bonding
• Single bonds are longer (length between
  the atoms) than double and triple bonds.
• Double bonds are longer than triple bonds.
• Single bonds are not as strong as double
  bonds, and can be broken much easier
  than double bonds.
• Triple bonds are stronger than double
  bonds.
             Bonding Theory
• The valence-shell electron pair repulsion
  (VSEPR) model predicts the shapes of
  molecules and ions by assuming that the
  valence shell electron pairs are arranged as far
  from one another as possible.

   – To predict the relative positions of atoms around a
     given atom using the VSEPR model, you first note
     the arrangement of the electron pairs around that
     central atom.
 Predicting Molecular Geometry
• The following rules and figures will help
  discern electron pair arrangements.
     1. Draw the Lewis structure
     2. Determine how many bonding pairs are around
        the central atom. Count a multiple bond as one
        pair.
     3. Determine how many lone pairs, if any, are around
        the central atom.

     All diatomic molecules have a linear shape.
          Arrangement of Electron
            Pairs About an Atom
2 pairs              3 pairs              4 pairs
Linear           Trigonal planar        Tetrahedral




      5 pairs                       6 pairs
Trigonal bipyramidal               Octahedral
 Molecular Geometry Examples
• NH3          • NOCl

• PO43-        • CO2

• CHFClBr      • SF2

• PF5          • N2

• SeF6         • SiO3-2
   Polar Bonds and Molecules
• Nonpolar covalent bond – equal sharing of
  electrons between two atoms.
• Polar covalent bond – unequal sharing of
  electrons between two atoms.
• In polar covalent bonds the electrons are
  pulled closer to the atom with the larger
  electronegativity value.
• This creates a partial positive and a partial
  negative pole within the bond.
   Polar Bonds and Molecules
• Polar bonds can create polar or nonpolar
  molecules and ions.
• If the centers of partial positive and partial
  negative charge are in the same location,
  the molecule or ion is nonpolar.
• If the centers of partial positive and partial
  negative charge are in different locations,
  the molecule or ion is polar.
   Polar Bonds and Molecules
• If the central atom has lone pairs of electrons,
  the molecule or ion is polar.
• If the central atom does not have any lone
  pairs of electrons, the molecule or ion is
  nonpolar (as long as the atoms around the
  central atom are all the same).
• If there is more than one type of atom around
  the central atom, the molecule or ion will be
  polar regardless of whether or not it has lone
  pair(s) on the central atom.
 Examples: Polar or Nonpolar?
• Determine whether       • CN-1
  each of the following
  molecules or ions are
  polar or nonpolar:      • CH4
• NO2-1
                          • SO3-2

• CFCl3
                          • NOCl
• N2
   Polar Bonds and Molecules
Attractions Between Molecules
• Molecules are attracted to one another by
  a variety of forces.
• These intermolecular forces are weaker
  than ionic or covalent bonds.
• These forces are responsible for whether
  or not a molecular compound is a solid,
  liquid, or a gas.
   Polar Bonds and Molecules
• van der Waals forces – consist of
  dispersion forces and dipole interactions
  (dipole-dipole moments).
• Dispersion forces – weakest of all
  intermolecular forces. They are caused by
  the motion of electrons. The strength of
  dispersion forces increases with the
  increasing number of electrons in a
  molecule.
   Polar Bonds and Molecules
• All molecules contain dispersion forces.
• As molar mass and the number of
  electrons increase, dispersion forces
  increase.
• Halogens are the most common molecules
  to have dispersion forces. Fluorine is a
  gas, Bromine is a liquid and Iodine is a
  solid.
   Polar Bonds and Molecules
• Dipole interactions – occur when polar
  molecules or ions are attracted to one
  another. This occurs when a partial
  positive charge and a partial negative
  charge come close to each other.
• Dipole interactions are very similar to, but
  much weaker than ionic bonds.
   Polar Bonds and Molecules
• Hydrogen bonds – force exerted between
  a hydrogen atom bonded to an F, O, or N
  atom in one molecule and an unshared
  pair on another F, O, or N atom in a
  nearby molecule.
• Hydrogen bonds can have a great effect
  on the boiling point of a substance.
Intermolecular Forces Examples

				
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