Chapter 6 - Covalent Bonds and M

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					Covalent Bonds and Molecular
          Chapter 6
           Sharing electrons
• Sodium atom reacts with chlorine gas to form
  the ionic compound sodium chloride, NaCl, is
  an example of this type of reaction. The
  reaction of hydrogen and oxygen to form
  water is another kind of rearrangement where
  electrons are shared.
    Molecular and Atomic orbitals
• The simplest example of sharing electrons occurs
  mainly in diatomic molecules such as H₂, and O₂.
• When two hydrogen atoms approach each other, the
  positive nucleus of each atom attracts its own electron
  and the electrons of the other atom. At the same time
  the positive nuclei of the two atoms repel each other.
  Likewise the electron cloud of the atoms repel. Since
  they are both of the same atom neither has enough
  attraction to take an electron from the other. Instead of
  forming ions , the 2 hydrogen atoms share electrons.
  The shared electrons moving about in space
  surrounding the the two nuclei are in molecular orbital.
• Molecular orbital is a region where an
  electron pair is most likely to exist as it travels
  in the three dimensional space around the
• A bond formed when two or more valence
  electrons are attracted by the positively
  charged nuclei of two atoms and are thus
  shared between both atoms.
    Potential energy curve for H₂
• As a covalent bond forms between two atoms,
  they reach a distance from each other at
  which the attractive and repulsive forces are
  balanced and the energy is at the minimum.
• As the two hydrogen atoms come nearer the
  potential energy of the combination becomes
  lower and lower until it reaches the minimum
  value of -436kJ/mol at a distance of 75pm.At
  the lowest energy, the H-H combination is
  most stable because lower energy means
  greater stability. At the distance of 75pm, the
  repulsion between the like charges equals the
  attraction of the opposite charges.This is the
  bond length.
          Diatomic Molecules
• Bond length- The distance between two bonded
  atoms at their minimum potential energy; the
  average distance between two bonded atoms.
• The energy
required to break a
bond between two
atoms is the bond
   Electronegativity and bonding
• The tendency of an atom to attract bonding
  electrons to itself when it bonds with another
  atom is electronegativity.
• To help explain why some combinations of
  atoms form ionic bonds and some form
  covalent bonds, this concept was developed
  by Linus Pauling. In general electronegativity
  decreases down a group and increases across
  a period.
Polar and non polar covalent bonds
• In a molecule such as H₂, the atoms are
  identical, so they pull on the bonding
  electrons with the same force. The electrons
  are shared equally. Such a covalent bond, in
  which the bonding electrons are shared
  equally, is called a nonpolar covalent bond. In
  other words the electronegativities of two
  atoms are equal. If the electronegativities are
  greatly different, an ionic bond is formed.
• There is a bond between these two extremes
  in which electrons are shared but not equally.
  These bonds are called as polar covalent
• In the previous example of polar covalent
  bond oxygen attracts electrons more strongly
  than the other atoms.
• Polar molecules have both positive and
  negative charges. Example hydrogen fluoride.
  The electronegativity of fluorine is much
  higher than the electronegativity of hydrogen.
  The fluorine atoms attracts electron much
  more than hydrogen atoms.
• The hydrogen having its electron pulled away
  has a partial positive charge and the fluorine
  has a partial negative charge. This is not an
  ionic bond. A molecule that has a partial
  positive charge on one end and partial
  negative charge on the other end is called a
• Page 201
• Q.5 and Q.6
        Electron Dot Structures
• Valence electrons are electrons in the
  outermost energy level of an atom, where it
  can participate in bonding.
• Lewis Structure is a structure in which atomic
  symbols represent nuclei and inner shell
  electrons, and dots are used to represent
  valence electrons.
• Consider a chlorine atom, which has the
  electronic configuration 1s²2s²2p⁶3s²3p⁶.
• Only the electrons in the outermost energy
  level are involved in bonding., so in the Lewis
  structure only seven valence electrons are
  represented by dots.
  Rules for Drawing Lewis Structures
           with many Atoms
• 1. Hydrogen or halogen atoms often bind to only
  one other atom and are usually on the outside or
  the other end of the molecule.
• 2. The atom with the lowest electronegativity is
  often the central atom. These atoms often have
  fewer than seven electrons and may form more
  than one bond.
• 3. When placing valence electron s around an
  atom, place one electron on each side before
  pairing any electrons.
              Class Practice
• Draw Lewis structure for iodine monochloride,
  ICl and hydrogen bromide HBr.
• Draw Lewis Structure for formaldehyde
         Resonance Structures
• A possible Lewis dot structure of a molecule
  for which more than one Lewis structure can
  be written.
              Class Practice
• Page #211
• Q.4 all
    Naming Covalent compounds
• The most common naming system uses
  prefixes, roots and suffixes.
• Example Carbon dioxide and carbon
• Prefixes and suffixes are usually attached to
  root words. For binary compounds the root
  word is the name of the element. The first
  element named is usually the one first written
  in the formula which is the least
  electronegative element.
• If the molecule contains only one atom of the
  first element given in the formula, the prefix
  mono is omitted in the name of the
  compound. For example , to distinguish
  between the two oxides of carbon, the
  prefixes mono and di are used.
                Home work
 page 213
Q.7. b and d
Q.9. all
Q.11. a and c
           Molecular shapes
• The shape of a molecule can be predicted by
  the Lewis Structure.
• In a molecule of only two atoms, such as HF,
  or H₂, only as linear shape is possible.
• Molecules of more than two atoms, molecular
  shapes will vary. Example CO₂ and SO₂ their
  formulas are similar then why carbon dioxide
  is linear, while sulfur dioxide is bent?
Different possible shapes
Molecular geometry based on
       electron pairs

    CO₂            SO₂
• There is a simple model that can be used to
  determine the three dimensional arrangement
  of the atoms in a molecule. This model is
  based on the valence shell electron pair
  repulsion (VSEPR) theory.
• According to this theory you can predict the
  shape of a molecule by knowing the electron
  pairs around a central atom.
Steps in determining the geometry
 of a molecule or polyatomic ion.
• 1.For a molecule ,count the number of
  electron pairs surrounding the central atom.
  Each single or multiple bond counts as one
  electron group. Each nonbonding electron pair
• There are two double bonds around the
  central carbon atom. Therefore there are two
  electron groups around this atom. Electron
  groups have negative charge, and like charges
  would repel each other and will remain as far
  apart as possible. If a central atom has 2
  electron groups they will be linear.
• In SO₂ one of the electron pairs is a lone pair.
  The nonbonding electrons repel the bonding
  electrons , causing the three electron pairs to
  orient in a trigonal planar geometry.
• When a molecule consists of a central atom
  bonded to three other atoms its shape will be
  trigonal planar as long as there are only three
  electron groups determining the geometry.
• SO₃ can have resonance structures. Because of
  the resonance structure the three electron
  clouds surrounding the sulfur atom in SO₃ are
  identical in order to be as far as possible the
  groups arrange like three spokes of a wheel,
  extending out from the sulfur atom. This
  geometry is called as trigonal planar.The angle
  between them will be 120⁰.
• When there are 4 molecules surrounding a
  central atom the electron pairs are farthest
  away when they orient themselves towards
  the corner of a tetrahedron.
• If the pairs are all equivalent (if there are four
  identical bonds on the central atom) all four of
  the angles between the bonds are 109.5⁰.
               Class Practice
• Page 217
• Q.1 & Q.2.
          Shape and property
• The shape of a molecule affects the chemical
  properties. Many of these properties depend
  on the polarity of the molecule. For molecules
  containing more than two atoms, molecular
  polarities depends on both the polarity of
  each atom and its orientation.
• The bond polarities in a water molecule add
  together, causing a molecular dipole. In
  carbon dioxide, bond polarities extend in
  opposite direction, cancelling each other.
H2O, CO2, and CH4 molecules
• The double bonds between C and CO2 are
  polar because oxygen attracts electrons more
  strongly than does carbon. However , the
  linear shape of the molecule causes the two
  bond dipoles to act in opposite directions,
  canceling each other and causing the
  molecular polarity to be zero. In the water
  molecule, the polar H-O bonds are oriented at
  a 105⁰ angle to each other, which creates a
  dipole for the molecule.
      Properties Of Compounds
• Covalent compounds melt at a lower
  temperature than ionic compounds.
• Ionic compounds consists of ions each of
  which is attracted to all ions of opposite
  charges. These attractions hold the ions tightly
  in a crystal lattice that can be disrupted only
  by heating to very high temperatures.
          Intermolecular forces
• The attraction that exists between molecules
  are called as intermolecular forces. If there are
  no intermolecular forces between molecules
  then the substance exists as gases.
            Intramolecular force
• An intra molecular force is any force that holds
  together the atoms making up a molecule.
• Intra molecular forces of attraction (covalent) are
  stronger than the intermolecular forces of
  attractions. The stronger the intermolecular forces,
  the higher the melting and boiling point of the
              Dipole Forces
• Dipole forces affect the melting and boiling
  points. In a polar molecule we have one end
  of the molecule having partial positive charge
  and the other end having a partial negative
  charge.The positive end of a molecule can
  attract the negative end of another molecule
  holding the two molecules together. This force
  that exists between the two positive and
  negative ends is called as dipole force.
• The dipole forces help the molecules to exist
  as a solid or a liquid. Oxygen and methane are
  non polar molecules but molecules of water
  and ethyl acetate are polar. Because of the
  dipole forces these have higher melting and
  boiling points.
    Hydrogen bonds are stronger
           dipole forces.
• Hydrogen atom bonded to an atom that is
  more electronegative.
• HF has a very high boiling point
 and HCl has the lowest. There
 is a strong dipole force between
 HF molecules due to large electronegativity
 difference between H and F.
• Hydrogen also has just one electron, and when that electron
  is pulled away then there are no electrons to protect the
  nucleus so the proton in the nucleus is attracted to the
  electron rich fluorine end of another HF molecule.
• Hydrogen bonds are usually formed with small atoms with a
  high electronegativity, like oxygen, fluorine and nitrogen. The
  HCl, HI, and HBr molecules are polar but they are much larger
  than HF, the distance between the molecules is greater and
  so the hydrogen bonds are weaker.
      Water’s unique properties.
• Water has a high boiling point as the
  molecules of water are held together by
  hydrogen bonds.
• Water (H₂O) has a higher b.p than hydrogen
  sulfide (H₂S) as the electro negativity
  difference between H and S is 0.4 and that of
  H and O is 1.2. As a result, the hydrogen bonds
  in H₂O is stronger than H₂S.
   London forces/ Van der Waal’s
• In case of noble gases the boiling points
  increases in this
 Higher boiling point indicate
 the addition of electrons in
the atoms and hence strong
bonds. This was
explained by Fritz London.
• London forces are an attraction between
  atoms and molecules caused by the formation
  of instantaneous dipoles in the atoms and
  molecules because of the unequal distribution
  of electrons around the nucleus or nuclei.
•   Page 227
•   Term Review all
•   Page 228
•   14, 16 and 17.