Covalent Bonds and Molecular Forces

1
Covalent Bonds and Molecular Forces

• Chapter 6

2
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.

3
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.

4

• Molecular orbital is a region where an electron
  pair is most likely to exist as it travels in the
  three dimensional space around the nuclei.
• 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.

5
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.

6

• 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.

7
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
• energy.

8
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.

9
(No Transcript)
10
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.

11

• There is a bond between these two extremes in
  which electrons are shared but not equally. These
  bonds are called as polar covalent bonds.

12

• 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.

13

• 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 dipole.

14
Homework

• Page 201
• Q.5 and Q.6

15
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.

16

• 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.

17
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.

18
Class Practice

• Draw Lewis structure for iodine monochloride, ICl
  and hydrogen bromide HBr.
• Draw Lewis Structure for formaldehyde CH3OH.

19
Resonance Structures

• A possible Lewis dot structure of a molecule for
  which more than one Lewis structure can be
  written.

20
Class Practice

• Page 211
• Q.4 all

21
Naming Covalent compounds

• The most common naming system uses prefixes,
  roots and suffixes.
• Example Carbon dioxide and carbon monoxide.
• 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.

22

• 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.

23
Home work

• page 213
• Q.7. b and d
• Q.9. all
• Q.11. a and c

24
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?

25
Different possible shapes
26
Molecular geometry based on electron pairsCO2
SO2
27

• 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.

28
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 counts.

29

• 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.

30

• 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.

31

• 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.

32

• 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
  120.

33

• 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.

34

• 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 109.5.

35
Class Practice

• Page 217
• Q.1 Q.2.

36
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.

37
H2O, CO2, and CH4 molecules
38

• 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 105 angle to each other, which
  creates a dipole for the molecule.

39
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.

40
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.

41
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 substance.

42
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.

43

• 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.

44
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.

45

• 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.

46
Waters 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.

47
London forces/ Van der Waals forces

• In case of noble gases the boiling points
  increases in this
• way.
• Higher boiling point indicate
• the addition of electrons in
• the atoms and hence strong
• bonds. This was
• explained by Fritz London.

48

• 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.

49
Homework

• Page 227
• Term Review all
• Page 228
• 14, 16 and 17.