Chemical bonding

1. Chemical Bonding

2. Definition
 bond: forces that hold one atom to
another in a compound

3.  To break a bond requires energy to be
put in to overcome the forces of
attraction
 Bond breaking is endothermic.

4.  To make bonds causes a release of
energy.
 Bond making is exothermic.

5.  Compounds have less energy (more
stable) than the substances from which
they form.
 Ex: water has less energy than the
hydrogen and oxygen from which it
formed.
 The energy stored in a bond is potential.

6. Three Types of Bonds
1. Metallic
2. Ionic
3. Covalent

7. Metallic Bonds
 Definition: bonds between atom in a
metal; ions held together in a crystalline
lattice in a “sea of mobile electrons”

8. Metallic Bonds
 Conduct electricity because of freely
moving electrons.
 Any substance that has moving charged
particles (typically either mobile electrons
or ions) will conduct electricity.
 High MP and high BP because bonds
are strong.
 malleable

9. Diagrams

10. Ionic Bonding:
What does the word “ionic” mean?
 Comes from the word ion meaning
charged particle.
 Ions are formed when atoms gain or lose
e-.

11. Electrical Attraction
 Opposites attract.

12. Ionic Bonds
 Formed when electrons are transferred
between atoms.

13. Ionic Bonds: Metals and Non-metals
 Metals… lose e-.
 Non-metals… gain e-.
 They can exchange electrons to form a
bond.

14. The Octet Rule
 Atoms are happy with a full valence
shell.
 This achieved by gaining or losing
electrons.

15.  Metals tend to have low ionization
energies, so they lose e- easily.
 Non-metals tend to have high
electronegativities, so they gain e’
readily.

16. To add to notes…
 The two Elements involved in an ionic
bond have a difference in
electronegativity (E.N.D.) that is greater
than or equally to 1.7
 Ex: NaCl E.N.D.= 2.3 (How do you find
this? Look up electronegativity of each
element on Table S and subtract.)
 Na = 0.9 Cl = 3.2
 3.2 – 0.9 = 2.3

17. Example:
sodium reacts with chlorine
Na Cl
Has one valence
electron, wants
to get rid of it so
that its valence
shell is full
Has 7 valence
electrons, needs
one more to
have a full
valence shell This reactions makes sodium
chloride or table salt.
Na+1[ Cl ]-1

18. See the reaction
 Reaction to form an ionic compound
 In your notes, describe the reaction.
 Are the reactive properties of these
elements consistent with what we
learned last chapter about the groups of
the periodic table?

19. Summary of Concepts:
Ionic Bonds
 Occur between metals (+) and non-
metals (-)
 Involve a transfer of electrons from the
metal to the nonmetal.
 The two elements have an
electronegative difference of 1.7 or
greater
 We call ionic compounds “salts”

20. Covalent Bonds vs. Ionic Bonds

21. Two Hydrogen Atoms
 The valence shells overlap and the
electrons are shared making a more
stable molecule.

22. Covalent Bonds
 Involves SHARING electrons because both
elements have high electronegativities.
 Sharing of electrons can be equal (non-polar)
or unequal (polar)
 Usually is between two NON-METALS (ex. H,
C, O, N…)
 Covalent bonding like Ionic bonding results in
a more stable compound, because the atoms
involved meet the “octet rule”.

23. Covalent Bonding (cont.)
 A “shared” pair of electrons makes a SINGLE BOND
(2 total e-).
 2 “shared” pairs makes a DOUBLE BOND (4 total e-).
 3 “shared” pairs makes a TRIPLE BOND (6 total e-).
 We call them MOLECULES.
 Covalent bonds are very strong bonds.

24. Ionic Bonding
 Involves a transfer of electrons because atoms
have an E.N.D. greater than or equal to 1.7.
 These bonds are always polar because it
involves a + and – ion.
 Usually involves a metal(+) and a nonmetal(-).
(ex. NaCl)
 The oppositely charged ions create an
electrical attraction which forms the bond.
 Ionic bonds result in compounds
(not molecules) that are more stable because
the atoms meet the octet rule.

25. Simple Molecules
(covalent bonds)
 Diatomic molecules are atoms of the
same element that covalently bond to
meet the octet rule.
 The following elements are diatomic:
 H2, O2, F2 Br2, I2, N2, Cl2
 A little trick to remember them:
HOF BrINCl

26. Lewis Dot Diagram of Atoms
and Ions (Review)
Sodium atom Potassium ion
Magnesium atom Calcium ion
Chlorine atom Iodide ion
Aluminum atom Oxide ion
Sulfur atom Aluminum ion

27. 1. NaI
2. Na2S
3. RbBr
4. CaF2
5. AlCl3
6. BaO
7. Li3N
8. K3P
9. MgO
10. BaCl2
Dot Diagrams for Ionic
Compounds

28. Drawing Lewis Dot Diagrams for
Covalently Bonded Molecules
1) Add up all the valence electrons for an atom in the
molecule.
2) The “central atom” is the one that is most
electronegative or there will only be one of this atom.
3) Add the electrons until you reach the total
remembering that all the atoms must obey the octet
rule except for hydrogen which obeys the duet rule.
4) Also remember that one pair of electrons between
two atoms is a single bond two is a double and three
pairs is a triple bond. These are known as bonded
or shared pairs of electrons.
5) Not every electron has to be bonded. We call these
the lone or non-bonded pairs because they are
“lonely”.

29. Examples
H2 F2 Cl2
H2O NH3 CH4

30. Practice on your own …

31. Covalently Bonded-Molecular
Cmpds

32. Coordinate Covalent Bonds
When a shared pair of electrons comes from only one atom,
not two its known as a coordinate covalent bond.
Polyatomic Ions have these.
+1
+1

33. Polyatomic Ions dot diagrams

34. Metal and a Polyatomic Ion Dot
Diagrams

35. Bond Polarity
Polar Covalent Bonds (between two atoms)
 Covalent bonds unlike ionic bonds, are NOT
composed of oppositely charged ions.
 However, we find that since the electrons are shared
and each element has a different electronegativity,
the shared electrons are often shared unequally.
 Meaning the electrons spend more time around the
more electronegative atom.
 This causes one atom to be slightly negative and one
atom to be slightly positive.
 This is known as a dipole moment. One positive end
and one negative end 2 oppositely charged poles.

36. Bond Polarity
Non-polar Covalent Bond (between two
atoms)
 If both of the atoms involved in a bond have
the same electronegativity then the electrons
are shared equally.
 This will happen if a bond forms between two
of the same atoms. Since the atoms are the
same they will have the same electronegativity
and share the electron(s) equally.
 When electrons are shared equally their will
never be a dipole moment and the bond will be
non-polar.

37. Molecule Polarity
Polar Molecule
 A molecule that is polar (also known as
a dipole) is asymmetrical which results
in an uneven distribution of charge
throughout the entire molecule.
 Ex:
H2O which has a bent shape and is
assymetrical

38. Molecule Polarity
Non-Polar Molecules
 A molecule that is non-polar is
symmetrical in shape which results in an
even distribution of charge.
 Ex
CH4 which is tetrahedral symmetrical in
shape

39. Bond Polarity vs. Molecule
Polarity
 Bond Polarity is based on the electronegativity
between the bonded atoms.
 If they have different electronegativities (different
atoms) the bond must be polar.
 If the have the same electronegativity (same atoms)
the bond must be non-polar
 Molecule polarity is based on the shape symmetry
of the entire molecule.
 If the molecule has a shape that is symmetrical then
the charge distribution is even and the molecule is
non-polar.
 If the molecule has a shape that is asymmetrical then
the charge distribution is uneven and the molecule
is polar.

40. Shapes of Molecules
 The shape of a molecule will help
determine whether or not the molecule is
polar or nonpolar
 Examples of Shapes of Molecules
CO2

41. Shapes Continued
CH4

42. Shapes Continued
HCl

43. Shapes Continued
NH3

44. Shapes Continued
H2O

45. Summary of Shapes
1 or 2 bonds only linear
2 bonded pairs and 2 lone
pairs
bent
3 bonded pairs and 1 lone
pair
pyramidal
4 bonded pairs and 0 lone
pairs
tetrahedral
If the central atom has ... the shape is example

46. Intramolecular and
Intermolecular Forces
 Atoms involved in a bond are held together by
what is known as an Intramolecular Force.
 An intramolecular force is just simply a bond
so it can be either ionic, polar covalent or
nonpolar covalent.
 When substances are in a solid or liquid state
the compounds or molecules are held together
by a force as well.
 The force that holds molecules together in a
solid or a liquid sample is known as an
Intermolecular Force.

47. Ionic Compounds (ionic bonds)
 For solid or liquid Ionic compounds the only
intermolecular force that is possible is created by the
attraction between the oppositely charged ions. This is
known as a dipole-dipole attraction between ions.
 For example in a Salt Crystal (NaCl) the positive and
negative ions will alternate as shown in the diagram
below. It is the opposite charges of the ions that hold the
crystal together.

48. Molecular Compounds
(polar covalent or nonpolar covalent)
 For solid or liquid molecular the compounds
there are three main intermolecular forces that
hold the molecules together.
 Some of the intermolecular forces for
molecules are strong, because like the ionic
solids it is due to a dipole-dipole attraction.
 Some of the intermolecular forces are weak,
because they are only based on the size of the
molecules.
 Intermolecular forces for molecules that
involve a dipole-dipole attraction only exist
in polar molecules.

49. Polar Molecules
 Hydrogen bonding is one type of intermolecular force
it is formed between a hydrogen atom in one molecule
and a nitrogen, oxygen, or fluorine atom in another
molecule.
 Hydrogen bonding is the strongest intermolecular
force for all molecular substances.
 Substances with H-bonding have relatively high
boiling points, because the force is so strong.
 The most important example of this is found in water.
The strength of Hydrogen bonding is the reason why
water has a relatively high boiling point compared to
other molecular compounds.

50. Polar Molecules (Continued)
 If water did not have H-bonding it would not exist
on earth as a liquid only a vapor. The diagram
below shows how the H-bonding occurs.
 H-bonding also occurs between ammonia
molecules and Hydrogen Fluoride molecules.

51. Polar Molecules (Continued)
 All other intermolecular forces for polar
molecules are due to the dipole-dipole
attractions between the molecules, and
the force is not quite as strong as
Hydrogen bonding.
 Remember H-Bonding only occurs
between the HYDROGEN atom of one
molecule and either an OXYGEN,
NITROGEN or FLUORINE atom of
another molecule.

52. Nonpolar Molecules
 For nonpolar molecules there is no dipole so the force
of attraction between the molecules is based on the
molecular mass of the molecules.
 The main rule for non-polar molecules is the smaller
the molecules the weaker the intermolecular forces,
and the larger the molecules the stronger the
intermolecular forces.
 Smaller non-polar molecules tend to be gases at room
temperature, because the weak intermolecular forces
give them low boiling points.
 Examples: methane (CH4) and Hydrogen (H2)are
gases at room temp.
 Larger non-polar molecules will have stronger
intermolecular forces so they will usually be liquids or
solids at room temperature.
 Examples: Glucose (C6H12O6) and Gasoline (C8H18)

53. Summary
1 Intramolecular forces are chemical bonds between
the atoms that make up an individual molecule.
2. Intermolecular forces are attractions between two
or more molecules.
3. H-bonding is one type of intermolecular force that
occurs between the hydrogen of one molecule and
the oxygen, nitrogen or fluorine of another molecule.
4. H-bonding is the strongest intermolecular force for
polar molecules, which is why these substances will
have high boiling points.
5. All other polar molecules have a dipole-dipole
attraction.

54. Summary (continued)
6. Nonpolar molecules have intermolecular forces that
depend on the mass of the molecules.
7. Molecules with a high molecular mass will have
strong intermolecular force.
8. Molecules with a low molecular mass will have weak
intermolecular force.
9. A substances boiling point and its intermolecular
forces are directly related.
10.Substances with high boiling points have strong
intermolecular forces
11.Substances with low boiling points have weak
intermolecular forces.

55. Properties of Bonds
 Metallic, covalent and ionic bonds all
have varying properties.
 These properties can be used to identify
the bonding of unknown substances.
 The three main properties that can be
used to distinguish bond type are
Melting & Boiling Points, Hardness in
the solid state, and Conductivity in the
solid, liquid and aqueous states.

56. Properties of Metallic, Ionic and
Covalent Bonds Summary
Hard
High
Ionic
Soft
Low
Covalent
Hard
High (except Hg)
Metallic
Hardness
Melting & Boiling
Points
Bond
Type
Conductivity
Aqueous
Liquid
Solid
Yes
Yes
No
No
No
No
Yes
Yes
Yes

57. Properties of Metallic, Ionic and
Covalent Bonds Summary
Bond
Type
MP
&
BP
Hardness
Conductivity
Solid
State
Liquid
State
Aqueous
Metallic High (except
Hg)
Hard Yes Yes Yes
Ionic High Hard No Yes Yes
Covalent Low Soft No No No

58. Why do some substance have
conductivity and others do not?
 In order for a substance to conduct electricity and heat the substance
must have charged particles that are free to move or mobile.
 Metallic substances can always conduct electricity and heat because
the bonds are created by free flowing or mobile valence electrons
between the metal atoms. These mobile electrons are what move the
electrical current through the substance.
 Ionic substance are made up of positive and negative ions (charged
particles) but they are not free to move in the solid phase so they can
not conduct as soilds.
 When Ionic substances are melted or dissolved in water the ions
are free to move so they have the ability to conduct in the liquid and
aqueous states, because of the mobile ions.
 Molecular substances are made up of atoms sharing electrons in a
covalent bond and since they do not have any charged particles they
can NEVER conduct electricity.

59. Shapes of Molecules
 CO2 – linear, space-filling model vs. ball and
stick
 Polar bonds (C-O), Non-polar molecule b/c it is
symmetrical

60. Shapes of Molecules
 CH4 – tetrahedral shape, ball and stick model
 Polar bonds (C-H), Non-polar molecule b/c it is
symmetrical

61. Shapes of Molecules
 HCl – linear shape, ball and stick model
 Polar bonds- Polar Molecule, asymmetrical

62. Shapes of Molecules
 NH3 – trigonal pyramidal shape
 Polar bonds (N-H) and polar molecule b/c
asymmetrical.

63. Shapes of Molecules
 H2O – bent shape
 Polar bonds (O-H) and polar molecule
b/c asymmetrical.

64. Solid State Chemistry-
Crystalline Solids

65. Ionic Solids
 Repeating particles in crystals are ions
(+/-).
 Examples: NaCl, KNO3, CaCl2, etc.
 Diagram

66. Ionic Solids
 Properties:
 High melting point (strong bonds)
 Electrically conductive ONLY when
dissolved in water or in liquid form (mobile
charged particles)
 Water soluble (exceptions on Table F)
 Relatively hard

67.  Molecular Solids: Repeating particles
are atoms or molecules.
 Examples: CO2, H2O, S8
 Diagram:

68. Molecular Solids
 Properties:
 Relatively low MP and BP (based on
Intermolecular forces)
 Soft
 Non-conducting
 Only polar molecules are soluble in water

69. Covalent-Network Solids:
 Repeating particles are atoms held
together by covalent bonds NOT
intermolecular forces.
 Examples: Diamond, quartz, SiO2, SiC
 Diagram

70. Covalent Network Solids
 Properties:
 Very strong covalent bonds
 Very high melting point
 Very hard
 Not soluble in water

71. Graphite is weird!

72.  Allotrope: different molecular structure of the same
element; also has different properties
 Examples: Carbon (diamond vs. graphite (C8) vs.
buckminster fullerenes (C60), Oxygen vs. Ozone

73. Metallic Solids
 Repeating particles are metal atoms with
mobile valence electrons moving
throughout the crystal; a “sea of mobile
electrons”
 Examples: Ag, Au, Na, Cu, Zn
 Diagram

74. Metallic Solids
 Properties:
 Some are hard, others are soft
 Conduct electricity and heat
 Malleable/ductile
 insoluble