1. Chapter 1: Chemical Bonding 1.1 Forming and Representing Compounds
A. The Basics
scientists have studied the way in order to
elements and compounds appear in nature
categorize
chemical bonding 
most are combined with in nature (called ) and are
metals
non-metals
ores
solids

2. 
a few are found in their
metals
pure form
…precious metals 
metals (except ) in pure form are Hg
solids

3. 
non-metals
combine with one another to form
solids, liquids or gases 
the only elements that are found in in nature are the
never
combined form
noble gases

4. atoms in such a way that they create a
gain, lose or share electrons
full outer energy level…
called the
octet rule
the can only hold and therefore satisfies the octet rule when it has
first energy level
two e
two e in it
octet rule is a
…not all elements follow it at all times
guideline

5. are the electrons in the
valence e
outermost
energy level of an atom
they are the only electrons involved in
chemical bonding
for representative elements ( ) group number (ignore the “1” in front of groups above 10) tells you the number of
groups 1,2 and 13-18
valence electrons
period number tells you the number of
energy levels occupied by electrons

6. for many , the number of valence electrons
transition metals
is not as predictable… it depends on the
environment
around the ion
eg)
iron can be Fe3+ or Fe2+
the can be used to determine
number of valence e
ion charge
eg)
Fe3+ had 3 valence electrons

7. B. Electron Dot Diagrams
you can’t see atoms and electrons, therefore it is convenient to to show the structure and formation of
draw models
chemical bonds
an is one such model
electron dot diagram
consists of the with
symbol for the element
dots
representing the
valence e
when drawing the diagrams, look up the , then place
number of valence e
dots around the symbol clockwise
for a maximum of
four dots

8. if you have more electrons to place, go back to the
of the symbol and
start pairing up the e     Si Na Mg Al               S   Ar  P   Cl             

9.     O   a orbital is called a and is (at this level) full
lone pair
not involved in bonding
a orbital contains a
half full
bonding electron  
lone pairs   O  
bonding e
the of an atom is the maximum number of that it can form (equals the number of )
bonding capacity
single covalent bonds
bonding e

10. H C Mg P He F K Be S Br Try These
Draw the Lewis diagram (electron dot diagram) for each of the following: H C Mg P He F K Be S Br

11. Na Na C. Ionic Bonding + an is the
electrostatic attraction between oppositely charged ions
ionic bond
most have
three or fewer valence e
metals
they tend to these electrons and become
lose
positive ions (cations) + Na Na

12. O O 2- most have more than four valence e non-metals
they tend to and become
gain electrons
negative ions (anions) 2- O O

13. after ions form, the attraction between the
positive charge
and
negative charge
draws the
ions
together, forming an
ionic bond
when drawing the electron dot diagrams for ionic compounds:
the number of electrons by the must the number of electrons by the
lost
metal
equal
gained
non-metal
the on the compound must be
net charge
zero
you may have to have of the to balance out the
more than one
metal and/or non-metal
charges

14. Examples Na Cl
NaCl Mg O
MgO

15. Examples F Ca F
CaF2 K S K
K2S

16. O Fe O Fe O
Fe2O3 Mg N Mg N Mg
Mg3N2

17. notice the following about the diagrams:
the metal has (since they them)
the non-metal has the valence level
both ions have and the
charges = charges
no valence electrons
lose
filled
square brackets
charge
positive
negative

18. D. Covalent Bonding a is formed when
two non-metals share a pair of electrons
covalent bond
compounds containing are called
covalent bonds
molecular compounds
ions are not formed!!!
electron dot diagrams used to show molecular compounds are called
Lewis structures

19. instead of transferring electrons, valence electrons are now to satisfy the
shared
octet rule
the electrons that are shared are called
a bonding pair
sharing two or three pairs of electrons between two atoms results in a
double or triple bond,
respectively

20. to draw the structures:
place the atom with the in the
most bonding electrons
centre
arrange all other atoms around it as as possible
symmetrically
to make sure that all atoms have the (remember that hydrogen only needs electrons to be satisfied)
share electrons
octet rule satisfied
two

21. eg) PH3 H P H H P H

22. 1. HCl 4. NBr3 2. CH4 5. C2H4 3. F2 6. N2 Try These
Draw the Lewis diagram (electron dot diagram) for each of the following:
1. HCl
4. NBr3
2. CH4
5. C2H4
3. F2
6. N2

23. H Cl N Br C H C H F N

24. P H P H E. Structural Formulas a is another way of drawing molecules
structural formula (diagram)
to draw them, figure out the Lewis structure then replace all
shared pairs of e with a line and leave off the lone pairs
eg) PH3 Lewis Diagram Structural Diagram P H P H

25. 1. HCl 4. NBr3 2. CH4 5. C2H4 3. F2 6. N2 Try These
Draw the structural formula for each of the following:
1. HCl
4. NBr3
2. CH4
5. C2H4
3. F2
6. N2

26. H Cl N Br C H C H F N

27. F. Metallic Bonding
most metals are at room temperature which means that there must be
solids
strong attractive forces
holding the atoms of a pure metal together
metals form or with
ionic bonds
DO NOT
covalent
other metal atoms
in all the atoms
share all the valence e
metallic bonding
the valence electrons are , which means they are from one atom to another
delocalized
free to move

28. metallic bonds are made up of a network of
positive metal ions
in a
“sea” of electrons
a is the
electrostatic force of attraction between the positive metal ions and the negative sea of electrons
metallic bond
this theory helps explain the of metals
properties
eg) good conductors of electricity and heat, ductility, malleability

29. “sea” of delocalized electrons
Metallic Bond Model
metal cations
“sea” of delocalized electrons

30. 1.2 The Nature of Chemical Bonds A. Electronegativity
the of an element is the relative measure of the ability of an atom to
electronegativity
attract electrons in a chemical bond
there is an attraction between the of an atom and the
nucleus (protons)
valence e in an adjacent atom
nucleus
electrons

31. each element is designated a number to represent
how strong it’s nucleus is at attracting another atom’s valence e

32. electronegativity means
higher
greater attraction
(affinity)
trend on periodic table – electronegativity and
increases across period
decreases down group
since do not readily react with other substances, electronegativities have been assigned to them
noble gases
not
understanding electronegativity has contributed to the knowledge of bonding in ionic and molecular compounds

33. Electronegativity and the Periodic Table
decreases
increases

34. B. Size & Electronegativity
as you move from left to right across a period, both the and
electronegativity,
atomic number increase
however size of the atom
decreases

35. here is why size decreases across a period:
the size of an atom depends on the of the containing the
radius
energy level
valence e
in any given period, the valence e of each atom occupy the
same energy level
as you move across the period, the and thus the in the nucleus
atomic number increases
number of protons
increases
there is a between the and when there are more , therefore the atom is
greater amount of attraction
nucleus e
protons
smaller

36. Li C F Period 2 Elements 3 p+ 6 p+ 9 p+ 1 valence e 4 valence e

37. so, the next question is “why does electronegativity increase when atomic size across a period decreases?”
the strength of the attraction (and therefore electronegativity) between oppositely charged particles depends on two factors:
the between the charges – the attractive force between opposite charges with the between them
distance
decreases
square of the distance
the of the charges – the attractive force is to the
magnitude
directly proportional
amount of charge…(simply put the greater
difference in electronegativity the greater
the attractive force!)

38. this means that an atom that is and has lots of (like fluorine) will have a amount of electrostatic attraction (electronegativity) for the of another atom
small
protons
very large e
big atoms have but they are by the therefore have a amount of attraction (electronegativity) for the of another atom
lots of protons
shielded
inner levels of e
small e

39. Cs Si  F Si  nucleus of cesium nucleus of fluorine
distance between nucleus of cesium and valence electrons of silicon
distance between nucleus of fluorine and valence electrons of silicon
nucleus of cesium
nucleus of fluorine
valence electrons of silicon
valence electrons of silicon

40. C. Bond Type & Electronegativity
electronegativities can be related to bond types:
ionic bonds occur between
metals and non-metals
metals have electronegativities and will while non-metals have and will
low
lose e
high
electronegativities
gain e
the two ions that are formed will
attract each other
and form a
chemical bond

41. covalent bonds occurs between
non-metallic atoms
if you look at two atoms that have the electronegativity, like in H2(g), the two nuclei of the atoms will attract the
same
electrons
with exactly the
same strength
the electrons are
shared equally between the two atoms

42. when two non-metals that have
when two non-metals that have electronegativities share electrons, the sharing is
different
no longer equal
the element with the electronegativity pulls the
higher
e closer to itself

43. this results in one end of the bond having a
this results in one end of the bond having a and the other end of the bond having a
slightly negative charge
()
slightly positive charge
(+) + 
bonds that have are called
unequal sharing of electrons
polar covalent bonds
also called since the bonds have
bond dipoles
oppositely charged ends

44. Electronegativity RECAP
The real deal:
Extra help (on website)

45. Calculate the Electronegativity
Subtract the element with the lower electronegativity from the higher electronegative element
Ex: N-H
3.0 – 2.2= 0.8

46. H – F + - Bond Dipole Arrows
“arrow” points towards element with higher electronegativity (-)
“+” at the end that is +
H – F + -

47. - + - + + - + - + - - + 6. C – H 1. H – H 7. Cl – Cl
Try These:
Draw the bond dipole arrow, label the + and  ends, and state the bond type (polar, nonpolar, ionic)
0.4
C – H - +
H – H
polar
nonpolar
0.8 - +
Cl – Cl
N – H
nonpolar
polar
1.3
2.0
Si – Cl + -
B – F + -
polar
polar
0.8
1.2
S – O + -
O – H - +
polar
polar
P – H
Na – Cl
nonpolar
ionic

48. Difference in Electronegativity
you can use the difference in electronegativity between two atoms to determine
bond character
Difference in Electronegativity
3.3
1.7
0.5
mostly ionic
slightly polar covalent
polar covalent
non-polar
covalent

49. bond classification is not simple
as you can see,
bond classification is not simple
bonding is considered a and there is
continuum
ionic and covalent bonding
no clear distinction
between

50. B. Structure of Molecules
in molecular compounds, covalent bonds exist between
specific pairs of atoms
these compounds exist as that have a and therefore are not necessarily written with the
molecules
given number of atoms
lowest whole number ratio

51. C. VSEPR & Structure of Molecular Compounds
the states that molecules adjust their shapes so that valence e-
valence shell electron pair repulsion (VSEPR) theory
are as far away from each other as possible
electron pair repulsion is
not always equal…
it is greatest between two
lone pairs (LP),
less between a and a LP
bonding pair (BP),
and lowest between two
BP’s
shape is determined around the
central atom

52. shapes can be classified into five categories:1.
linear –
two other atoms
central atom is bonded to and has lone pairs, or there is only two atoms in the molecule
zero
eg) CO2(g), HCN(g), HCl(g)  C O  C N H  Cl H

53. central atom is bonded to and has lone pairs trigonal planar – 2.
central atom is bonded to and has lone pairs
trigonal planar –
three other atoms
zero
eg) CH2O(l) O H C  3.
tetrahedral –
central atom is bonded to and has lone pairs
four other atoms
zero H C
eg) CH4(g)

54. central atom is bonded to and has lone pair pyramidal – one4.
three other atoms
central atom is bonded to and has lone pair
pyramidal –
one H N 
eg) NH3(g) 5.
bent –
two other atoms
central atom is bonded to and has either lone pairs
one or two
eg) H2O(l), HNO(g) H O  O H N 

55. we can use a to determine the shape of a molecule around the code
central atom
the code has two numbers:
1. the number of attached to the central atom
atoms
2. the number of on the central atom
lone pairs
CH4
eg) NH3(g) H C H N 
3 - 1
4 - 0
pyramidal
tetrahedral

56. 4 – 0 tetrahedral CH4 3 – 0 trigonal planar CH2O 3 – 1 pyramidal NH3
Code
Shape
Example
4 – 0
tetrahedral
CH4
3 – 0
trigonal planar
CH2O
3 – 1
pyramidal
NH3
2 – 1
bent
HNO
2 – 2
bent
H2O
***all other codes are
linear

57. Pg. 2-4 columns 3 & 4 in Workbook. Lets do the first one together.
Your assignment
Pg. 2-4 columns 3 & 4 in Workbook.
Lets do the first one together.
NH CBr4
Shape Code= Shape Code= 4-0
Shape name= pyramidal Shape Name= tetrahedral
Shape Diagram Shape Diagram

58. D. Polar Bonds & Polar Molecules
a molecule that contains can be overall
polar covalent bonds
nonpolar
the individual are that can be to each other
bond dipoles
vectors
added
if the bond dipoles are and , they each other out resulting in a
equal in strength
opposite in direction
cancel
nonpolar molecule
this canceling happens in
symmetrical molecules
if the bond dipoles , the entire molecule will have a
do not cancel
slightly positive and slightly negative end…
called dipoles

59. general rules:
tetrahedral: if all atoms attached have the same pull (in or out), if different atoms attached
nonpolar
polar
trigonal planar: if all atoms attached have the same pull (in or out), if different atoms attached
nonpolar
polar
pyramidal: as long as there is a difference in electronegativity between the atoms
polar
bent:
polar
linear: …look at electronegativity difference
polar or nonpolar

60. Examples 1. H2O 2. HCl O H H Cl polar polar 3. C2H2 4. C2HI I H C H C
nonpolar

61. Try These H F C H 1. HF 2. CH4 polar nonpolar 3. N2 P I N 4. PI3

62. 2.2 Intermolecular Forces
A. Types of Forces
are the forces of attraction
intramolecular forces
within molecules
(eg. ionic or covalent bonding)
are the forces of attraction
intermolecular forces
between molecules
they are the
weakest of all forces
responsible for
state, melting point, boiling point etc.
there are three types of intermolecular forces that we will look at:

63. 1. Dipole-Dipole Forces electrostatic force of attraction between the
dipoles of polar molecules
attract the in other molecules and vice versa
slightly negative “poles”
slightly positive “poles” + -

64. 2. Hydrogen Bonding
this is a special type of dipole-dipole interaction that is
very strong
hydrogen bonding is the attraction between a
hydrogen on one molecule which is bonded to O, F or N, to the O, F or N of an adjacent molecule
when hydrogen is bonded to a such as O, F or N, the electrons are pulled from it
highly electronegative element
far away

65. O H since hydrogen doesn’t have any other its is basically exposed
electrons,
proton
this proton is then able to be attracted not only to the
 pole
but also to the
lone pairs O H 

66. 3. London (Dispersion) Forces
the attractive force that occurs between molecules is called
all
London Dispersion force
the result of the electrostatic attraction of
induced dipoles
electrons in atoms and molecules are always in
constant rapid motion
for brief instances, the distribution of electrons becomes which produces very weak
distorted
dipoles

67. this induces a dipole in the when
temporary dipole
like charges repel each other
adjacent molecule
this process throughout the substance, causing that
“disperses”
flickering dipoles
attract each other
even though this force lasts and is , the overall effect in a substance is
only a moment
very weak
significant

68. LD forces are affected by two factors:1.
size of the atoms –
means higher probability of creating
more e
temporary dipoles 2.
shape of the molecule –
the more between molecules, the the force of attraction
contact
higher

69. Scale of Forces Intermolecular Forces (between) Intramolecular Forces
very high
very
low LD DD
network
covalent HB
ionic
covalent
Intermolecular Forces
(between)
Intramolecular Forces
(within)
London Dispersion
metallic ** wide range
Dipole – Dipole
ionic
Hydrogen Bonding
covalent
network covalent
eg) diamond, SiC, SiO2

70. 2.3 Relating Structures and Properties
A. States of Matter – Read p. 72 – 74!!!
the of a substance depends on the between its particles
strength of the attractive forces
state
solids have the forces of attraction, liquids have the and gases have intermolecular attractions between the particles
greatest
very few if any
next greatest

71. B. Melting & Boiling Points
both melting and boiling points are indicators as to the
strength of attractions within or between molecules
when melt or boil, must be broken
metals
metallic bonds

72. when melt or boil, must be broken
ionic bonds
ionic compounds

73. when melt or boil, must be broken
molecular compounds
only intermolecular forces
(covalent bonds DO NOT break)
these forces are much therefore it takes a lot less
weaker
energy to separate the molecules
as you the number of intermolecular forces, the melting and boiling points
increase
increase

74. Order of bp’s
Using the scale of forces you can order compounds based on their relative bp’s
Ex. From Highest to Lowest
Network covalent compound (ex. SiO2)
Ionic compound
Molecular compound with HB, DD, LD
Molecular compound with DD, LD
Molecular compound with LD (if 2 molecular compounds have LD only then bigger molecule or molecule with more electrons has higher bp)

75. Chapter 2: Chemical Bonding 2.1 Three Dimensional Structures
A. Ionic Crystals
ionic compounds have
crystal structure
they form so that are as as possible
oppositely charged ions
close together
this is called a
3-D array of alternating positive and negative ions
crystal lattice

76. sodium chloride since all the in the lattice are the attractive forces
same,
you cannot call them
molecules…
each positive ion is attracted to of the negative ions around it
all
(and vice versa)
the chemical formula is the
lowest whole number ratio
for that type of crystal
eg) NaCl has a 1:1 ratio of Na ions to Cl ions
sodium chloride

77. there are many different and they all depend on the way the ions
crystal shapes
pack together
shape also depends on
the relative size of the ions and the charges on the ions

78. C. Mechanical Properties of Solids
the mechanical properties of solids are determined by the types of bonds in the substance
delocalized e cause bonds to be non-directional, which allows a solid to be malleable and ductile
eg) metals
solids that do not have delocalized e have directional bonds, which causes them to be brittle and hard
eg) ionic compounds

79. D. Conductivity
electric current is the directional flow of electrons or ions
metals are good conductors of electricity because the delocalized valence e are free to move
solid ionic compounds have valence electrons that are held solidly in place therefore they cannot conduct electricity
when ionic compounds melt or dissolve in water, the ions are able to move past one another which allows them to carry an electric current

80. in most molecular compounds, valence electrons are not free to move through the molecule therefore they are not able to conduct electricity
when molecular compounds melt or dissolve in water, they do not form ions and therefore they do not carry an electric current

81. How does a gecko climb?