Stoichiometry

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
chemical bonding
categorize
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
Hg
metals (except ) in pure form are
?
solids
3
combine with one
another to form
non-metals
?
solids, liquids or gases
never
the only elements that are found
in
in nature are the
?
combined form
noble gases
4

• atoms
  in such a way that they
  create a

gain, lose or share electrons
full outer energy level
octet rule
called the

• 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

guideline

not all elements follow it at all times
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
transition metals

• for many , the
  number of valence electrons

is not as predictable it depends
on the
environment
around the ion
eg)
iron can be Fe3 or Fe2
ion charge

number of valence e?

• the can be used to
  determine

eg)
Fe3 had 3 valence electrons
7
B. Electron Dot Diagrams

• you cant see atoms and electrons, therefore it
  is convenient to
  to show the structure and formation of

draw models
chemical bonds
electron dot diagram

• an
  is one such model

• 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

start pairing up the e?
top
of the symbol and
?
?
?
?
Si
Na
Al
Ca
?
?
?
?
?
?
?
?
?
?
?
?
?
?
Ar
?
?
?
O
Cl
?
?
?
?
?
P
?
?
?
?
?
?
?
?
?
?
9
full

• a orbital is called a
  and is
  (at this level)

lone pair
not involved in bonding

• a orbital contains a

bonding electron
half full
?
?
lone pairs
?
?
O
?
?
bonding e?
bonding capacity

• the of
  an atom is the maximum number of
  that it can form
  (equals the number of )

single covalent bonds
bonding e?
10
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
C. Ionic Bonding

• an is the

electrostatic attraction between oppositely
charged ions
ionic bond
metals
three or
fewer valence e?

• most have

• they tend to these electrons and
  become

lose
positive ions (cations)

Na
Na
12

more than four valence e?

• most have

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
MgO
15
Examples
F
Ca
F
CaF2
K
S
K
K2S
16
Fe2O3
Mg3N2
17

• notice the following about the diagrams

no valence electrons

• the metal has
  (since they them)
• the non-metal has the valence level
• both ions have
  and the
• charges
  charges

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
covalent bonds

• compounds containing
  are called

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
22
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
(No Transcript)
24
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
25
Try These
Draw the structural formula for each of the
following
1. HCl
4. NBr3
2. CH4
5. C2H4
3. F2
6. N2
26
(No Transcript)
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

DO NOT
covalent

ionic bonds
other metal atoms

• in all the
  atoms

metallic bonding

share all the valence e?

• 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

metallic bond

electrostatic force of attraction between the
positive metal ions and the negative sea of
electrons

• this theory helps explain the
  of metals

properties
eg) good conductors of electricity and heat,
ductility, malleability
29
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 its nucleus is at attracting another
atoms valence e?
32
higher

• electronegativity means

greater attraction
(affinity)

• trend on periodic table electronegativity
  
  and

decreases down group

increases across period

• 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
Period 2 Elements
3 p
6 p
9 p
1 valence e?
4 valence e?
7 valence e?
Li
C
F
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

distance

• the between the charges the
  attractive force between opposite charges
  with the
  between them

decreases
square of the distance

• the of the charges
  the attractive force is
  to the

magnitude
directly proportional
amount of charge
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
inner levels of e?
shielded
small
e?
39
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
gain e?
electronegativities

attract each other

• the two ions that are formed will

and form a
chemical bond
41
non-metallic atoms

• covalent bonds occurs between

• 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
  electronegativities share electrons,
  the sharing is

different
no longer equal
higher

• the element with the
  electronegativity pulls the

e? closer to itself
43

• 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
Bond Dipole Arrows
arrow points towards element with higher
electronegativity (?-)
at the end that is ?
H F
?
?-
45
Try These
Draw the bond dipole arrow, label the ? and ??
ends, and state the bond type (polar, nonpolar,
ionic)
0.4
?
?-
6. C H
polar
1. H H
nonpolar
0.8
7. Cl Cl
?
?-
nonpolar
2. N H
polar
1.3
2.0
?
?-
8. Si Cl
?-
polar
?
3. B F
polar
0.8
1.2
?
?
?-
?-
4. S O
9. O H
polar
polar
5. P H
10. Na Cl
nonpolar
ionic
46

• you can use the difference in electronegativity
  between two atoms to determine

 
bond character
Difference in Electronegativity
3.3
1.7
0.5
0
slightly polar covalent
mostly ionic
polar covalent
non-polar covalent
47
bond classification
is not simple

• as you can see,

• bonding is considered a
  and there is

continuum
no clear distinction
between

ionic and covalent bonding
   
48
Chapter 2 Chemical Bonding
2.1 Three Dimensional Structures
A. Ionic Crystals

• ionic compounds have

crystal structure
oppositely charged ions

• they form so that
  are as
  as possible

close together

• this
  
  is called a

3-D array of alternating positive and
negative ions
crystal lattice
49

• since all the
  in the lattice are the

attractive forces
molecules
same,
you cannot call them
all
each positive ion is attracted to of the
negative ions around it
(and vice versa)

• the chemical formula is the

lowest whole number ratio
for that type of crystal
eg) NaCl has a 11 ratio of Na ions to Cl ions
sodium chloride
50

• 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
51
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
52
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
BPs

• shape is determined around the

central atom
53

• shapes can be classified into five categories

1.
linear
central atom is bonded to
and has lone
pairs, or there is only two atoms in the molecule

two other atoms
zero
eg) CO2(g), HCN(g), HCl(g)
54
2.
trigonal planar
central atom
is bonded to and has
lone pairs
three other atoms
zero
eg) CH2O(l)
3.
tetrahedral
central atom is
bonded to and
has lone pairs

four other atoms
zero
eg) CH4(g)
55

three other atoms
4.
pyramidal
central atom is bonded
to and has
lone pair
one
eg) NH3(g)
5.
bent
central atom is bonded to
and has either
lone pairs

two other atoms
one or two
eg) H2O(l), HNO(g)
56

• 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)
3 - 1
4 - 0
pyramidal
tetrahedral
57
Code Shape Example





tetrahedral
4 0
CH4
3 0
trigonal planar
CH2O
3 1
pyramidal
NH3
2 1
bent
HNO
2 2
bent
H2O
all other codes are
linear
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
cancel
opposite in direction
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
polar
polar
4. C2HI
3. C2H2
np
np
polar
nonpolar
61
Try These
1. HF
2. CH4
polar
nonpolar
3. N2
4. PI3
nonpolar
polar
5. C2H6
nonpolar
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
slightly negative poles

• 
  attract the
  in other molecules and vice
  versa

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
electrons,

• since hydrogen doesnt have any other
  its
  is basically exposed

proton

• this proton is then able to be attracted not only
  to the

?? pole
but also to the
lone pairs
66
3. London (Dispersion) Forces
all

• the attractive force that occurs between
  molecules is called

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
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

state

strength of the
attractive forces

• solids have the forces of
  attraction, liquids have the
  and gases have
  intermolecular attractions between the particles

greatest
next greatest

very few if any
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 compounds

ionic bonds
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 bps

• Using the scale of forces you can order compounds
  based on their relative bps
• 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
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
76
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

77

• 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

78
How does a gecko climb?