A_-_102_F11_Notes

A_-_102_F11_Notes - A)
Introduction
to
Chemistry


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Unformatted text preview: A)
Introduction
to
Chemistry
 Petrucci
et
al.,
10th
Edition:
1.4‐1.5,
2.7‐2.8,
3.1‐3.3,
4.1‐4.5
 
 Note
that
this
is
not
a
comprehensive
introduction
but
rather
highlights
the
minimum
required
 background
for
the
course.

If
any
of
the
information
here
is
not
familiar
to
you,
read
the
sections
in
 the
textbook.

If
you
require
more
assistance,
please
contact
your
instructor
and/or
the
TAs.
 
 Dalton’s
Atomic
Theory:
 • An
atom
is
a
minute,
indivisible
particle
that
is
the
smallest
unit
of
a
chemical
element.
 • Atoms
are
neither
created
nor
destroyed
in
chemical
reactions.
 • For
a
given
element,
atoms
have
the
same
properties
 • For
atoms
of
different
elements,
these
properties
vary
 • Atoms
combine
in
fixed
proportions
to
form
compounds
 
 
 Physical
VS
Chemical
Change
 • Physical
Change
 • Sample
may
go
from
a
liquid,
to
a
solid,
or
to
a
vapour,
or
to
some
combination
 thereof
 • Chemical
make‐up
unchanged
 • Chemical
Change
or
Reaction
 • A
change
in
composition
occurs
 • Atoms
are
added
to
or
lost
from
the
initial
sample
 
 Law
of
Conservation
of
Mass
 The
total
mass
of
substances
present
before
a
chemical
reaction
is
the
same
as
the
 total
mass
after
the
reaction.
 
 
 Compounds
and
Molecules
 • A
molecule
is
the
smallest
unit
that
has
the
same
proportion
of
elements
as
the
overall
 compound
 • For
example:

a
water
molecule
(H2O)
is
composed
of
two
hydrogen
atoms
(H)
and
 one
oxygen
atom
(O)
 • Table
salt
is
sodium
chloride:
 • Formula
NaCl
(the
ratio
of
sodium
to
chlorine
is
1:1)
but
it
is
not
possible
to
 identify
a
molecule
of
sodium
chloride
 
 Law
of
Constant
Composition
 All
samples
of
a
compound
have
the
same
proportion
by
mass
of
the
constituent
 elements.
 ChE102
Fall
2011
Class
Notes
‐
Introduction
 A ‐1 
 Isotopes
 • All
atoms
of
the
same
element
have
the
same
number
of
protons,
but
they
may
have
 different
numbers
of
neutrons.

 • Atomic
number
is
what
controls
most
of
the
properties
of
the
atom
 • For
example,
chlorine
always
has
17
protons
but
it
commonly
has
either
18
or
20
 neutrons.
 • Therefore,
there
are
two
isotopes
of
chlorine,
both
with
atomic
number
17
but
 having
mass
numbers
of
35
and
37.
 
 
 Ions
 Atom
is
neutral
and
has
the
same
number
of
protons
(positive
charge)
as
electrons
 (negative
charge)
 Ion
is
an
atom
that
has
gained
electrons
(____________
=
negative
charge)
or
lost
electrons
 (______________
=
positive
charge)
 • Therefore,
an
ion
has
a
charge
 For
example:

 
 
 Chemical
Bonding
 • Minimal
knowledge
required
for
this
course:
 • Elements
combine
to
form
compounds
by
bonding
 • Bonding
involves
electrons
 • In
a
covalent
bond,
one
or
more
atoms
share
one
or
more
electrons.
 • For
example,
in
one
molecule
of
methane
(CH4),
one
atom
of
carbon
shares
 electrons
with
four
atoms
of
hydrogen
 
 • In
an
ionic
bond,
one
atom
donates
one
or
more
electrons
to
another
atom:
 • For
example,
in
sodium
chloride
the
sodium
Na
donates
an
electron,
becoming
Na+
 • The
chlorine
accepts
it,
becoming
Cl–
 • The
formula
for
sodium
chloride
isNaCl,
but
it
does
not
exist
as
a
molecule.
 Instead,
it
forms
an
ordered
crystal
of
alternating
Na+
and
Cl–
ions.
 
 
 Polyatomic
Ions
 • Two
or
more
atoms,
joined
by
covalent
bonds,
may
form
a
charged
polyatomic
ion.

 For
example:


 Ammonium
 
 Carbonate
 
 Sulphate
 ChE102
Fall
2011
Class
Notes
‐
Introduction
 A ‐2 
 Identifying
Compounds
 • An
ionic
compound
is
usually
a
metal
combined
with
a
non‐metal.
 • One
of
the
parts
may
be
one
the
polyatomic
ions
 For
example:

 
 
 • A
covalent
compound
tends
not
to
consist
of
a
metal
with
a
non‐metal
 • Large,
organic
(carbon
containing)
molecules
are
usually
covalent.

 For
example:

 
 
 Dissolution
 • When
an
ionic
compound
dissolves
in
water,
it
(wholly
or
partially)
dissociates
into
 ions.

(www.northland.cc.mn.us/biology/Biology1111/animations/dissolve.html)
 For
example:

 
 
 • A
covalent
compound
does
not
dissociate
 For
example:

 
 
 Composition
 The
empirical
formula
is
the
simplest
possible
formula
for
a
particular
compound
 • Gives
the
relative
number
of
atoms
in
a
molecule
 • Different
substances
can
have
the
same
empirical
formula
 • For
example,
empirical
formula
of
hexene
is
CH2
 The
molecular
formula
gives
the
actual
number
of
atoms
in
a
molecule

 • Molecular
formula
can
be
obtained
from
the
empirical
formula
if
the
molar
mass
of
 the
compound
is
known
 • For
example,
molecular
formula
of
hexene
is
C6H12
 
 
 Question
A1
 Choose
the
INCORRECT
statement.

 A)
A
molecule
is
a
group
of
bonded
atoms
that
exist
as
an
entity.
 B)
Ionic
compounds
result
from
combinations
of
metals
and
non‐metals.
 C)
An
anion
is
a
positive
ion.

 D)
The
empirical
formula
is
the
simplest
formula
of
the
ratio
of
atoms.
 E)
The
molecular
formula
is
the
listing
of
the
atoms
in
an
actual
molecule.

 ChE102
Fall
2011
Class
Notes
‐
Introduction
 A ‐3 
 Unit
Conversions
 Mass
 • Metric
units

kg
 • American
Engineering
units

lbm

 Weight
=
mass
x
acceleration
=
force
 • Metric
units:

1
N
=
1
kg
m/s2
 • American
Engineering
units:

1
lbf
=
32.2
lbm
ft/s2
 
 
 Temperature
Conversions
 • To
convert
a
temperature
from
one
scale
to
another,
use
the
following
equations:
 T(K)
=
T(°C)
+
273.15
 T(°R)
=
T(°F)
+
459.67
 T(°F)
=
1.8
T(°C)
+
32
 T(°R)
=
1.8
T
(K)

 
 250
 degrees
Farenheit
 200
 150
 100
 50
 0
 0
 20
 40
 60
 80
 100
 120
 degrees
Celcius
 • 
 
 
 
 
 To
convert
changes
in
temperature,
use
the
following
relationships,
which
give
the
 rate
of
change
per
degree
of
temperature
(aka
slope
of
the
line):
 o o 1C1F , , 1 K 1 oR 1K , 1K o , 1C o , 1C 1.8 o R 1.8 o F 1.8 o F 1.8 o R € ChE102
Fall
2011
Class
Notes
‐
Introduction
 A ‐4 
 Question
A2
 On
a
very
cold
day
in
Alaska,
the
temperature
was
‐40°F.

True
or
false:

the
temperature
 in
degrees
Celsius
on
the
same
day
was
also
‐40.
 
 
 Density
 • Density
is
the
mass
per
unit
volume
(e.g.
g/mL,
____________)
 • Allows
conversion
between
volume
and
mass
of
a
substance
or
solution
 
 
 Question
A3
 1.00
oz
of
gold
(density
=
19.3
g/mL)
canbe
hammered
into
a
sheet
that
would
cover
a
 100
ft2
area.

How
thick
would
such
a
sheet
be
in
centimeters?

 
 
 Percent
Composition
 • Percent
from
the
Latin
“Per
centum”
 • means
per
100
of
something
 • Percent
composition
can
be
given
in
terms
of
mass
(sometimes
“weight”
is
used),
 volume,
mols,
etc
 • To
use
this
information,
convert
the
percentage
into
a
fraction
or
ratio
 • A
compound
that
is
7.2
mol%
C
could
also
be
written
as
 mols C 
 mols compound 
 
 Question
A4
 € A
fertilizer
contains
21%
nitrogen
by
mass.

What
mass
of
this
fertilizer,
in
kilograms,
is
 needed
for
an
application
requiring
225
g
of
nitrogen?
 
 
 Avogadro’s
Number
 • The
number
of
carbon
atoms
in
exactly
12.0
g
of
12C
is
called
Avogadro’s
number
(NA).

 
 
 NA
=
6.022
x
1023
 • The
mole
is
a
measurement
of
quantity:
 
 
 1
___‐mol
=
6.022
x
1023
units
 
 
 Molar
Mass
 • Molar
Mass
is
the
mass
of
one
mol
of
a
substance

 • 6.022
x
1023
atoms
(1
mol)
of
C
weigh
12
g
and
therefore
molar
mass
of
C
is
12
g/mol.
 • 1
g‐mol
of
C
=
12
g
of
C
 • 1
lb‐mol
of
C
=
12
lb
of
C
 ChE102
Fall
2011
Class
Notes
‐
Introduction
 A ‐5 
 Question
A5
 Determine
the
name
of
the
mystery
element
if
a
2.80
x
1022
atom
sample
of
the
element
 has
a
mass
of
5.23g.
 
 
 Molar
Mass
of
Compounds
 • Calculated
from
the
molecular
formula
 • For
example:
Water
(H2O)
 • The
molar
mass
of
H
is
1
g/mol
and
O
is
16
g/mol
 • 



 • Therefore,
the
molar
mass
of
H2O
is
18
g/mol


 
 
 Question
A6
 You
work
for
a
copper
mining
company
and
are
trying
to
decide
which
mine
to
buy.

One
 mine
is
rich
in
chalcolite
(Cu2S)
and
the
other
is
rich
in
Athabascaite
(Cu5Se4).

How
 much
more
copper
(in
g)
can
be
extracted
from
the
chalcolite‐containing
mine
per
kg
of
 mineral?
 
 Composition
of
Compounds
 • Calculating
percent
composition
 • Determining
empirical
formulas
for
new
compounds
using
combustion
analysis
 • Based
on
the
knowledge
of
the
products
made
when
a
substance
is
burned
(reacted
 with
oxygen)
 
 
 Question
A7
 A
certain
compound,
used
as
a
welding
fuel,
contains
only
carbon
and
hydrogen.

 Burning
a
small
sample
of
the
compound
completely
in
oxygen
only,
gives
3.38
g
of
 carbon
dioxide,
0.692
g
of
water,
and
no
other
products.

What
is
the
empirical
formula
 of
the
welding
fuel?
 
 
 Question
A8
 Air
contains
approximately
79
mol%
N2
and
21
mol%
O2.

What
is
the
average
molecular
 mass
of
air?

 ChE102
Fall
2011
Class
Notes
‐
Introduction
 A ‐6 
 Chemical
Reaction
Definitions
 • A
Chemical
Reaction
is
a
process
by
which
a
chemical
change
occurs
 • Reactants
–
substances
that
are
transformed
 • Products
–
substances
that
are
produced
 • Evidence
of
a
chemical
reaction:
 • 

 • 

 • 

 • 

 
 
 Chemical
Equations
&
Stoichiometry
 • Chemical
equations
describe
the
changes
that
occur
in
chemical
reactions:

 N2O5
→
NO2
+
O2
 • Chemical
equations
must
be
balanced,
so
that
the
number
of
atoms
on
the
reactant
 side
are
equal
to
the
number
of
atoms
on
the
product
side:
 
 
 • Stoichiometric
chemical
equation
describes
only
the
net
changes
taking
place
in
a
 reaction
and
does
not
necessarily
correspond
to
what
occurs
at
the
molecular
level
 
 
 The
Mechanistic
Equation
 • Mechanistic
equations
describe
the
steps
required
to
go
from
reactants
to
products
at
 the
molecular
level

 • For
the
above
equation,
a
possible
mechanistic
equation
could
be:

 2
N2O5
 →


2
NO2
+
2
NO3
 NO3
+
NO2
 →


NO
+
NO2
+
O2
 NO3
+
NO

 →


2
NO2

















.
 2
N2O5
 • →


4
NO2
+
O2




 Mechanistic
equations
must
add
up
to
the
stoichiometric
chemical
equation.

 ChE102
Fall
2011
Class
Notes
‐
Introduction
 A ‐7 
 Stoichiometric
Calculations
 • Stoichiometric
relationships
are
used
to
calculate
the
proportions
of
atoms
in
a
 molecule
or
the
proportions
of
products
and
reactants
in
a
given
chemical
reaction

 • For
the
reaction:


aA
+
bB
→
cC
+
dD
 
 
 • 

 
 For
example:

2N2O5
→
4NO2
+
O2

 2
moles
N2O5
:
4
moles
NO2
:
1
mole
O2
 
 
 
 Question
A9
 Sodium
bromide,
used
to
produce
silver
bromide
for
use
in
photography,
can
be
made
as
 follows:
 

Fe
+
Br2
→
FeBr2
 

3
FeBr2
+
Br2
→
Fe3Br8
 

Fe3Br8
+
4
Na2CO3
→
8
NaBr
+
4
CO2
+
Fe3O4
 How
many
kilograms
of
iron
are
consumed
to
produce
2500
kg
of
NaBr?

 
 
 Question
A10
 A
0.578
g
sample
of
pure
tin
(Sn)
is
treated
with
gaseous
fluorine
until
the
weight
of
the
 resulting
compound
is
constant
at
a
value
of
0.944
g.

What
is
the
empirical
formula
of
the
 compound
formed
during
the
reaction?

 
 
 Chemical
Reactions
in
Solution
 • A
solution
is
a
homogeneous
system
that
contains
two
or
more
substances
 • • The
solvent
is
the
major
component
of
the
solution
 • Aqueous
is
used
for
solutions
where
the
solvent
is
water,
but
not
how
much
water.
 The
solute(s)
is/are
the
minor
component(s)
of
the
solution.

 ChE102
Fall
2011
Class
Notes
‐
Introduction
 A ‐8 
 Molarity
&
Molality
 • Concentration,
or
molarity,
is
a
solution
property
defined
as
the
number
of
moles
of
 solute
per
liter
of
solution:

 
 
 
 • Or
molality,
which
eliminates
temperature
effects:
 
 
 
 
 Mass
&
Mol
Fractions
 • Mass
fraction,
which
is
also
sometimes
converted
to
mass
percent
by
multiplying
by
 100%,
is:

 
 
 • Similarly,
mole
fraction
is:
 
 
 
 Question
A11
 A
solution
is
prepared
by
dissolving
22.4
g
of
MgCl2
in
0.200
L
of
water.

Taking
the
 density
of
pure
water
to
be
1.00
g/cm3
and
the
density
of
the
resulting
solution
to
be
 1.089
g/cm3,
calculate
the
molarity
of
MgCl2
in
this
solution.
 
 
 Question
A12
 An
excess
of
aluminum
foil
is
allowed
to
react
with
225
mL
of
an
aqueous
solution
of
HCl
 that
has
a
density
of
1.088
g/mL
and
contains
18%
HCl
by
mass.

What
mass
of
H2
is
 produced
in
the
reaction?



 2
Al
+
6
HCl
→
2
AlCl3
+
3
H2
 
 
 Dilutions
 
 
 [Petrucci
et
al.,
Figure
4­6]
 • Visualize
dilutions
as
the
transfer
of
a
certain
number
of
solute
moles
between
two
 vessels
 • The
number
of
moles
of
a
particular
solute
is
constant
between
the
sample
withdrawn
 from
the
initial
solution
and
in
the
final
solution:
ni
=
nf
 ChE102
Fall
2011
Class
Notes
‐
Introduction
 A ‐9 
 Question
A13
 What
mass
of
potassium
hydrogen
carbonate
(KHCO3)
is
needed
to
prepare
500mL
of
a
 0.100M
aqueous
solution?


 How
much
water
do
you
need
to
add
to
this
solution
if
you
were
to
make
it
0.0400M?
 
 
 Limiting
Reactant
 • Suppose
arbitrary
amounts
of
reactants
are
mixed
and
allowed
to
react.
 • The
limiting
reactant
is
used
up
first
 • Everything
left
at
the
end
is
present
in
excess
 • Test
your
skills:
chemistry.csudh.edu/lechelpcs/limitreagentcsn7.html
 
 
 Question
A14
 If
1.00
kg
each
of
PCl3,
Cl2,
and
P4O10
are
allowed
to
react,
how
many
kilograms
of
POCl3
will
 be
formed,
given
the
following
unbalanced
reaction?
 PCl3
(l)
+
Cl2
(g)
+
P4O10
(s)

POCl3
(l)
 
 
 Theoretical,
Actual
and
Percent
Yield
 • Theoretical
yield
is
the
amount
of
product
possible
if
everything
reacted
ideally
 • However,
we
get
an
actual
yield
because
nothing
goes
exactly
as
planned
 • Why
doesn’t
theoretical
yield
=
actual
yield?
 o 


 o 


 o 


 
 
 
 
 Question
A15
 The
iron
oxide
Fe2O3
reacts
with
carbon
monoxide
to
produce
iron
and
carbon
dioxide:
 Fe2O3
+
3
CO
→
2
Fe
+
3
CO2
 The
reaction
of
450g
of
Fe2O3
with
excess
CO
yields
275g
of
iron.

Calculate
the
 percentage
yield
of
iron
(Fe)
in
this
reaction.
 ChE102
Fall
2011
Class
Notes
‐
Introduction
 A‐10
 ...
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This note was uploaded on 12/28/2011 for the course CHE 102 taught by Professor Simon during the Winter '08 term at Waterloo.

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