Transcript The Mole - Aparima

The Mole
Mass
Relative Atomic Mass
Amount
Calculations
Author: J R Reid
Mass
Mass is the fancy scientific name for the amount of
matter or stuff in an object
Mass has the symbol m, and is measured in grams
(g)
Mass is not the same thing as weight
For example: If you have 70kg of mass here on earth and
someone shoots you into space where there is no gravity,
you do not lose any matter (stuff). However, in space you
may weigh nothing at all – but you still have mass.
Weight is measured in Newtons – it is the effect of
gravity on a mass.
Amount
‘Amount’ is the scientific name for the
number of particles in an object. The
particles could be atoms or molecules
Amount has the symbol n, and is
measured in moles (mol)
One mole is equal to 6x1023 particles.
That is
600,000,000,000,000,000,000,000.
This number is called Avogadro’s
number
Atomic Mass
Each element on the periodic table has a different
mass. Hydrogen is the smallest, helium is next…
The mass is caused by the number of protons and
neutrons each element has in its nucleus
The mass number is the amount of grams per mole
an element has:
For example: Uranium can have a mass number of 235
grams per mole. This means that
600,000,000,000,000,000,000,000 Uranium atoms has a
mass of 235 grams. This can be simplified by writing it
like this: 235gmol-1
Relative Atomic Mass
Unfortunately when we look up the periodic table
we find that the mass numbers are not usually
whole numbers. This is because here on Earth each
element can have a number of isotopes. An isotope
is a different arrangement of the nucleus, for
example hydrogen can have a mass number of 1, or
2, or 3. This is because they can have different
amounts of neutrons but still be hydrogen
The Relative Atomic Mass (Ar) on the periodic table
is the mean mass of that element here on Earth
For example: Hydrogen has a Ar of 1.0079 which shows
that most Hydrogen atoms have a mass number of 1 but
there are some with a mass number of 2. The mean
(average) mass number is 1.0079
Relative Molecular Mass
Atoms can bond together to form molecules or
ionic compounds. When this happens we can
add up all the Relative Atomic masses to get a
Relative Molecular mass (Mr) – also called a
molar mass
H2 is two hydrogen atoms combined. The Relative
Molecular mass of H2 is:
2 x 1.0079 = 2.0158
H2O is two hydrogen atoms and one oxygen
combined. The Mr(H2O) is:
2 x 1.0079 + 15.999
NOTE: Molecular and Atomic mass can also
be called Molar mass (M).
A shorthand way of writing “the molar mass of
H2O” is M(H2O)
Calculations
The mass, amount or relative atomic mass can be calculated. This is the
equation:
m = Ar x n
Ar or Mr can be used in this equation
The Ar or Mr can always be worked out from a periodic table
In other words we need two values before we can calculate the missing one:
For example:
m = 100g, Ar = 1.008 gmol-1
n = 100/1.008
= 99.2 mol
In other words 100g of hydrogen atoms is 99.2 moles
Note – all masses MUST be measured in grams before you can use this
equation
This equation can also be written using the following symbol:
m = Mn
(M = molar mass, another way of writing Ar or Mr)
Example: Calculating an
amount
Joan has a 56g mass of copper sulphate. What
amount does she have?
What do we know?
Mass (m) = 56g
Molar mass (Mr) = ?159.5 gmol-1
Amount (n) = ?0.35 mol
We
nowbe
use
Thecan
Mr can
the
equation
to
calculated
if we
calculate
the
have a periodic
amount
table
Cu
m ==Mr
63.5
xn
But first we need
to rearrange it so
n=
= 32.0
m/Mr
S
that it starts with
we add the
nx
=O
56/159.5
4
= 4 x 16.0
“n Now
=
=”64.0
values that we
n = 0.35
mol gmol-1
Total
= 159.5
know…
Chemical Molar mass
Mass
O
100g
O2
5.0g
Amount
Na
1.2mol
NaOH
0.0023mol
Ca(OH)2
1.3kg
KMnO4
100mg
20g
1.25mol
0.38g
0.010mol
Using Mass, Amount and
Empirical Formulae
Sometimes we are given percentage mass and we are asked to
turn it into a empirical formula. This involves a few steps –
turn the mass into amounts, then turn the amounts into a
ratio (empirical formula)
1.
2.
3.
4.
5.
Find the percentage mass for each element (you may be asked to
calculate this i.e. element mass/total mass)
Pretend that you have 100g – now your percentages convert
directly to grams (mass) i.e. 12.5% = 12.5g
Turn the mass into moles i.e. 12.5g of Carbon = 1.04moles (n =
m/Mr)
Find the ratio of the moles of each element (divide each amount
by the smallest number, then check to see if it can be simplified
any further)
Now put the mole ratios into a formula i.e. C:H:O = 1:2:2
therefore the empirical formula is CH2O2
Note: This empirical formula can only be converted into the
molecular formula if you know the total mass number of the
substance
Some Examples Percentages
Lets say “Fred” tries to extract Ca
From CaCO3. If he has 30 grams of
CaCO3 what mass of Ca is present?
1. What percentage of CaCO3 is Ca?
- The M(CaCO3) = 100gmol-1
- The M(Ca) = 40gmol-1
- Therefore 40/100 (40%) is Ca
2. Now if 40% of CaCO3 is Ca then:
- 40/100 x 30g = 12g
Some Examples –
Percentages (Part 2)
1.
2.
Lets say “Fred” now tries to extract Ca From CaCO3. If
he wants 18 grams of Ca what mass of CaCO3 will he
need to use?
What percentage of CaCO3 is Ca?
- The M(CaCO3) = 100gmol-1
- The M(Ca) = 40gmol-1
- Therefore 40/100 (40%) is Ca
Now if 40% of CaCO3 is Ca then:
- 100/40 x 18g = 45g
Note – for the mathematically challenged
- Did you see the difference on this slide?
– Did I want to make the number smaller or larger? For
example, I want an 18g sample of Ca. If I’m extracting it
from CaCO3. The CaCO3 is going to be a larger number
than 18g. Therefore my multiplier needs to be 100/40
(the opposite to the percent)
Balanced Equations
Balanced equations like the one below show
the proportion of amounts being reacted:
2H2 + O2 -> 2H2O
In other words 2 moles of H2 reacts with 1
mole of O2 to make 2 moles of H2O.
These are proportions and can be changed
e.g. what happen if I only have 0.2 moles of
H2?
0.2H2 + 0.1O2 -> 0.2H2O
The proportions of each chemical are still the
same (but each one is now 10x less than
before)
Balanced Equations 2
Here is how we can use balanced equations:
2H2 + O2 -> 2H2O
“Jill” has 18g of H2 – what amount of H2O will
be made?
- First we convert the mass of the H2 into
amount (moles)
n = m/M = 18/2 = 9moles
- Next we work out the proportions in the
balanced equation:
9H2 + 4½O2 -> 9H2O
Note: If we need to, we can now calculate the
other masses using: m = Mn
e.g. m(H2O) = 2 x 9 = 18g
Challenges…
Try these:
Joe burns 100g of carbon in oxygen to
make CO2. What mass of CO2 will he
make?
Julie reacts Sodium with Chlorine gas
(Cl2) to make 20g of NaCl. What masses of
each reactant did Julie use?
Dave extracts 10g of Ca from a 100g
sample of limestone (impure CaCO3).
What percentage of the sample was pure
CaCO3? (What assumptions must we
make?)
See: www.webelements.com