Transcript Ch 10 The Mole

THE
MOLE
Chapter 10: Chemical Quantities
10.1 Measuring Matter
What is a mole? It is
the SI unit that
measures the amount
of substance.
Mass
s
Mole
Volume
Number
of
Particles
Avogadro’s
number
Simply multiply by
23
6.02 x 10
 How
many moles of
magnesium is 1.25 x 1023 atoms
of magnesium?
 How
many atoms are in 2.12
mol of propane? (C3H8)

Multiply by the molecular mass of the
substance.

Molar mass: mass (in grams) of one mole
of the substance

The Mass of 1 mole (in grams)
› Molecular mass, molecular weight, formula mass,
formula weight

Equal to the numerical value of the average
atomic mass (get from periodic table)
1 mole of C atoms
=
12.0 g
1 mole of Mg atoms
=
24.3 g
1 mole of Cu atoms
=
63.5 g

Find the molar mass
(usually we round to the tenths place)
A. 1 mole of Br atoms = 79.9 g
 B. 1 mole of Sn atoms = 118.7 g


Mass in grams of 1 mole is numerically
equal to the sum of the atomic masses
1 mole of CaCl2
= 111.1 g/mol
1 mole Ca x 40.1 g/mol
+ 2 moles Cl x 35.5 g/mol
= 111.1 g/mol CaCl2
 How
many grams are in 9.45
mol of dinitrogen trioxide
(N2O3)?
 How
many moles in 92.2 g of
iron (III) oxide (Fe2O3)?
The Special Number
22.4

volume varies with changes in
temperature and pressure
› usually measured at STP (0oC; 101.3 kPa)
› at STP, one mole of any gas occupies a
volume of 22.4 L
 Determine
the volume, in liters,
of 0.60 mol SO2 gas at STP.
 The density of a gaseous
compound containing carbon
and oxygen is 1.964 g/L at STP.
Determine the molar mass of
the compound.
 Mass
/ molecular mass
 Volume / 22.4
 Number of particles / Avogadro’s
number

How many molecules are in 5.0 g of
N2O5?

What is the volume of 10.0 g of CO2?

What is the mass of 1.5 x 1025 molecules
of C3H8?
 Percent
by mass of each
element in a compound

% mass of element =
Grams of element X X 100
Grams of compound
 9.03
g Mg combine
completely with 3.48 g N to
form a compound. What is
the % composition of this
compound?
 Calculate the % composition
of ethane (C2H6).
 Calculate
the mass of
hydrogen in 350 g C3H8.
Lowest
whole number ratio
of atoms of an element in a
compound.
› May or may not be the same
as the molecular formula!
› If not given an amount, you
may assume it is a 100g.
1.
2.
3.
Find moles of each element
Divide by the smallest number of moles
If not a whole number, multiply to
obtain a whole number.
Give the empirical formula for a compound
which is 25.9% Nitrogen and 74.1% Oxygen.
 Assume it is a 100g sample, so 25.9g N and
74.1g of O
 25.9/14 =1.85 moles of N
74.1/16 = 4.63 moles of O
1.85/1.85 = 1
4.63/1.85 = 2.5
Double both numbers to get a whole number
ratio…
N2O5



Calculate the empirical formula of a
compound that is 79.8% C and 20.2% H.

Calculate the molecular formula of a
compound whose molecular mass is 62
g/mol and empirical formula is CH3O.
Compounds with specific numbers of
water molecules bound to their atoms
are called hydrates.
 In the formula, a dot is used to show that
water is bonded. Use a prefix to name
the hydrate.

› Na2CO3 ∙ 10H2O
Sodium carbonate decahydrate

In order to analyze a hydrate, you must
first find the number of moles of water
associated with one mole of the hydrate.
› Heat the sample to drive off the water, then
you can mass the anhydrous compound
and determine the moles of water.

A mass of 2.50 g of blue, hydrated
copper sulfate (CuSO4· xH2O) is placed
in a crucible and heated. After heating,
1.59 g of white anhydrous copper sulfate
(CuSO4) remains. What is the formula for
the hydrate? Name the hydrate.

An 11.75g sample of a common hydrate
of cobalt (II) chloride is heated. After
heating, 0.0712 mol of anhydrous cobalt
chloride remains. What is the formula
and the name of this hydrate?
THE
END