CHAPTER 10
Chemical Quantities
Before We Begin…
a)
b)
c)
d)
I can write numbers in scientific notation.
I can write numbers in standard notation.
I can multiply numbers written in scientific notation.
I can divide numbers written in scientific notation.
Before We Begin…


We need to review some scientific notation.
Scientific notation is a way of writing very large
and very small numbers.
How to Write Numbers in Scientific
Notation

Always written as a coefficient multiplied by 10
raised to a power.
3.5 x 1034
coefficient
power
Examples:
Write the following in scientific notation:

a)
234560000
b)
0.00056974
c)
8524000000
d)
0.000000044258
How to Multiply in Scientific Notation

To multiply numbers written in scientific notation you
multiply the coefficients and add the powers.
(2.35x1014) x (3.25x10-23)
Multiply
Add
(2.35x3.25) x 1014+-23
How to Multiply in Scientific Notation

To multiply numbers written in scientific notation you
multiply the coefficients and add the powers.
(2.35x1014) x (3.25x10-23)
Multiply
Add
Answer = 7.64x10-9
Examples:

Multiply the following numbers:
a)
(1.23x104) x (4.56x107)
b)
(7.89x10-1) x (1.23x1010)
c)
(4.56x107) x (7.89x10-10)
d)
(1.23x10-11) x (4.56x10-23)
How to Divide in Scientific Notation

To divide numbers written in scientific notation you
divide the coefficients and subtract the powers.
(2.35x1014) ÷ (3.25x10-23)
Divide
Subtract
 2 . 35 
14    23 

 x10
 3 . 25 
How to Divide in Scientific Notation

To divide numbers written in scientific notation you
divide the coefficients and subtract the powers.
(2.35x1014) ÷ (3.25x10-23)
Divide
Subtract
Answer =0.72x1037
Examples:

Divide the following numbers:
a)
(1.23x104) ÷ (4.56x107)
b)
(7.89x10-1) ÷ (1.23x1010)
c)
(4.56x107) ÷ (7.89x10-10)
d)
(1.23x10-11) ÷ (4.56x10-23)
Section 1
The Mole: A Measurement of Matter
Section 1 Learning Targets
10.1.1 – I can describe methods of measuring the
amount of something.
10.1.2 – I can define Avogadro’s number as it relates
to a mole of a substance.
10.1.3 – I can distinguish between the atomic mass of
an element and its molar mass.
10.1.4 – I can describe how the mass of a mole of a
compound is calculated.
Measuring Matter

You often measure the amount of something by one
of three different methods – by count, by mass, and
by volume.
Example:

If 0.20 bushel is 1 dozen apples and a dozen
apples has a mass of 2.0kg, what is the mass of
0.50 bushel of apples?
What Is a Mole?


Mole (mol) – 6.02x1023 representative particles of
that substance (SI unit for measuring the amount of
something).
Avogadro’s number - 6.02x1023 named after
Amadeo Avogadro di Quarenga (1776-1856)

A mole of any substance contains Avogadro’s
number of representative particles, or 6.02x1023
representative particles.
Converting Number of Particles to
Moles

You can use Avogadro’s number as a conversion
factor.
Example:

How many moles is 2.80x1024 atoms of silicon?
Converting Moles to Number of
Particles

The reverse also works.
Example:

How many molecules are in 5.6 moles of NO2?
The Mass of a Mole of an Element



The atomic mass of an element expressed in grams
is the mass of a mole of the element.
Molar mass – the mass of a mole of an element.
Find the element on the periodic table and the mass
that’s listed is the mass of one mole.
The Mass of a Mole of a Compound


To calculate the molar mass of a compound, find the
number of grams of each element in one mole of
the compound.
Then add the masses of the elements in the
compound.
Example:

What is the mass of 1.00 mol of sodium hydrogen
carbonate?
Section 2
Mole-Mass and Mole-Volume Relationships
Section 2 – Learning Targets
10.2.1 – I can describe how to convert the mass of a
substance to the number of moles of a substance,
and moles to mass.
10.2.2 – I can identify the volume of a quantity of
gas at STP.
The Mole-Mass Relationship

Use the molar mass of an element or compound to
convert between the mass of a substance and the
moles of a substance.
Example:

Find the mass, in grams, of 4.52x10-3mol of C20H42.

The reverse is also true.
Example:

Calculate the number of moles in 75.0g of
dinitrogen trioxide.
The Mole-Volume Relationship

Avogadro’s hypothesis – states that equal volumes
of gases at the same temperature and pressure
contain equal numbers of particles.

Standard temperature and pressure (STP) – means
a temperature of 0°C and a pressure of 101.3kPa
or 1 atmosphere (atm).


At STP, 1 mole or 6.02x1023 representative
particles, of any gas occupies a volume of 22.4L
Molar volume – the 22.4L of a gas.
Calculating Volume at STP

22.4L = 1 mol at STP provides a nice conversion
factor.
Example:

What is the volume of 3.70 mole N2 at STP?
Example

How many moles are in 102 L of carbon dioxide, CO2?
Calculating Molar Mass from Density

Different gases have different densities and is
usually measured in g/L so we can calculate
different things using density as a conversion factor.
Example:

A gaseous compound composed of sulfur and oxygen,
which is linked to the formation of acid rain, has a density
of 3.58 g/L at STP. What is the molar mass of this gas?
The Mole Road Map

A helpful tool to figure out easily which conversion
factor to use.
This can also be found
on page 303 in your
Chemistry book