Transcript The Mole-Ch 11-Chem L1 notes.PART1

The Mole
Chapter 11 – Chemistry L1
LSM High School
Section 11.1: Measuring Matter
Objectives:
 Describe how a mole is used in chemistry
 Relate a mole to common counting units
 Covert moles to number of representative particles and number of
representative particles to moles.
How do Chemists measure
how much of a substance?
Chemists can measure mass or volume or
they can count pieces.
 Chemists can measure mass in grams.
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Chemists can measure volume in liters.
No, not that kind of mole!!!
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Chemists can count pieces in MOLES.
What are MOLES?
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Moles are defined as the number of carbon
atoms in exactly 12 grams of the carbon-12
isotope.
1 mole of _____ = 6.02 x 1023 particles
 Mole: unit = “mol”
 Avogadro’s number
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dozen, baker’s dozen, pi
A Little History
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Amedeo Avogadro was born in
1776 in Turin, Italy.
He went on to study molecular
theory and helped other
scientists distinguish between
atoms and molecules.
Because of his
accomplishments in this field,
the variable that tells the
number of molecules in one
mole was named after him
What are Representative
Particles?
These particles are the smallest pieces of a
substance.
• The types of representative particles that
chemists generally work with are:
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• atoms – the smallest particle of an element
• ions – atoms with positive or negative charges
• molecules – two or more covalently bonded atoms
• formula units – the simplest ratio of ions that make up
an ionic compound
Converting
Moles to Particles
and Particles to Moles
Using Avogadro’s Number as a Conversion Factor:
Practice Problem 1
How many atoms are in 2.50 mol of zinc?
 K:
UK:
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Answer: 1.51 x 1024 atoms Zn
Practice Problem 2
How many molecules of CO2 are there in
4.56 moles of CO2 ?
 K:
UK:
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Answer: 2.75 x 1024 molecules of CO2
 How many atoms is this?
Practice Problem 3
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How many moles of water is 5.87 x 1022
molecules of water?
K:
UK:
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ANSWER: 0.0975 moles of water
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Practice Problem 4
Given 3.25 mol AgNO3, determine the
number of formula units.
 K:
UK:
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ANSWER: 1.96 x 1024 formula units AgNO3
Section 11.2:
Mass and the Mole
Objectives:
• Relate the mass of an atom to the mass of a
mole of atoms.
• Calculate the number of moles in a given mass
of an element, and the mass of a given number
of moles of an element.
• Calculate the number of moles of an element
when given the number of atoms of an element.
• Calculate the number of atoms of an element
when given the number of moles of the element.
Let’s Look at the Periodic
Table!
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Atomic Numbers - always increase across a row.
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Atomic Mass - usually increase across a row
Why do they have decimal values?
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The atomic number is the number of protons in an atom of that
element.
This number identifies it as an atom of a particular element.
The atomic mass (sometimes called average atomic mass) is the
weighted average of the masses of all the naturally occurring
isotopes of that element.
A relative scale: Uses isotope carbon-12 as the standard

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Each atom of carbon-12 has a mass of exactly 12 amu (atomic
mass units)
Ex: One atom of hydrogen-1 has a mass of 1 amu, meaning 1 atom
of hydrogen-1 is one-twelfth the mass of one atom of carbon-12
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The mass in grams of one mole of ANY pure
substance is its molar mass.
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Same value as atomic mass - has units of g/mol
Occasionally referred to as:
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12.01 grams of carbon has the same number of
particles as 1.01 grams of hydrogen and 55.85 grams
of iron.
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Gram atomic mass (gam) – for atoms
Gram molecular mass (gmm) – for molecules
Gram formula mass (gfm) – for formula units (ionic compounds)
The molar mass is found in the periodic table!
Avogadro’s number tells us the number of particles.
Using Molar Mass as a
Conversion Factor:
# of grams
1 mol
or
1mol
# of grams
Practice Problems:
1.
K:
What is the mass, in grams, of 2.34 moles
of carbon?
UK:
28.1 g carbon
2. How many moles of magnesium are in
4.61g of Mg?
K:
UK:
0.190 mol Mg
Section 11.3:
Moles of Compounds
Objectives:
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•
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Recognize the mole relationships shown by a
chemical formula.
Calculate the molar mass of a compound.
Calculate the number of moles of a compound
from a given mass of the compound, and the
mass of a compound from a given number of
moles of the compound.
Determine the number of atoms or ions in a
mass of a compound.
Enough about atoms:
What about compounds?

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The chemical formula for a compound tells us the types of
elements and the number of each element contained in one
unit of the compound.
Ammonia (NH3)
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1 molecule contains:
 1 atom of nitrogen and 3 atoms of hydrogen
Baking soda (sodium hydrogen carbonate,
NaHCO3)
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1 formula unit contains:
 1 atom of sodium, 1 atom of hydrogen, 1 atom of
carbon, and 3 atoms of oxygen
Example Problems:
1. Calculate the number of moles of hydrogen
found in 3.50 moles of NH3.
K:
UK:
10.5 mol Hydrogen
2. Calculate the number of moles of carbon
found in 9.85 moles of C6H12O6 (sugar).
K:
UK:
59.1 mol carbon
The Molar Mass of Compounds
Summary of Getting the Molar
Mass of Compounds:
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The mass of a mole of a compound equals the sum of the
masses of every particle that makes up the compound.
Use the formula to tell you how many of each element that
is in the compound
Use the periodic table to get the masses of each element
Add them all up and you get the molar mass of the
compound in units of g/mol
(NH4)2SO4
Example Problems of
Molar Masses:
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1) What is the molar mass of NH3?
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17.04 g/mol
2) What is the molar mass of Sr(NO3)2?
211.64 g/mol
Example Problems of MoleMass Conversions:
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1) How many moles is 4.56 g of CO2 ?
K:
UK:
0.104 moles CO2
 2) How many moles is 46.8 g of CH4?
K:
UK:
2.92 mol CH4
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3) How many grams is 9.87 moles of H2O?
K:
UK:
178g H2O
 4) How many grams is 0.157 mol Fe2O3?
K:
UK:
25.0 g Fe2O3
Using Molar Volume as a
Conversion Factor:

Molar Volume:
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Standard Temperature and Pressure is 0°C or 273 K and 101.3 kPa or 1 atm
22.4 L
1 mol
for any gas at STP, 1 mol = 22.4 L
or
1mol
22.4 L
Practice Problem:
What is the volume of 1.5 moles of nitrogen gas?
K:
UK:
34 L N2 (17 2-Liter bottles!)
REVIEW:
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What types of particles are contained in
covalent compounds?
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What types of particles are contained in ionic
compounds?
Multi-step Conversions
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You must first convert to moles and then
convert to the desired unit either using molar
mass or Avogadro’s number or molar
volume.
Example Problems:
1. What is the volume of 45.6 g of water
vapor?
K:
UK:
? L H2O(g)
2. How many atoms are in 0.120 kg Ti?
K:
UK:
1.51 x 1024 atoms Ti
4. What is the mass, in grams, of 1.50 x 1015
atoms uranium?
K:
UK:
5.93 x 10-7 g U
3. What is the mass, in grams, of 1.50 x 1015
formula units of NaCl?
K:
UK:
? g NaCl
More Example Problems…
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1) How many molecules in 6.8 g of CH4?
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2a) How many formula units are there in 4.9 g
of NaNO3?
2b) How many ions if the compound is made of
Na+ and NO3- ions?
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