Download UNIT 2 – Chemical Quantities

UNIT 3 –Quantities in Chemical Reactions
THE MOLE!
In chemistry as in other aspects of life it is sometimes more convenient to count in groups
of items rather than count items individually.
Quantity
Amount
In chemistry chemical reactions are __________________________________________.
To count the # of atoms in a chemical reaction individually is inconvenient so chemists
group atoms into groups.
Chemists count in ____________________________________________
________________________________________________________________________
This value is known as the __________________________________________________
The mole is defined as:
________________________________________________________________________
________________________________________________________________________
How big is the Avogadro constant?
Try the Thought lab on pg. 175.
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Canada covers 9976140 km2
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Distance to moon is 384,403 km
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Length of $5 is 152.4 mm
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The worlds oceans hold 1.37x109 km3 water.
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One teaspoon is equal to 5ml
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Average size apple has a mass of 0.15 kg
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Mass of the Earth is 5.9742 × 1024 kg
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Current population of Earth is 6,654,181,537
Converting Moles to Number of Particles
N=
n=
NA =
Problem 1. A sample contains 2.25 moles of carbon dioxide, CO2.
a. How many molecules are in the sample?
b. How many atoms are in the sample?
Problem 2. How many moles of nitrogen dioxide are there in a sample containing 4.35 x
1024 molecules?
Molar Mass
When we measure the mass of a substance we express its mass in grams.
You have already learned that one atom of ______________ has a mass of exactly _____
and that ______________ atoms has a mass of exactly _______.
This means:
__________________________________________________________________
_________________________________________________________________.
This is known as ________________________________________________________.
Molar mass can be applied to any substance and is equivalent to the
_______________________________________________________________________.
Molar mass of Al
Molar mass of HNO3 or MHNO
Using Molar Mass
The following equation will allow you to solve problems involving
________________________________________________________________________
Problem 1.
A flask contains 0.950 mol of carbon monoxide, CO. What is the mass of CO in this
sample?
Problem 2.
How many moles of acetic acid CH3COOH are in a 15.2g sample?
Problem 3.
What is the mass of 6.35 x 1024 molecules of copper (II) nitrate.
Number
of
Particles
Moles
Mass of
substance
Chapter 6 – Chemical Proportions in Compounds
In the late 1700’s Joseph Proust discovered that samples of copper (II) carbonate always
contained the same proportion of Cu, C and O.
This led to the ___________________________________________________________
– the elements in a chemical compound are always present
_______________________________________________________________________
The mass of an element in a compound, expressed as a % of the total mass is known as an
elements ________________________________.
Ex H2O
The law of definite proportions _____________________ that the elements in
compounds are always the same.
CO
CO2
If we state all of the mass percents for a compound we call it
_______________________________________________________________________
A compound with a mass of 48.72g is found to contain 32.69g of Zn and 16.03g of S.
What is the % composition?
Calculating % composition from a formula
Pyrite (aka _________________________) is a compound with the chemical formula
_______________. If you wanted to extract the iron and know how much to expect, you
first need to know the _____________________________.
Calculate the % composition for ________________.
______________________________ C9H8O is the molecule responsible for the
____________________________________________________________________.
Find the % composition of this compound.
Empirical Formula of a Compound
The _____________________________________ of a compound shows the
____________________________________________ of the elements in the compound.
Ex H2O2
The ____________________________________________ of a compound shows the
___________________________________________ of each element in the compound.
The relationship between the 2 types of formulas is:
Determining Empirical Formulas
To determine an empirical formula you need to use the concepts of
________________________________________________________________________
Ex1. Calculate the empirical formula for a compound that is 85.6% C and 14.4% H.
Ex2. The percentage composition of a fuel is 81.7% carbon and 18.3% hydrogen. Find
the empirical formula.
Molecular Formula of a Compound
Being able to identify unknown substances as many uses. Think forensic science.
Since many substances have the same empirical formulas, we need to be able to
determine the _____________________________________ for a substance. To do this
we use _________________________.
Chemists use a ___________________________________ to determine the molar mass
of a substance. Then:
Ex. 1 The empirical formula for ribose (a sugar) is CH2O. The molar mass of ribose is
150g/mol (by experiment). What is the molecular formula of ribose?
Finding Formulas by Experiment
There are many ways to determine formulas by experiment.
The Carbon-Hydrogen Combustion Analyzer
This is a device used to ____________________________________________________.
How it works

Stream of ______________________________________ containing compounds.

CH compound will react completely to produce __________________________

All the H2O is absorbed by ___________.

Final mass – initial mass = mass of H2O. All ____________________________.

Use ______________________ of H in H2O to determine mass of H

All the CO2 is absorbed by __________________. All C ends up here.

Use _____________________________ of CO2 to determine mass of C.

If original compound contained a third element, then:
Ex1. A 1.000g sample of a pure compound containing only C and H was combusted.
0.6919g of water and 3.338g of carbon dioxide were produced.
a. Calculate the masses of C and H in the sample.
b. Find the empirical formula of the compound.
Hydrated Ionic Compounds
Many ______________________ crystallize from an aqueous solution, with H2O
molecules ___________________________________________________________.
These are called ____________________________.
_______________________ have a specific # of _______________________ chemically
bonded to each formula unit.
Chemical Name
Formula
Compounds that have no H2O molecules are called ______________________________
Just remember: when doing calculations you must
________________________________________________________________________
Ex 1.
A 50.0g sample of a hydrate of barium hydroxide Ba(OH)2xH2O contains 27.2g of
Ba(OH)2.
a. Calculate the percent by mass of water in Ba(OH)2xH2O
b. Find the value of x in Ba(OH)2.
Chapter 7 Quantities in Chemical Reactions
Stoichiometry
Stoichiometry is the study of the
________________________________________________________________________
It all begins with ___________________________________________, which are
essential for making calculations related to chemical reactions.
For example, the __________________________________ (used to make ammonia gas)
is based on the following balanced equation:
What does this equation tell you?
1. s
2. s
3. If we multiply each by 12 we could say: 1 dozen molecules of nitrogen gas reacts
with 3 dozen molecules of hydrogen gas to produce 2 dozen molecules of
ammonia gas.
4.
The relationship between moles in a balanced chemical equation are known as MOLE
RATIOS
We use mole ratios to solve problems.
1.
From the equation above, how much ammonia gas would be produced from 3.5
mol of hydrogen gas?
2.
When carbon and oxygen combine 2 reactions are possible. The product can be
CO2 or CO.
a. Write the balanced chemical equations.
b. How much oxygen is needed to react with 0.75mol carbon to produce
CO2?
c. How much oxygen is needed to react with 0.75mol carbon to produce CO?
Mole Ratios can also be used to find the mass ratios for chemical equations.
N2(g) + 3H2(g)  2NH3(g)
Balanced

Equation
# of Particles

Amount

Mass = n x M

Total mass

Solving Stoichiometric Problems
Follow this process:
1.
2.
3.
4.
Ex 1. Mass to Mass
Astronauts must be able to remove the CO2 from their spacecraft. If not hypoxia (lack of
O2) would occur. CO2 is removed using the following chemical reaction: CO2(g) +
2LiOH(s)  Li2CO3(g) + H2O(g) . If an astronaut produces 1.50x103 g of CO2 a day, what
mass of LiOH would be needed per day?
Ex 2. On Feb, 20, 2008 the US Navy shot down a spy satellite in space from a war ship
in the Pacific ocean. They claimed the satellite contained hydrazine (rocket fuel), that
could harm people if the satellite fell back to Earth. When hydrazine is combined with
dinitrogen tetroxide it produces nitrogen gas and water. If 100g of hydrazine is reacted,
how much nitrogen gas would be produced?
Ex 3. During WWI, chlorine gas was used as a weapon. When chlorine gas enters the
lungs it reacts with water and forms corrosive hydrochloric acid and oxygen gas. How
many molecules of chlorine react to produce 5.0g of HCl?
The Limiting Reactant
Chemical reaction equations give the ideal
_______________________________________________________________________.
However, the reactants for a reaction in an experiment are not necessarily a
_______________________________________________________________________.
In a chemical reaction, one of the reactants may ___________________ when the
reaction is finished. These are called ___________________. The reagent that is
______________________________ or reacted is called the ______________________,
because its quantity limit the amount of products formed.
Think of when you put gas in your car:
2C8H18 + 25O2  16CO2 + 18H2O
Ideally, ____________ of oxygen react with _________________ of octane. But in
reality __________________________________________so ___________ would be the
________________________________________. When you run out of gas
____________________________
In stoichiometric problems, _________________________________________________
________________________________________________________________________
________________________________________________________________________
Ex 1. Glucose reacts with oxygen gas to produce carbon dioxide and water during
cellular respiration. If 1.25 g of glucose reacts with 7.51g of oxygen gas, find the
limiting reactant.
Now that you know the limiting reactant above, how much carbon dioxide would be
produced?
Ex 3.
Percentage Yield
Chemists use _____________________ to predict the ________________________ that
can be __________________ from a chemical reaction.
This amount is called the __________________________________________.
This amount is _____________________________. The ____________________
obtained in an experiment is called the ______________________________.
The ____________________ is usually less then the _______________________ and for
a variety of reasons. For example:
a.
b.
c.
d.
_________________________________________________ compares the
______________________________________________________________ (actual) to
the ____________________________________________ (theoretical).
Ex1. When 75g of nitrogen gas reacts with sufficient hydrogen gas, the theoretical yield
of ammonia is 9.10g. If 1.72 g of ammonia is obtained what is the percentage yield?
Ex 2. When Calcium carbonate is decomposed to calcium oxide and carbon dioxide there
is a 92.4% yield of calcium oxide. How many grams of calcium oxide can you expect if
12.4g of calcium carbonate is heated.
Many times substances may not be pure. The substance might contain
_______________________. If gold is 99% pure then it must contain 1% other
substances. ________________________________________ describes what proportion
_________________ of a sample is composed of a specific compound or element.
Ex 1. Your have a 13.9 g sample of “fools gold”. You heat the sample in oxygen and
produce 8.02g of iron (III) oxide. What % purity of iron pyrite was in the original
sample?
4FeS2 + 11O2  2Fe2O3 + 8SO2