Download Slide 1

2.C – Conserving
Matter
Objectives
State and apply the Law of Conservation of
Matter.
 Learn how to balance chemical equations.
 Explain the mole concept.
 Calculate molar mass of a compound.
 Calculate percent composition of elements.
 Distinguish between renewable and
nonrenewable resources.

C.1: Keeping Track of Atoms

Complete the reading guide handout
C.1: Keeping Track of Atoms
In some chemical reactions, matter seems to be
created.
 When a nail rusts.
 In other reactions, matter seems to disappear.
 When paper burns and apparently vanishes
 However, neither creation or destruction of
matter occurs.
 Matter may undergo physical or chemical
changes.

C.1: Keeping Track of Atoms
In a car engine gasoline is burned. What
happens to the molecules of gasoline?
 Gasoline is made up of carbon and hydrogen
atoms (C and H atoms)
 When gasoline burns these atoms react with
oxygen atoms in air to form carbon dioxide
(CO2), carbon monoxide (CO) and water
(H2O).
 The original atoms of gasoline are not
destroyed but become rearranged.


Basic chemical equation:
CxHx + O2
reactants
CO2 + CO + H2O
products
Molecules can be converted or
decompose by chemical reactions; but
the atoms remain.
C.1: Keeping Track of Atoms
Law of Conservation of Matter: Matter is
neither created nor destroyed.
 Since chemical reactions cannot create or
destroy atoms, chemical equations
representing the reactions must always be
BALANCED.
 This means that the number of atoms of
each element is the same on the reactant
and product sides of a chemical equation.

Example:
The burning of coal is a reaction
where carbon reacts with oxygen to
produce carbon dioxide. The
number of carbon and oxygen
atoms is the same on both sides of
the equation.
C
+
O2
CO2
Atomic Perspective:
C
1 Carbon atom
O2
1 oxygen molecule


CO2
1 carbon dioxide molecule
1. What are the reactants in this chemical equation?
2. What are the products in this chemical equation?
3. Are there the same number of atoms on both sides of the equation?
a. Were any atoms destroyed or created?
b. Was the Law of Conservation of Matter maintained?
Atomic Perspective:
Cu (s) + O2 (g)  CuO (s)
From the burning of copper lab:
Is this reaction balanced?
Does it obey the Law of Conservation of
Matter?
2 Cu (s) + O2 (g)  2 CuO (s)
COEFFICIENTS - indicates
the number of units of each
substance involved.
1. Does the oxygen molecule have a coefficient?
2. What do the subscripts represent?
3. Can subscripts be removed from chemical equations?
The above reaction reads: 2 copper atoms react with 1 oxygen
to produce 2 molecules of copper oxide.
C.2: Accounting for Atoms
To obey the Law of Conservation of Matter
atoms can not be created or destroyed.
All atoms must be accounted for on both
sides of the reaction arrow.
Accounting for Atoms
Homework:
 Complete C1 Supplement worksheet


Due:
C.2: Accounting for Atoms
Classwork:
 Page 155-157
 Questions 1-5, parts A, C and D only

C.3: Balancing Chemical
Equations
Balancing Equations
Balancing Equations
Law of Conservation of
Atoms:
 The
number of atoms of
each type of element
must be the same on
each side of the
equation.
Balancing Equations
 Balancing
 Balance
hints:
the metals first.
 Balance the non metals next.
 Save the oxygen and hydrogen
atoms until the end.
How do we Balance
Equations?
Below is the chemical equation for the
reaction of hydrogen gas with oxygen gas
to produce water.
H2 + O2  H2O
Is this reaction balanced?
Balancing Equations
Hydrogen + oxygen
H 2 + O2
water
H2O
Hydrogen and oxygen are diatomic elements.
 Their subscripts cannot be changed.
 The subscripts on water cannot be changed.

Balancing Equation
H2 + O2
 Count
H2O
the atoms on each side.
 Reactant
side: 2 atoms H and 2 atoms O
 Product side: 2 atoms H and 1 atom O
Balancing Equations
H2 + O 2
H2O
 If
the subscripts cannot be
altered, how can the atoms be
made equal?
 Adjust
the number of molecules
by changing the coefficients.
Balancing Equations
H2 + O2




2H2O
Balance one atom at a time: Let’s start with
balancing oxygen!
Reactants: 2 atoms of H and 2 atoms of O
Products: 4 atoms of H and 2 atoms of O
H is no longer balanced!
Balancing Equations
2H2 + O2




2H2O
Now try to balance the hydrogens by placing a 2
in front of the reactant H2
Reactant side: 4 atoms of H and 2 atoms of O
Product side: 4 atoms of H and 2 atoms of O
It’s Balanced!
How do we Balance
Equations?
Number of
compounds in
the reaction
Coefficients
2 H2 + O2  2 H2O
Subscripts
# of atoms in a
compound
Subscripts can not be changed.
Coefficients balance atoms in an equation
How to Balance By
Inspection:
1
Make a table of elements
_____ CH4 + _____ O2  _____H2 O + _____ C O2
Reactants
C
H
O
Products
How to Balance By
Inspection:
2
Count the number of each element or ion on the
reactants and products side.
Don’t forget to add all the atoms of the same element together—
even if it appears in more than one compound!
_____ CH4 + _____ O2  _____H2 O + _____ C O2
Reactants
Products
C
1
1
H
4
2
O
2
3
How to Balance By
Inspection:
3
Add coefficients to balance the numbers
Each time you add a coefficient, update your table with the new
quantities of each atom.
_____ CH4 + _____
2 O2 _____
2 H2 O + _____ C O2
Reactants
Products
C
1
1
H
4
2
4
O
2
3
4
4
How to Balance By
Inspection:
4
Place a “1” in any empty coefficient location
Filling each coefficient location lets you and the grader know that you
finished the problem rather than you left some blank because you
weren’t done!
_____
1 CH4 + _____
2 O2 _____
2 H2 O + _____
1 C O2
Reactants
Products
C
1
1
H
4
2
4
O
2
3
4
4
What do Coefficients Really Mean?
CH4 + 2 O2  CO2 + 2 H2O
H
H
C
H
O
O
O
H
O
O
C
O
H
O
O
H
Total:
1C
4H
4O
The equation is balanced.
H
Total:
1C
4H
4O
H
Balancing Equations
N 2 + H2
Nitrogen + hydrogen
 Count
NH3
ammonia
atoms.
 Reactants:
2 atoms N and 2 atoms H
 Products: 1 atom N and 3 atoms of H
Balancing Equations
 Nothing
is balanced.
 Balance the nitrogen first by placing a
coefficient of 2 in front of the NH3.
N 2 + H2
2NH3
Balancing Equations
Hydrogen is not balanced.
 Place a 3 in front of H2.
 Reactant side: 2 atoms N, 6 atoms H
 Product side: 2 atoms N, 6 atoms H
 It’s balanced!

N2 + 3H2
2NH3
Balancing Equations
Cu + H2SO4
CuSO4 + H2O + SO2
Count Atoms:
 Reactants: Cu – 1, H – 2, S – 1, O – 4
 Products: Cu – 1, H – 2, S - 2, O - 7

Balancing Equations
 Sulfur
is not balanced.
 Place a two in front of sulfuric acid.
 Count Atoms:
Reactants: Cu – 1, H – 4, S – 1, O – 8
 Products: Cu – 1, H – 2, S - 2, O - 7

Cu + 2H2SO4
CuSO4 + H2O + SO2
Balancing Equations
 Hydrogen
needs to be balanced so
place a 2 in front of the H2O.
 Count the number of atoms.
Cu + 2H2SO4
CuSO4 + 2H2O + SO2
Balancing Equations
Cu + 2H2SO4
CuSO4 + 2H2O + SO2
Reactants: Cu – 1, H – 4, S – 2, O – 8
 Products: Cu – 1, H – 4, S – 2, O – 8


It’s balanced!
Balancing Equations
Ca3(PO4)2 + H2SO4
 Count
CaSO4 + H3PO4
atoms.
 Reactants: Ca – 3 atoms, P – 2
atoms, O – 12 atoms, H – 2 atoms, S
– 1 atom
Balancing Equations
 Side
 The
note on Ca3(PO4)2
subscript after the
phosphate indicates two
phosphate groups.
 This means two P and eight O
atoms.
Balancing Equations
Ca3(PO4)2 + H2SO4
 Count
CaSO4 + H3PO4
atoms in the product.
 Ca atoms – 1, S atom – 1, O atoms – 8;
H atoms – 3, P atom – 1,
Balancing Equations
 Balance
the metal first by placing a
coefficient of 3 in front of CaSO4.
Ca3(PO4)2 + H2SO4
3CaSO4 + H3PO4
Balancing Equations
 Now
balance the S atoms followed by
the P atoms.
Ca3(PO4)2 + 3H2SO4
3CaSO4 + H3PO4
Balancing Equations
A
coefficient of 2 placed in front of
H3PO4 which balances both hydrogen
and phosphate.
Ca3(PO4)2 + 3H2SO4
3CaSO4 + 2H3PO4
Balancing Equations
 This
method of
balancing
equations is trial
and error.
 Practice.
Let’s Practice #1
Example:
Balance the
following
equation
__ HCl + __ Ca(OH)2  __ CaCl2 + __ H2O
Let’s Practice #2
Example:
Balance the
following
equation
__ H2 + __ O2  __ H2O
Let’s Practice #3
Example:
Balance the
following
equation
__ Fe + __ O2  ___ Fe2O3
C.4: Balancing Equations
Classwork:
 Complete questions 1-6 on page 160161


Due:
Balancing Equations
Homework:
 Complete C3 Supplement worksheet


Due:
Balancing Equations

Additional Worksheets
C.5 The Mole Concept
Definition:
Mole – Unit for counting for chemists.
Counting atoms is impractical because atoms are
so small and are not visible to the naked eye.
You can not weigh individual atoms on
laboratory balance.
Relating Mass to Numbers of
Atoms
Introduction of three very important concepts:
1) The mole
2) Avogadro’s number
3) Molar mass
What is a counting unit?
You’re already familiar with one counting unit…a
“dozen”
A dozen = 12
“Dozen”
12
A dozen doughnuts
12 doughnuts
A dozen books
12 books
A dozen cars
12 cars
A dozen people
12 people
A Mole of Particles
Contains 6.02 x 1023 particles
Avogadro’s Number
1 mole C = 6.02 x 1023 C atoms
1 mole H2O = 6.02 x 1023 H2O molecules
1 mole NaCl= 6.02 x 1023 NaCl “molecules”
How big is a mole?




Enough soft drink cans to cover the
surface of the earth to a depth of over
200 miles.
If you had Avogadro's number of
unpopped popcorn kernels, and
spread them across the United States
of America, the country would be
covered in popcorn to a depth of over
9 miles.
If we were able to count atoms at the
rate of 10 million per second, it would
take about 2 billion years to count the
atoms in one mole.
One mole of paper clips would wrap
around the earth 400 trillion times!
What does a “mole” count
in?
A mole = 6.02  1023 (called Avogadro’s number)
6.02  1023 = 602,000,000,000,000,000,000,000
“mole”
6.02  1023
1 mole of doughnuts
6.02  1023 doughnuts
1 mole of atoms
6.02  1023 atoms
1 mole of molecules
6.02  1023 molecules
1. Mole of a substance = grams of
substance/MW of substance
2. The mole enables chemists to move
from the microscopic world of
atoms and molecules to the real
world of grams and kilograms.
Molar Mass
Definition
Molar Mass – The mass for one mole
of an atom or molecule.
Other terms commonly used for the same meaning:
Molecular Weight
Atomic Weight (used for atoms)
Mass for 1 mole of atoms
The average atomic mass = grams for 1 mole
Average atomic mass is found on the periodic table
Element
Mass
1 mole of carbon atoms
12.01 g
1 mole of oxygen atoms
16.00 g
1 mole of hydrogen
atoms
1.01 g
Unit for molar mass: g/mole or g/mol
One mole of carbon (12 grams) and
one mole of copper (63.5 grams)
Molar Mass
Examples:
Molar mass of lithium (Li) = 6.94 g/mol
Molar mass of helium (He) = 4.00 g/mol
Molar mass of mercury (Hg) = 200.6 g/mol
Molar Mass
A molar mass of an element contains one
mole of atoms.
4.00 g helium = 1mole = 6.02 x 1023 atoms.
6.94 g lithium = 1mole = 6.02 x 1023 atoms.
200.6 g mercury = 1mole = 6.02 x 1023
atoms.
Molar mass for molecules
The molar mass for a molecule = the
sum of the molar masses of all the
atoms
Calculating a Molecule’s
Mass
To find the molar mass of a molecule:
1
Count the number of each type of atom
2
Find the molar mass of each atom on the periodic
table
3
Multiple the # of atoms  molar mass for each atom
4
Find the sum of all the masses
Example: Molar Mass
Example:
Find the
molar
mass for
CaBr2
Example: Molar Mass
1
Example:
Find the
molar
mass for
CaBr2
Count the number of each type of atom
Ca
1
Br
2
Example: Molar Mass
2
Example:
Find the
molar
mass for
CaBr2
Find the molar mass of each atom on the periodic table
Ca
1
40.08 g/mole
Br
2
79.91 g/mole
Example: Molar Mass
3
Example:
Find the
molar
mass for
CaBr2
Multiple the # of atoms  molar mass for each atom
Ca
1  40.08 g/mole =
40.08 g/mole
Br
2  79.91 g/mole =
159.82 g/mole
Example: Molar Mass
4
Example:
Find the
molar
mass for
CaBr2
Find the sum of all the masses
Ca
1  40.08 g/mole =
40.08 g/mole
Br
2  79.91 g/mole =
+ 159.82 g/mole
199.90 g/mole
1 mole of CaBr2 molecules would have a mass of
199.90 g
Molar Mass
A molar mass of a compound is the sum
of the molar masses of the elements.
Example: Water, H2O:
2 H = 2 x 1 g/mole = 2 g/mole
1 O = 1 x 16 g/mole = 16 g/mole
molar mass of H20 =18 g/mol
Molar Mass
A molar mass of a compound is the sum
of the molar masses of the elements.
Example: methane, CH4:
4 H = 4 x 1 g = 4 g/mole
1 C = 1 x 12g = 12 g/mole
molar mass of CH4 =16 g/mol
Example: Molar Mass &
Parenthesis
Be sure to distribute the subscript outside the
parenthesis to each element inside the parenthesis.
Example:
Find the
molar
mass for
Sr(NO3)2
Let’s Practice #1
Example:
Find the
molar
mass for
CH2Cl2
Let’s Practice #2
Example:
Find the
molar
mass for
Al(OH)3
Classwork: C.6: Molar Masses
Page 163; questions1-9
Homework: C.5 Supplement
Worksheet
Due:
C.8: Molar
Relationships
Molar mass for molecules
The molar mass for a molecule = the
sum of the molar masses of all the
atoms in the molecule.
Example:
1 mole of H2O = 18 grams
1 mole of H2O weighs 18 grams.
What happens when you do not have 18 grams
of water?
How do you calculate the number of moles of
water?
Conversion Factors
Conversion factor – a way that can be
used to convert from one unit to the other.
Example: the conversion between
quarters and dollars:
4 quarters
1 dollar
or
1 dollar
4 quarters
Conversion Factors
Example:
Determine the number of quarters in 12
dollars?
(conversion factor)
? Quarters = 12 dollars x
4 quarters
= 48 quarters
1 dollar
Conversion Factors
Lets return to the term dozen.
12 roses = 1 dozen roses (conversion factor)
6 roses = ½ dozen
Calculations:
6 roses____
12 roses/dozen
= 0.5 dozen
Molar masses can be used as
a conversion factor in chemical
calculations.
Molar masses allow the
chemist to convert from grams
to moles of a compound and
from moles to grams of a
compound.
molar mass
Grams
Moles
Gram/Mole Conversions
When converting between grams and
moles, the molar mass is needed
Example: The molar mass of helium is 4.00
g/mol. To calculate how many grams of
helium are in 2 moles of helium:
amount of He in moles
amount of He in grams
4.00 g He
2.00 mol He x
1 mol He
= 8.00 g He
Example: Moles to Grams
Example:
How many
grams are
in 1.25
moles of
water?
Example: Moles to Grams
When converting between grams and moles, the
molar mass is needed
Example:
How many
grams are
in 1.25
moles of
water?
1.25 mol H2O
H 2  1.01 g/mole =
2.02 g/mole
O 1  16.00 g/mole = + 16.00 g/mole
18.02 g/mole
1 mole H2O molecules = 18.02 g
18.02 g H2O
1 mol H2O
22.53 g H2O
= _______
Problem
How many grams are in
2.25 moles of iron, Fe?
Problem
How many grams are in
2.25 moles of iron, Fe?
Answer: 126 grams Fe
Problem
How many grams are
in 0.375 moles of
potassium, K?
Problem
How many grams are
in 0.375 moles of
potassium, K?
Answer: 14.7 grams
Problem
How many grams are
in 20 moles of
methane, CH4?
Problem
How many grams are
in 20 moles of
methane, CH4?
Answer: 320 grams
Grams to Moles
If you are given a quantity in grams
you can determine how moles you
have by using the molar mass.
molar mass
Grams
Moles
Let’s Practice
Example:
How many
moles are
in 25.5 g
NaCl?
Let’s Practice
Na 1  22.99 g/mole = 22.99 g/mole
Cl 1  35.45 g/mole = + 35.45 g/mole
58.44 g/mole
Example:
How many
moles are
in 25.5 g
NaCl?
25.5 g NaCl
1 mole NaCl molecules = 58.44 g
1
mole NaCl
58.44 g NaCl
0.44
= _______
mole NaCl
Example: Grams to Moles
A chemist produced 11.9 g of aluminum, Al.
How many moles of aluminum were
produced?
mass of Al in grams
amount of Al in moles
1 mol Al
11.9 g Al x
27 g Al
= 0.44 mol Al
Problem
How many moles of
calcium, Ca, are in 5.00
grams of calcium?
Problem
How many moles of
calcium, Ca, are in 5.00
grams of calcium?
Answer: 0.125 moles
Problem
How many moles of
sodium chloride, NaCl, are
in 75 grams of NaCl?
Problem
How many moles of
sodium chloride, NaCl, are
in 75 grams of NaCl?
Answer: 1.28 moles
C.8: Molar Relationships
Classwork: C.8 Problems
Page 166 questions 1-3
(hold on 4)
Homework: C.7-1 Supplement
Worksheet
Due:
C.7:Equations and Molar
Relationships
How are chemical equations and molar
masses related?
2 CuO
2 moles
+
C
2 Cu +
CO2
1 mole
2 moles
1mole
C.7:Equations and Molar
Relationships
2 CuO
2 moles
+
C
1 mole
2 Cu + CO2
2 moles
1mole
159.1 g + 12.01 g
127.1 g + 44 g
171.1 g total
171.1 g total
Same total mass on the reactant and product side
C.7:Equations and Molar
Relationships
Stoichiometry
PowerPoint
Example: Grams to
Molecules
Example:
How many
molecules
are in
25.5 g
NaCl?
Skip
Example: Grams to
Molecules
Example:
How many
molecules
are in
25.5 g
NaCl?
Na 1  22.99 g/mole = 22.99 g/mole
Cl 1  35.45 g/mole = + 35.45 g/mole
58.44 g/mole
1 moles NaCl molecules = 58.44 g
1 mol = 6.021023 molecules
25.5 g NaCl
1
mol NaCl
58.44 g NaCl
6.021023 molecules NaCl
1
mol NaCl
2.63  1023 molecules NaCl
= _________
Let’s Practice #4
Example:
How many
grams is a
sample of
2.75 × 1024
molecules of
SrCl2?
Skip
Homework: C.7-2 Supplement
Worksheet
Due:
C.9: Percent Composition
Percent Composition - The percent mass
of each element found in a mixture
Mixtures
Sample problem 1: A post-1982 penny
with a mass of 2.50 g is composed of 2.44
g of zinc and 0.06 g copper. What is the
percent composition of each element?
The percent composition of the penny
can be found by dividing the mass of
each element by the total mass of the
penny and multiplying by 100%.
Zinc:
2.44 g zinc
2.50 g total
Copper:
x 100% = 97.5%
0.06 g Cu x 100% = 2.5%
2.50 g total
Total percent must equal 100%
Relating molar mass and
percent composition
Sample problem 2: The formula for the
copper containing mineral chalcocite is
Cu2S.
1) What percentage of copper (Cu) is in
this mineral?
2) What percentage of sulfur (S) is in this
mineral?
The formula for chalcocite indicates one
mole of Cu2S contains two molecules of
Cu (127.1 g) and one mole of S (32 g).
The molar mass of Cu2S = 159.2 g.
% Cu = mass of Cu x 100%
mass of Cu2S
% Cu:
127.1 g Cu x 100% = 79.9%
159.2 g Cu2S
The formula for chalcocite indicates one
mole of Cu2S contains two molecules of
Cu (127.1 g) and one mole of S (32 g).
The molar mass of Cu2S = 159.2 g.
% S = mass of S
x 100%
mass of Cu2S
% Cu:
32 g Cu
x 100% = 20.1%
159.2 g Cu2S
C.10: Percent Composition
Classwork: page 168,
problems 2 and 4
C.10: Percent Composition
Homework: C.10: Worksheet
percent composition
Due:
C.11:
Lab: Retrieving Copper
Read over the lab procedure
C.12: Conservation
Renewable resources: resources that can
be replenished.
Examples include fresh water, air, fertile
soil, plants and animals.
As long as natural cycles are not disturbed
too much, supplies of renewable resources
can be maintained indefinitely.
Nonrenewable resources: resources that
can not be readily replenished.
Examples include metals, natural gas, coal
and petroleum.
The length of time to replenish these
resources are very large.
End of Unit
Replacing: replace a resource by finding a
substitute, preferably from renewable
resources.
Reusing: refurbish or repair an item to use
again. Examples include car parts and printer
cartridges.
Recycle: reprocess to use again. Examples
include aluminum cans, newspaper and glass
bottles.
C.12: Conservation
C12: Complete worksheet.
These terms will appear again on a test or
final exam!