Download 06 Chapter 1.3 Powerpoint Chem Quant 1.3 2013

IB Topic 1: Quantitative Chemistry
1.3: Chemical Equations
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1.3.1 Deduce chemical equations when all
reactants and products are given
1.3.2 Identify the mole ratio of any two species
in a chemical reaction.
1.3.3 Apply the state symbols (s), (l), (g) and
(aq)
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1.3.1 Deduce chemical equations when all
reactants and products are given.
What is a chemical reaction?
Rearrangement of atoms forming new substances
Reactants  Products
Some reactions are desirable…
Glucose + oxygen  Carbon dioxide + water
…some are not.
Iron + oxygen  iron (III) oxide
(a.k.a. rust)
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Chemical Reactions
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A few ways to determine
whether a chemical reaction
has taken place:
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Heat is absorbed or given off
Change in color
Change in odor
Production of a gas or solid
(precipitate)
Not easily reversible (it won’t
recreate the reactants)
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1.3.1 Deduce chemical equations when all
reactants and products are given.
Chemical formulas are easier and more informative to use
in equations than words.
A skeleton equation is a chemical equation that
shows what reactants and products are involved. It
does not necessarily indicate the relative amounts of
reactants and products.
Fe + O2  Fe2O3
Sometimes the skeleton equation is balanced (all
coefficients = 1).
SnO2(s)  Sn(s) + O2(g)
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Identify types of chemical reactions
Types of Chemical Reactions
1)
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Synthesis or Combination Reaction: Two
or more substances combine to form a single
substance.
2Na(s) + Cl2(g)  2NaCl(s)
CaO(s) + H2O(l)  Ca(OH)2(aq)
2H2(g) + O2(g)  2H2O (l)
4Fe(s) + 3O2(g)  2Fe2O3(s)
N2(g) + 3H2(g)  2NH3(g)
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Identify types of chemical reactions
Types of Chemical Reactions
2)
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Decomposition Reactions: A single substance is
broken down into two or more substances.
CaCO3(s)  CaO(s) + CO2(g)
2H2O(l)  2H2(g) + O2(g)
2H2O2(aq)  2H2O(l) + O2(g)
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Identify types of chemical reactions
Types of Chemical Reactions
3)
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Single-Replacement Reactions: One element
replaces another in a compound.
Mg(s) + Zn(NO3)2(aq)  Mg(NO3)2(aq) + Zn(s)
2K(s) + 2H2O(l)  2KOH (aq) + H2(g)
Cu(s) + 2AgNO3(aq)  Cu(NO3)2(aq) + 2Ag(s)
2Al(s) + Fe2O3(s)  Al2O3(s) + 2Fe(s)
Cl2(g) + 2NaBr(aq)  2NaCl(aq) + Br2(g)
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Identify types of chemical reactions
Types of Chemical Reactions
4)
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Double-Replacement Reactions: Ions of two
reacting compounds trade places.
Na2CO3(aq) + CaCl2(aq)  CaCO3(s) + 2NaCl(aq)
Na2S(aq) + Cd(NO3)2(aq)  CdS(s) + 2NaNO3(aq)
NaOH(aq) + HCl(aq)  NaCl(aq) + H2O(l)
H2SO4(aq) + 2NaCN(aq)  Na2SO4(aq) + 2HCN(g)
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Identify types of chemical reactions
Types of Chemical Reactions
5)
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Combustion Reactions: A compound reacts with
oxygen. Products are usually carbon dioxide and
water.
CH4(g) + O2(g)  CO2(g) + H2O(g)
C3H8(g) + 5O2(g)  3CO2(g) + 4H2O(g)
2C8H18(l) + 25O2(g)  16CO2(g) + 18H2O(g)
C2H5OH(l) + 3O2(g)  2CO2(g) + 3H2O(g)
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1.3.1 Deduce chemical equations when all
reactants and products are given.
Balancing chemical equations
The Law of Conservation of Matter: In a chemical (non-nuclear)
reaction, atoms are neither created nor destroyed.
For an equation to be balanced the number of atoms of each element is the same
on both sides of the equation.
H2(g) + O2(g)  H2O (l)
(unbalanced)
2H2(g) + O2(g)  2H2O (l)
(balanced)
K(s) + H20(l)  KOH (aq) + H2(g) (unbalanced)
2K(s) + 2H20(l)  2KOH (aq) + H2(g)
(balanced)
C6H12O6 + O2  CO2 + H2O
C6H12O6 + 6O2  6CO2 + 6H2O
(unbalanced)
(balanced)
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Counting Atoms
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Subscripts indicate how many of a specific
atom is present in a compound
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Coefficients tell us how many units of the
compound we have
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This is the small number to the bottom right of an
element in a compound.
Ex. H2O… 2 is the subscript
Ex. 3 H2O means that we have 3 water molecules
So, how many hydrogen atoms do we have?
Oxygen atoms?
Let’s practice
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1.3.1 Deduce chemical equations when all
reactants and products are given.
Rules for Balancing Equations
1) Write the correct formulas for the reactants on the left side of the yield
sign and products on the right side.
2) Count the number of atoms of each element in the products and the
reactants.
3) Balance the elements one at a time by using coefficients. Do not change
the subscripts in the chemical formulas.
4) Check each atom or polyatomic ion to make sure the equation is
balanced.
5) Make sure all coefficients are in the lowest possible ratio.
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1.3.1 Deduce chemical equations when all
reactants and products are given.
N2 + H2  NH3
Rules for Balancing
Equations
1) Formulas given
1) Write the correct formulas for
2) Reactants
Products
the reactants on the left side of 2N 2H  1N 3H
the yield sign and products on
the right side.
3) Balance N by putting 2 in front of NH3
2) Count the number of atoms of
each element in the products
N2 + H2  2NH3
and the reactants.
Balance H by putting 3 in front of H2
3) Balance the elements one at a
N2 + 3H2  2NH3
time by using coefficients. Do
not change the subscripts in
4)
2N 6H  2N 6H
the chemical formulas.
4) Check each atom or polyatomic
ion to make sure the equation
5)
Lowest ratio of coefficients
is balanced.
5) Make sure all coefficients are in
the lowest possible ratio.
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1.3.1 Deduce chemical equations when all
reactants and products are given.
KClO3  KCl + O2
Rules for Balancing Equations
1) Write the correct formulas for the
reactants on the left side of the
yield sign and products on the right
side.
2) Count the number of atoms of each
element in the products and the
reactants.
3) Balance the elements one at a time
by using coefficients. Do not
change the subscripts in the
chemical formulas.
4) Check each atom or polyatomic ion
to make sure the equation is
balanced.
5) Make sure all coefficients are in the
lowest possible ratio.
1) Formulas given
2) Reactants
Products
1K 1Cl 3O  1K 1Cl 2O
3) Balance O by putting 2 in front of
KClO3
and 3 in front of O2
2KClO3  KCl + 3O2
Balance K & Cl by putting 2 in front of
KCl
2KClO3  2KCl + 3O2
4)
5)
2K 2Cl 6O  2K 2Cl 6O
Lowest ratio of coefficients
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1.3.1 Deduce chemical equations when all
reactants and products are given.
Balance the following equations
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Fe(s) + O2(g)  Fe2O3(s)
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CaCO3(s)  CaO(s) + CO2(g)
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Al(s) + Fe2O3(s)  Al2O3(s) + Fe(s)
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H2SO4(aq) + NaCN(aq)  Na2SO4(aq) + HCN(g)
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C2H5OH(l) + O2(g)  CO2(g) + H2O(g)
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1.3.1 Deduce chemical equations when all
reactants and products are given.
Balance the following equations
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4Fe(s) + 3O2(g)  2Fe2O3(s)
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CaCO3(s)  CaO(s) + CO2(g)
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2Al(s) + Fe2O3(s)  Al2O3(s) + 2Fe(s)
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H2SO4(aq) + 2NaCN(aq)  Na2SO4(aq) + 2HCN(g)
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C2H5OH(l) + 3O2(g)  2CO2(g) + 3H2O(g)
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1.3.1 Deduce chemical equations when all
reactants and products are given.
Aluminum bromide + chlorine yield
aluminum chloride + bromine
2AlBr3 + 3Cl2  2AlCl3 + 3Br2
Seven elements are diatomic: H2, N2, O2, F2, Cl2, Br2, I2
Copper + oxygen produces copper(I) oxide
4Cu + O2  2Cu2O
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1.3.1 Deduce chemical equations when all
reactants and products are given.
Sodium chlorate decomposes to sodium chloride and
oxygen gas
Aluminum nitrate plus sodium hydroxide yields
aluminum hydroxide plus sodium nitrate
Ethane (C2H6) burns in oxygen to produce
carbon dioxide and water vapor
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1.3.1 Deduce chemical equations when all
reactants and products are given.
Sodium chlorate decomposes to sodium chloride and
oxygen gas
2NaClO3  2NaCl + 3O2
Aluminum nitrate plus sodium hydroxide yields
aluminum hydroxide plus sodium nitrate
Al(NO3)3 + 3NaOH  Al(OH)3 + 3NaNO3
Ethane (C2H6) burns in oxygen to produce
carbon dioxide and water vapor
2C2H6 + 7O2  4CO2 + 6H2O
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1.3.3 Apply the state symbols (s), (l), (g) and
(aq)
Symbols used in chemical reactions
+
Used to separate two reactants or products
A “Yields” arrow () separates products from reactants
 Used in place of  for reversible reactions
(s) Designates a solid reactant or product
(l) Designates a liquid reactant or product
(g) Designates a gaseous reactant or product
(aq) Designates an aqueous reactant or product
 Indicates that heat is supplied to the reaction
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1.3.3 Apply the state symbols (s), (l), (g) and
(aq)
(s) Designates a solid reactant or product
(l) Designates a liquid reactant or product
(g) Designates a gaseous reactant or product
(aq)
Designates an aqueous reactant or
product
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1.3.3 Apply the state symbols (s), (l), (g) and
(aq)
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Rule 1. All compounds of Group IA
elements (the alkali metals) are soluble.
For example, NaNO3, KCl, and LiOH are all
soluble compounds. This means that an
aqueous solution of KCl really contains the
predominant species K+ and Cl- and,
because KCl is soluble, no KCl is present as
a solid compound in aqueous solution:
KCl(s) => K+(aq.) + Cl-(aq.)
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1.3.3 Apply the state symbols (s), (l), (g) and
(aq)
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Rule 2. All ammonium salts (salts of NH4+)
are soluble.
For example, NH4OH is a soluble
compound. Molecules of NH4OH completely
dissociate to give ions of NH4+ and OH- in
aqueous solution.
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1.3.3 Apply the state symbols (s), (l), (g) and
(aq)
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Rule 3. All nitrate (NO3-), chlorate (ClO3-),
perchlorate (ClO4-), and acetate (CH3COOor C2H3O2- are soluble.
For example, KNO3 would be classified as
completely soluble by rules 1 and 3. Thus,
KNO3 could be expected to dissociate
completely in aqueous solution into K+ and
NO3- ions: KNO3 => K+(aq.) + NO3-(aq.)
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1.3.3 Apply the state symbols (s), (l), (g) and
(aq)
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Rule 4. All chloride (Cl-), bromide (Br-),
and iodide (I-) salts are soluble except for
those of Ag+, Pb2+, and Hg22+.
For example, AgCl is a classic insoluble
chloride salt:
AgCl(s) <=> Ag+(aq.) + Cl-(aq.)
(Ksp = 1.8 x 10-10).
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1.3.3 Apply the state symbols (s), (l), (g) and
(aq)
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Rule 5. All sulfate ( SO4=) compounds are soluble
except those of Ba2+, Sr2+, Ca2+, Pb2+, Hg22+, and
Hg2+. Ca2+ and Ag+ sulfates are only moderately
soluble.
For example, BaSO4 is insoluble (only soluble to a
very small extent):
BaSO4(s) <=> Ba2+(aq.) + SO42-(aq.)
(Ksp = 1.1 x 10-10).
Na2SO4 is completely soluble:
Na2SO4(s) => 2 Na+(aq.) + SO42-(aq.).
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1.3.3 Apply the state symbols (s), (l), (g) and
(aq)
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Rule 6. All hydroxide (OH-) compounds are
insoluble except those of Group I-A (alkali metals)
and Ba2+, Ca2+, and Sr2+.
For example, Mg(OH)2 is insoluble
(Ksp = 7.1 x 10-12)
NaOH and Ba(OH)2 are soluble, completely
dissociating in aqueous solution:
NaOH(s) => Na+(aq.) + OH- (aq.), a strong base
Ba(OH)2(s) => Ba2+ (aq.) + 2OH- (aq.)
(Ksp = 3 x 10-4)
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1.3.3 Apply the state symbols (s), (l), (g) and
(aq)
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Rule 7. All sulfide (S2-) compounds are
insoluble except those of Groups I-A and
II-A (alkali metals and alkali earths).
For example,
Na2S(s) <=> 2Na+ (aq.) + S2- (aq.)
MnS is insoluble (Ksp = 3 x 10-11).
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1.3.3 Apply the state symbols (s), (l), (g) and
(aq)
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Rule 8. All sulfites (SO32-), carbonates
(CO32-), chromates (CrO42-), and
phosphates (PO43-) are insoluble except
for those of NH4+ and Group I-A (alkali
metals)(see rules 1 and 2).
For example, calcite,
CaCO3 (s) <=> Ca2+ (aq.) + CO32- (aq.)
(Ksp = 4.5 x 10-9).
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1.3.3 Apply the state symbols (s), (l), (g) and (aq)
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1.
2.
3.
4.
OR... Use a Solubility Chart 
Let’s use the following reactions to apply this
assessment statement
Hydrogen peroxide decomposes into water and
oxygen gas
Potassium phosphate reacts with magnesium
chloride to form magnesium phosphate and
potassium chloride
Iron reacts with copper (II) sulfate to form iron (II)
sulfate and copper
Sulfuric acid (hydrogen sulfate) reacts with copper
(II) oxide to produce copper (II) sulfate and water
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1.3.3 Apply the state symbols (s), (l), (g) and (aq)
1. Hydrogen peroxide decomposes into water and
oxygen gas
2 H2O2(aq)  2 H2O(l) + O2(g)
2. Potassium phosphate reacts with magnesium chloride
to form magnesium phosphate and potassium
chloride
2K3PO4 (aq) + 3MgCl2 (aq)  Mg3(PO4)2 (s) + 6KCl (aq)
3. Iron reacts with copper (II) sulfate to form iron (II)
sulfate and copper
Fe(s) + CuSO4(aq)  FeSO4(aq) + Cu(s)
4. Sulfuric acid (hydrogen sulfate) reacts with copper (II)
oxide to produce copper (II) sulfate and water
H2SO4 (aq) + CuO (s)  CuSO4 (aq) + H2O (l)
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Cookies and Chemistry…Huh!?!?
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Just like chocolate chip
cookies have recipes,
chemists have recipes as well
Instead of calling them
recipes, we call them reaction
equations
Furthermore, instead of using
cups and teaspoons, we use
moles
Lastly, instead of eggs,
butter, sugar, etc. we use
chemical compounds as
ingredients
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1.3.2 Identify the mole ratio of any two
species in a chemical reaction.
Coefficients are in particles or in moles
2 f.u. NaClO3  2 f.u. NaCl + 3 molecules O2
2 mol NaClO3  2 mol NaCl + 3 mol O2
We will use moles because we measure moles in grams
The coefficients give us mole ratios
Mole ratio NaClO3:NaCl is 2:2
Mole ratio NaCl:O2 is 2:3
Mole ratio O2:NaCl is 3:2
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1.3.2 Identify the mole ratio of any two
species in a chemical reaction.
Consider the following equation
2K2Cr2O7 + 2H2O + 3S  4KOH + 2Cr2O3 + 3SO2
What
What
What
What
is the KOH:S mole ratio?
is the K2Cr2O7:Cr2O3 mole ratio?
is the mole ratio between sulfur dioxide and water?
species will give mole ratios of 4:2?
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1.3.2 Identify the mole ratio of any two
species in a chemical reaction.
Consider the following equation
2K2Cr2O7 + 2H2O + 3S  4KOH + 2Cr2O3 + 3SO2
What is the KOH:S mole ratio? 4:3
What is the K2Cr2O7:Cr2O3
2:2
What is the mole ratio between sulfur dioxide and water? 3:2
What species will give mole ratios of 4:2?
KOH:K2Cr2O7
KOH:H2O
KOH:Cr2O3
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1.3.2 Identify the mole ratio of any two
species in a chemical reaction.
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Looking at a reaction tells us how much of
something you need to react with something
else to get a product (like the cookie recipe)
Be sure you have a balanced reaction before
you start!
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Example: 2 Na + Cl2  2 NaCl
This reaction tells us that by mixing 2 moles of
sodium with 1 mole of chlorine we will get 2 moles
of sodium chloride
What if we wanted 4 moles of NaCl? 10 moles?
50 moles?
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1.3.2 Identify the mole ratio of any two
species in a chemical reaction.
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These mole ratios can be used to calculate
the moles of one chemical from the given
amount of a different chemical
Example: How many moles of chlorine is
needed to react with 5 moles of sodium
(without any sodium left over)?
2 Na + Cl2  2 NaCl
5 moles Na 1 mol Cl2
2 mol Na
= 2.5 moles Cl2
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1.3.2 Identify the mole ratio of any two
species in a chemical reaction.


Looking at a reaction tells us how much of
something you need to react with something
else to get a product (like the cookie recipe)
Be sure you have a balanced reaction before
you start!

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
Example: 2 Na + Cl2  2 NaCl
This reaction tells us that by mixing 2 moles of
sodium with 1 mole of chlorine we will get 2 moles
of sodium chloride
What if we wanted 4 moles of NaCl? 10 moles?
50 moles?
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1.3.2 Identify the mole ratio of any two
species in a chemical reaction.

How many moles of sodium chloride will be
produced if you react 2.6 moles of chlorine gas
with an excess (more than you need) of sodium
metal?
2 Na + Cl2  2 NaCl
2.6 moles Cl2 2 mol NaCl
1 mol Cl2
= 5.2 moles NaCl
1.3.2 Practice
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Write the balanced reaction for hydrogen gas
reacting with oxygen gas.
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2 H2 + O2  2 H2O
2 mol H2
How many moles of reactants are needed?
1 mol O2
What if we wanted 4 moles of water? 4 mol H2
mol O2 hydrogen
What if we had 3 moles of oxygen, how2 much
would we need to react and how much water would we
get?
6 mol H2, 6 mol H2O
What if we had 50 moles of hydrogen, how much oxygen
would we need and how much water produced?
25 mol O2, 50 mol H2O
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