Download Chemical Reactions

Chemical Reactions
2/3- 2/4
Signs of a chemical reaction?
Indicators of chemical reactions
Emission of energy
Formation of a gas
Formation of a precipitate
Color change
What is a chemical reaction composed
of?
1) Contains reactants and products
2) Formulas must be written correctly with
symbols and subscripts
3) Law of conservation of matter requires that
coefficients be used to ensure that atoms
Describing chemical reaction
The way atoms are joined is changed
Atoms aren’t created or destroyed.
May involve a catalyst
Can be described several ways
In a sentence
• Solid Copper reacts with chlorine gas to form
aqueous copper (II) chloride.
In a word equation or formula equation
• Copper(s) + chlorine(g)  copper(II) chloride(aq)
• Cu(s) + Cl2(g)  CuCl2(aq)
•
•
•
•
Reaction Energy
All chemical reactions are accompanied by a change
in energy.
Exothermic - reactions that release energy to
their surroundings (usually in the form of heat)
ΔH (enthalpy) is negative – energy leaving system
Endothermic - reactions that need to absorb heat
from their surroundings to proceed.
ΔH (enthalpy) is positive – energy coming into the
system
Reaction Energy
•Spontaneous Reactions - Reactions that proceed
immediately when two substances are mixed
together. Not all reactions proceed spontaneously.
Some require…
•Activation Energy – the amount of energy that
is required to start a chemical reaction.
•Once activation energy is reached the
reaction continues until you run out of
material to react.
What is a catalyst?
• Does not cause a reaction to occur, but
speeds up the rate which a reaction
occurs
• Can be in the form of the following:
– Energy- light, heat
– Chemicals
– Enzymes are biological or protein catalysts.
Word and Formula Equations
Word Equations
• Word Equations: an equation in which the
reactants and products in a chemical reaction are
represented by words instead of chemical
formulas.
• The problem with word equations is they do not
actually show the number of atoms or molecules
of each substance… formulas would have to be
written out for this to happen.
(Absent? We looked at examples of these in class)
Formula Equations
• Represents reactants and products of a
chemical reaction by their symbols or formulas.
• Unbalanced- does not account for law of
conservation of matter
• Balanced- using coefficients show the
representative numbers
• See table 2 on page 266 for important symbols
Summary of Symbols
Word Formula practice
Fe(s) + O2(g)  Fe2O3(s)
Solid iron reacts with oxygen gas to
form solid iron III oxide (rust).
Formula Equation
Nitric acid dissolved in water reacts
with solid sodium carbonate to form
liquid water and carbon dioxide gas
and sodium nitrate dissolved in water.
HNO3
(aq)
+ Na2CO3 (s)  NaNO3
(aq)
+ H2O(l)
Convert this to a formula equation
Solid iron (III) sulfide reacts with
gaseous hydrogen chloride to form
iron (II) chloride and hydrogen
sulfide gas.
Fe2S3 (s) + HCl(g)  FeCl2 (s) + H2S(g)
A silver spoon tarnishes. The solid silver
reacts with sulfur in the air to make
solid silver sulfide, the black material we
call tarnish and water.
Ag (s) + H2S (g) + O2 (g)  Ag2S (s) + H2O
In-class work
• In-class we practiced writing word and
formula equations for different chemical
reactions.
• See chemical reactions sheet
Balancing Equations
2 H2O(l)
2 H2(g) + ___ O2(g) ---> ___
___
•What Happened to the Other Oxygen Atom?
•This equation is not balanced! Until…
CH4 + 2 O2  CO2 + 2 H2O
Reactants
1 C atom
4 H atoms
4 O atoms
Timberlake, Chemistry 7th Edition, page 167
Products
1 C atom
4 H atoms
4 O atoms
Reactants  Products
+
C(s)
+
O2(g)
carbon
oxygen
CO2(g)
carbon dioxide
Reactants
1 carbon atom
2 oxygen atoms
Product
1 carbon atom
2 oxygen atoms
catalyst – speeds up reaction
Pt
+
2H2(g)
hydrogen
+
O2(g)
oxygen
Reactants
2 hydrogen atoms
4
2 oxygen atoms
Timberlake, Chemistry 7th Edition, page 164
Pt
2 H2O (l)
water
Unbalanced
Product
2 hydrogen atoms
4
1 oxygen atoms
2
Unbalanced and Balanced Equations
H
Cl
Cl
H
H
H
H
Cl
H2 + Cl2  HCl
reactants
H
Cl
2
2
(unbalanced)
H
H
Cl
Cl
Cl
H2 + Cl2  2 HCl
(balanced)
reactants products
products
1
1
Cl
H
Cl
2
2
2
2
Types of Reactions
• There are millions of reactions
• Objectives:
– We will learn 6 types.
– We will be able to predict the
products.
– We will be able to predict whether
they will happen at all.
Synthesis Reactions
• Also called combination reactions
• 2 elements, or compounds combine to
make one compound.
• A + B

AB
• Na (s) + Cl2 (g)  NaCl (s)
• Ca (s) +O2 (g)  CaO (s)
• SO3 (s) + H2O (l)  H2SO4 (s)
• We can predict the products if they
are two elements.
• Mg (s) + N2 (g)  Mg3N2 (s)
A simulation of the reaction:
2H2 + O2

2H2O
Synthesis Reaction
Direct combination reaction (Synthesis)
2 Na +

Cl2
Cl
Na
2 NaCl

Cl
Na
General form: A
+
element or
compound
B

element or
compound
AB
compound
Synthesis Reaction
Direct combination reaction (Synthesis)
2 Mg +
Mg
Mg2+

O2
2 MgO
OO2OO2-
Mg
Mg2+
General form: A
+
element or
compound
B

element or
compound
AB
compound
Decomposition Reactions
• decompose = fall apart
• one compound (reactant) falls apart into
two or more elements or compounds.
• Usually requires energy
• AB  A + B
electricity
 Na + Cl2
• NaCl   

• CaCO3   CaO + CO2
Decomposition Reactions
• Can predict the products if it is a binary
compound
• Made up of only two elements
• Falls apart into its elements
electricity
 H2 (g) + O2 (g)
• H2O   
• HgO  
Hg (s) + O2 (g)

Decomposition Reactions
• If the compound has more than two
elements you must be given one of the
products
• The other product will be from the
missing pieces

• NiCO3 (aq)   CO2 (g) + Ni (s)
• H2CO3(aq)

H2
(g)
+ CO2 (g)
Decomposition Reaction
Decomposition reaction
2 H2O
2 H2
+
O2
H
O
H
+
H
O
H
General form: AB
compound
A
+
B
two or more elements
or compounds
Single Replacement
•
•
•
•
single displacement
One element replaces another
Reactants= an element and a compound.
Products= a different element and a different
compound.
• A + BC

AC
+ B
• Based on the activity series (in other words the
element replacing must be more active.
Activity Series
Foiled again –
Aluminum loses to Copper
Element Reactivity
Li
Rb
K
Ba
Ca
Na
Mg
Al
Mn
Zn
Cr
Fe
Ni
Sn
Pb
H2
Cu
Hg
Ag
Pt
Au
Halogen Reactivity
F2
Cl2
Br2
I2
Potassium reacts with Water
POW!
Double Replacement
• Two things replace each other.
• Reactants must be two ionic compounds or
acids.
• Usually in aqueous solution
AB + CD
 AD + CB
ZnS
+ 2HCl

AgNO3 + NaCl 
ZnCl + H2S
AgCl + NaNO3
Double Replacement Reaction
Formation of a solid AgCl
AgNO3(aq) + KCl(aq)  KNO3 (aq) + AgCl(s)
Single and Double Replacement
Reactions
Single-replacement reaction
Mg
+
CuSO4
General form:
A
+ BC


MgSO4
AC
+
+
Cu
B
Double-replacement reaction
CaCO3
+
2 HCl
General form:
AB
+ CD

CaCl2
+
H2CO3

AD
+
CB
Combustion
• A reaction in which a compound
(often carbon) reacts with oxygen
•
CH4 + O2

CO2 + H2O
•
C3H8 + O2

CO2 + H2O
•
C6H12O6 + O2 
CO2 + H2O
• The charcoal used in a grill is basically
carbon. The carbon reacts with oxygen to
yield carbon dioxide. The chemical equation
for this reaction is C + O2  CO2
Exploding Flour
FLOUR
6” PVC Pipe
EXPLOSION:
Fuel (flour)
Ignition (candle)
Oxygen (combustion)
As confinement increases,
EXPLOSION is greater.
Combustion of a Hydrocarbon
GENERAL FORMULA: CH + O2  CO2 + H2O
Many homes get heat from propane (C3H8) heaters.
Write a balanced chemical equation for the complete
combustion of propane gas.
C3H8(g) + O2(g)  CO2(g) + H2O(g)
5 2(g)  CO
3 2(g) +
C3H8(g) + O
4
H2O(g)
+ energy
Combustion of Hydrocarbon (cont.)
Ideal Stoichiometry

C3H8(g) + 5 O2(g)  3 CO2(g) + 4 H2O(g)
+ energy
Too ‘rich’ (not enough oxygen – too much fuel)

C3H8(g) + 3 O2(g)  3 CO (g) + 4 H2O(g)
C3H8(g) + 2 O2(g) 3 C (g) + 4 H2O(g)
SOOT
+ energy
+ energy
Combustion of
Methane Gas
Davis, Metcalfe, Williams, Castka, Modern Chemistry, 1999, page 245
Combustion of
Methane Gas
One methane
molecule
CH4
1 carbon
+
4 hydrogen
Two oxygen
molecules
One carbon
dioxide molecule
2 O2
CO2
+
4 oxygen
Davis, Metcalfe, Williams, Castka, Modern Chemistry, 1999, page 245
=
1 carbon
+
2 oxygen
Two water
molecules
+
2H2O
2 oxygen
+
4 hydrogen
Combustion of Glucose
Kelter, Carr, Scott, Chemistry A World of Choices 1999, page 130
Combustion of Iron
• Formation of Rust
4 Fe + O2  2 Fe2O3
• Thermite Reaction
• underwater welding
• Temp. = ~3500oC
Fe2O3 + 2 Al  2 Fe + Al2O3 + 199 kcal
Combustion of Copper
• Copper burns with a
green color
• Copper forms a
patina (oxide)
– green in color
• CuO2
– black in color
Statue of Liberty is covered with
copper that has oxidized to form
copper (II) oxide, CuO2.
• CuO
Acid/Base Reaction
• An acid and a base react to form a salt and
water.
• Always in aqueous solution
• Acid (H+) + Base (OH-) → Salt + H2O
NaOH + HCl → NaCl + H2O
NH4OH + H2SO4 →
(NH4)2SO4 + H2O
How to recognize which type
• Look at the reactants
• Element(E), Compound(C)
•
•
•
•
•
E+E
C
E+C
C+C
Acid + Base
Synthesis
Decomposition
Single replacement
Double replacement
Acid/Base reaction
• Look at the Products
• CO2 + H2O
Combustion
Redox
Examples
H2 + O2  Synthesis
H2O  Decomposition
AgNO3 + NaCl  Double replacement
Zn + H2SO4  Single replacement
HgO  Decomposition
KBr +Cl2  Single replacement
Mg(OH)2 + H2SO3  Double replacement
Examples
Acid/Base
HNO3 + KOH 
CaPO4  Decomposition
Single replacement
AgBr + Cl2 
Zn + O2 
Synthesis
HgO + Pb Single replacement
HBr + NH4OH Acid/Base
Cu(OH)2 + KClO3  Double replacement
Summary
An equation:
• Describes a reaction
• Must be balanced because to follow Law of
Conservation of Energy
• Can only be balanced by changing the
coefficients.
• Has special symbols to indicate state, and if
catalyst or energy is required.
• Can describe 5 different types of reactions.
Introduction to Stoichiometry
Stoichiometry is…
• Greek for “measuring elements”
Pronounced “stoy kee ahm uh tree”
• Defined as: calculations of the
quantities in chemical reactions,
based on a balanced equation.
• There are 4 ways to interpret a
balanced chemical equation
#1. In terms of Particles
• An Element is made of atoms
• A Molecular compound (made of
only nonmetals) is made up of
molecules (Don’t forget the diatomic elements)
• Ionic Compounds (made of a metal
and nonmetal parts) are made of
formula units
Example: 2H2 + O2 → 2H2O
• Two molecules of hydrogen and one
molecule of oxygen form two molecules of
water.
• Another example: 2Al2O3 Al + 3O2
2 formula units Al2O3 form 4 atomsAl
and 3 molecules O2
Now read this: 2Na + 2H2O  2NaOH + H2
#2. In terms of Moles
• The coefficients tell us how many
moles of each substance
2Al2O3 Al + 3O2
2Na + 2H2O  2NaOH + H2
• A balanced equation is a Molar
Ratio- We will look at this next.
#3. In terms of Mass
• The Law of Conservation of Mass applies
• We can check mass by using moles.
2H2 + O2  2H2O
2 moles H2
1 mole O2
2.02 g H2
1 mole H2
32.00 g O2
1 mole O2
= 4.04 g H2
+
= 32.00 g O2
36.04 gg HH22++ O
O22
36.04
reactants
In terms of Mass (for products)
2H2 + O2  2H2O
2 moles H2O
18.02 g H2O
1 mole H2O
= 36.04 g H2O
36.04 g H2 + O2 = 36.04 g H2O
36.04 grams reactant = 36.04 grams product
The mass of the reactants must
equal the mass of the products.
#4. In terms of Volume
• At STP, 1 mol of any gas = 22.4 L
2H2 + O2  2H2O
(2 x 22.4 L H2) + (1 x 22.4 L O2)  (2 x 22.4 L H2O)
67.2 Liters of reactant ≠ 44.8 Liters of product!
NOTE: mass and atoms are ALWAYS
conserved - however, molecules, formula
units, moles, and volumes will not
necessarily be conserved!
Mole Ratios
 Ratio between two of the substances in a
balanced equation
 Derived from coefficients of any two substances
in an equation.
69
Writing Mole Factors
4 Fe + 3 O2
2 Fe2O3
Fe and O2
4 mol Fe
3 mol O2
and
3 mol O2
4 mol Fe
Fe and Fe2O3
4 mol Fe
2 mol Fe2O3
and
2 mol Fe2O3
4 mol Fe
70
O2 and Fe2O3
3 mol O2
2 mol Fe2O3
and
2 mol Fe2O3
3 mol O2
71
Learning Check
3 H2(g) + N2(g)
2 NH3(g)
A. A mol factor for H2 and N2 is
1) 3 mol N2
2) 1 mol N2
3) 1 mol N2
1 mol H2
3 mol H2
2 mol H2
B. A mol factor for NH3 and H2 is
1) 1 mol H2
2 mol NH3
2) 2 mol NH3
3 mol H2
3) 3 mol N2
2 mol NH3
72
Answers:
3 H2(g) + N2(g)
2 NH3(g)
A. A mol factor for H2 and N2 is
2) 1 mol N2
3 mol H2
B. A mol factor for NH3 and H2 is
2) 2 mol NH3
3 mol H2
73
Converting Moles
2Al2O3 Al + 3O2
– each time we use 2 moles of Al2O3 we will also
make 3 moles of O2
2 moles Al2O3
3 mole O2
or
3 mole O2
2 moles Al2O3
Molar ratios can also be known as
conversion factors. We could use them to
solve calculations.
Example:
• How many moles of O2 are
produced when 3.34 moles of
Al2O3 decompose?
2Al2O3 Al + 3O2
3.34 mol Al2O3
3 mol O2
2 mol Al2O3
= 5.01 mol O2
Conversion factor from balanced equation
If you know the amount of ANY chemical in the reaction,
you can find the amount of ALL the other chemicals!
Practice:
4 Fe + 3 O2
2 Fe2O3
How many moles of Fe2O3 are produced when 6.0
moles O2 react?
6.0 mol O2 x
mol Fe2O3 = 4.0 mol Fe2O3
mol O2
76
More Practice:
4 Fe + 3 O2
2 Fe2O3
How many moles of Fe are needed to react with 12.0
mol of O2?
1) 3.00 mol Fe
2) 9.00 mol Fe
3) 16.0 mol Fe
77
Answer:
4 Fe + 3 O2
2 Fe2O3
4
12.0 mol O2 x
mol Fe = 16.0 mol Fe
3
mol O2
78
More Practice
4 Fe + 3 O2
2 Fe2O3
How many grams of O2 are needed to produce
0.400 mol of Fe2O3?
1) 38.4 g O2
2) 19.2 g O2
3) 1.90 g O2
79
Answer:
0.400 mol Fe2O3 x 3 mol O2 x 32.0 g O2
2 mol Fe2O3
1 mol O2
= 19.2 g O2
80
Converting Mass
 Balance equation
 Convert starting amount to moles
 Use coefficients to write a mol-mol factor
 Convert moles of desired to grams
81
Example:
The reaction between H2 and O2 produces 13.1 g of
water. How many grams of O2 reacted?
Write the equation
H2 (g) + O2 (g)
H2O (g)
Balance the equation
2 H2 (g) + O2 (g)
2 H2O (g)
82
Organize data
2 H2 (g) + O2 (g)
?g
Plan g H2O
Setup
mol H2O
2 H2O (g)
13.1 g
mol O2
O2
13.1 g H2O x 1 mol H2O x 1 mol O2 x 32.0 g O2
8.0 g H2O 2 mol H2O 1 mol O2
= 11.6 g O2
83
Points to Remember
1.
2.
3.
4.
Read an equation in moles
Convert given amount to moles
Use mole factor to give desired moles
Convert moles to grams
grams (given
moles (given)
grams (desired)
moles (desired)
84
Mass-Mass Problem:
6.50 grams of aluminum reacts with an excess of
oxygen. How many grams of aluminum oxide are
formed?
4Al + 3O2  2Al2O3
6.50 g Al
1 mol Al
2 mol Al2O3 101.96 g Al2O3
26.98 g Al
4 mol Al
(6.50 x 1 x 2 x 101.96) ÷ (26.98 x 4 x 1) =
1 mol Al2O3
= ? g Al2O3
12.3 g Al2O3
are formed
Another example:
• If 10.1 g of Fe are added to a solution
of Copper (II) Sulfate, how many
grams of solid copper would form?
2Fe + 3CuSO4  Fe2(SO4)3 + 3Cu
Answer = 17.2 g Cu
Practice
How many O2 molecules will react with 505 grams
of Na to form Na2O?
4 Na + O2
2 Na2O
Complete the set up:
505 g Na x 1 mol Na x ________ x _______
23.0 g Na
87
Answer:
4 Na + O2
2 Na2O
505 g Na x 1 mol Na x 1 mol O2 x 6.02 x 1023
23.0 g Na
4 mol Na 1 mol O2
= 3.30 x 1024 molelcules
88
More Practice:
Acetylene gas C2H2 burns in the oxyactylene torch for
welding. How many grams of C2H2 are burned if the
reaction produces 75.0 g of CO2?
2 C2H2 + 5 O2
4 CO2 + 2 H2O
75.0 g CO2 x _______ x _______ x _______
89
Answer:
2 C2H2 + 5 O2
4 CO2 + 2 H2O
75.0 g CO2 x 1 mol CO2 x 2 mol C2H2 x 26.0 g C2H2
44.0 g CO2 4 mol CO2
1 mol C2H2
= 22.2 g C2H2
90
Volume-Volume Calculations:
• How many liters of CH4 at STP are required
to completely react with 17.5 L of O2 ?
CH4 + 2O2  CO2 + 2H2O
1 mol O2 1 mol CH4 22.4 L CH4
17.5 L O2
22.4 L O2 2 mol O2 1 mol CH4
= 8.75 L CH4
Notice anything relating these two steps?
Avogadro told us:
• Equal volumes of gas, at the same
temperature and pressure contain the
same number of particles.
• Moles are numbers of particles
• You can treat reactions as if they happen
liters at a time, as long as you keep the
temperature and pressure the same.
1 mole = 22.4 L @ STP
Shortcut for Volume-Volume?
• How many liters of CH4 at STP are required
to completely react with 17.5 L of O2?
CH4 + 2O2  CO2 + 2H2O
17.5 L O2
1 L CH4
2 L O2
= 8.75 L CH4
Note: This only works for
Volume-Volume problems.
Limiting Reagants
 If the amounts of two reactants are given, the
reactant used up first determines the amount of
product formed.
 Limiting reagents- the reactant that is used up in a
chemical reaction.
 Excess reagents- the reactant that is left over
after chemical reaction.
94
Example:
Suppose you are preparing cheese sandwiches.
Each sandwich requires 2 pieces of bread and 1
slice of cheese. If you have 4 slices of cheese
and 10 pieces of bread, how many cheese
sandwiches can you make?
95
Cheese Sandwich Products
Sandwich 1
+
+
=
Sandwich 2
+
+
=
96
Practice:
How many sandwiches can you make?
____ slices of bread
+ ____ slices of cheese
= ____ sandwiches
What is left over? ________________
What is the limiting reactant?
97
Practice
How many sandwiches can you make?
__10__ slices of bread
+ __4__ slices of cheese
= __4__ sandwiches
What is left over? _2 slices of bread
What is the limiting reactant? cheese
98
Hints for LR Problems
1. For each reactant amount given, calculate the
moles (or grams) of a product it could produce.
2.The reactant that produces the smaller amount of
product is the limiting reactant.
3. The number of moles of product produced by the
limiting reactant is ALL the product possible. There is
no more limiting reactant left.
99
Limiting Reagents - Combustion
How do you find out which is limited?
• The chemical that makes the least
amount of product is the “limiting
reagent”.
• You can recognize limiting reagent
problems because they will give you 2
amounts of chemical
• Do two stoichiometry problems, one
for each reagent you are given.
• If 10.6 g of copper reacts with
3.83 g
is the
sulfur, how many gramsCu
of the
product (copper
Limiting
(I) sulfide) will be formed?
Reagent,
2Cu + S  Cu
2S
since it
1 mol Cu2S 159.16 g Cu2S
1
mol
Cu
10.6 g Cu
produced less
63.55g Cu 2 mol
Cu
1 mol Cu2S
product.
= 13.3 g Cu2S
1
mol
S
3.83 g S
32.06g S
1 mol Cu2S 159.16 g Cu2S
1 mol S
1 mol Cu2S
= 19.0 g Cu2S
Another example:
• If 10.3 g of aluminum are reacted
with 51.7 g of CuSO4 how much
copper (grams) will be produced?
2Al + 3CuSO4 → 3Cu + Al2(SO4)3
the CuSO4 is limited, so Cu = 20.6 g
• How much excess reagent will
remain?
Excess = 4.47 grams
The Stoichiometric Concept of :
Percent Yield
You prepared cookie dough to make 5 dozen cookies.
The phone rings while a sheet of 12 cookies is baking.
You talk too long and the cookies burn. You throw
them out (or give them to your dog.) The rest of the
cookies are okay.
How many cookies could you have made (theoretical
yield)?
How many cookies did you actually make to eat?
(Actual yield)
105
Types of Yield:
Actual yield is the amount of product actually
recovered from an experiment
Theoretical (possible) yield is the maximum
amount of product that could be produced from
the reactant.
Percent Yield is the actual yield compared to the
maximum (theoretical yield) possible.
106
Details on Yield
• Percent yield tells us how “efficient” a
reaction is.
• Percent yield can not be bigger than
100 %.
• Theoretical yield will always be larger
than actual yield!
– Why? Due to impure reactants; competing side
reactions; loss of product in filtering or
transferring between containers; measuring
Percent Yield Calculation
What is the percent yield of cookies?
Percent Yield = Actual Yield (g) recovered X 100
Possible Yield (g)
% cookie yield = 48 cookies x 100 = 80% yield
60 cookies
108
Example:
• 6.78 g of copper is produced when 3.92 g
of Al are reacted with excess copper (II)
sulfate.
2Al + 3 CuSO4  Al2(SO4)3 + 3Cu
• What is the actual yield?
= 6.78 g Cu
• What is the theoretical yield?
• What is the percent yield?
= 13.8 g Cu
= 49.1 %