Transcript Chapter 11 Chemical Reactions

Chapter 11
Chemical Reactions
Hingham High School
Mr. Dan Clune
All chemical reactions
• Two parts:
•
•
– what you start with
what you end with
• Reactants turn into the products.
• Reactants  Products
In a chemical reaction
• The way atoms are joined is
changed.
• Atoms aren’t
or
.
In a chemical reaction
• Can be described several ways:
1. In a sentence
Copper
with chlorine to
copper (II) chloride.
2. In a word equation
Copper + chlorine  copper (II) chloride
Symbols in equations-p.323
Cu + Cl2  CuCl2
•
() separates reactants from
products
• Read “reacts
”
• Plus (+) sign read “
”
Symbols used in equations
Cu(s) + Cl2(g)  CuCl2(s)
•
•
•
•
(s) =
(g) =
(l) =
(aq) - dissolved in water,
an
solution.
Symbols used in equations
• ↑ after product, indicates
produced
•same as (g) - H2↑
•after product, indicates
produced
•same as (s) - PbI2↓
Symbols used in equations
indicates
reaction

heat
shows that
is supplied to the reaction
Pt
  is - indicates a
is supplied, in this case, platinum.
  ,   
What is a catalyst?
•
up a reaction
• Is NOT
or
by the reaction.
• Enzymes are biological or protein
catalysts.
Skeleton Equation
• Uses formulas to describe a reaction
• doesn’t indicate how many.
•
Convert this to an equation
• Solid iron (III) sulfide reacts with
gaseous hydrogen chloride to form
iron (III) chloride and
hydrogen sulfide gas.
Nitric acid dissolved in water reacts with
solid sodium carbonate to form liquid
water, carbon dioxide gas,
sodium nitrate dissolved in water.
Now, read these:
• Fe(s) + O2(g)  Fe2O3(s)
• Cu(s) + AgNO3(aq) 
Ag(s) + Cu(NO3)2(aq)
• NO2 (g)
Pt
 
N2(g) + O2(g)
Balancing Chemical
Equations
Balanced Equation
• Atoms can’t be created or destroyed
• All the atoms we
with
we must
up with
• A balanced equation has the
of each element on
of the equation.
C
+
O
O

O C
O
•
C + O2  CO2
• This equation is already balanced
• What if it isn’t?
C
+
O
O

C
O
•
C + O2 
CO
• Need one more O in the
.
• Can’t change the
,
because it describes what it is (carbon
monoxide in this example)
C
+
O
O

C
O
C
O
• Must be used to make another
• But where did the other C come from?
C
+
C
O
O

C
O
C
O
• Must have started with two C
•
Finding the number of atoms
• The subscript in front of an element is the
number of atoms of that element / polyatomic.
• Ex) CO2
B2(SO4)3
• C=
O=
B=
SO4=
• A coefficient in front of the formula multiplies
the amount of elements by the coefficient.
• Ex) 3CO2
2B2(SO4)3
• C=
O=
B=
SO4=
Finding the number of atoms
H2 O
H=
O=
2H2O
H=
O=
B(NO3)2
B=
NO3=
3B(NO3)2
B=
NO3=
Pb3(PO4)4
Pb=
PO4=
2Pb3(PO4)4
Pb=
PO4=
Rules for balancing:
. Determine the
for reactants and products.
. Write a
equation.
Rules for balancing:
3. Count the
of atoms
of each
appearing
on
sides of the equation.
4. Balance the elements one at a
time by adding
(the numbers in front)
-Save
and
until LAST!
Rules for balancing:
5. Check to make sure it is balanced.
6. Make sure the coefficients are
in the
possible ratio.
Don’t you ever…
• Never change a
to balance an equation.
• H2O is a different compound than H2O2
• Never put a coefficient in the
of a formula
X
• 2 NaCl is okay, Na2Cl is not.
Example
H2 + O2 
R
H2O
P
H
O
2 H2 + O2  2 H2O
AgNO3 + Cu  Cu(NO3)2 + Ag
R
P
Ag
Cu
NO3
Mg + N2  Mg3N2
R
P
Mg
N
P + O2  P4O10
R
P
P
O
Na + H2O  H2 + NaOH
R
P
Na
H
O
CH4 + O2  CO2 + H2O
R
P
C
H
O
Section 8.2
Types of Chemical Reactions
• OBJECTIVES:
• Identify a reaction as combination,
decomposition, single-replacement, doublereplacement, or combustion
Types of Reactions
• 5 major types.
• predict the products
• predict whether or not they will happen at all
• How? We recognize them by their
#1 - Combination Reactions
• Combine - put together
•
substances combine to
make one
.
A + B 
• Ca +O2 CaO
• SO3 + H2O  H2SO4
#1 - Combination Reactions
• We can predict the products if
they are two elements.
Mg + N2 
Write and balance
Ca + Cl2 
Fe + O2  iron (II) oxide
Al + O2 
Remember that the first step is to write
the correct formulas
• Then balance by using
only
•
•
•
•
#2 - Decomposition Reactions
• Decompose = fall apart
•
reactant falls apart into
elements or
compounds.
AB  A + B
#2 - Decomposition Reactions
•NaCl   
Na + Cl2

•CaCO3   CaO + CO2
•Note that
is
usually required to
decompose
electricity
#2 - Decomposition Reactions
• Binary compounds (made of 2 elements)
falls apart into its elements
•
H 2O
• HgO
electricity



 
#3 - Single Replacement Reactions
•
element
another (new dance partner)
• Reactants are an
• Products will be a
element and different cmpd
• Li + KCl  K + LiCl
(Cations switched)
• F2 +2 LiCl  2LiF + Cl2
(Anions switched)
#3 Single Replacement Reactions
• Metals replace other metals (and H)
• K + AlN 
• Zn + HCl 
• Think of water as: HOH
• Metals replace first H, then combines w/
hydroxide (OH).
• Na + H2O 
#3 Single Replacement Reactions
• Sometimes, the reaction will
Some chemicals are more “
than others
•
active replaces
happen:
”
active
The “Activity Series” of Metals
Higher • Lithium
activity • Potassium
• Calcium
• Sodium
• Magnesium
• Aluminum
• Zinc
• Chromium
• Iron
• Nickel
• Lead
• Hydrogen
• Bismuth
• Copper
• Mercury
• Silver
Lower
• Platinum
activity
• Gold
If the lone metal is
the paired metal,
replacement
occur.
Ex) Li + NaCl  Na + LiCl
If the lone metal is
the paired metal,
replacement will
occur.
Ex) Na + LiCl  Na + LiCl
The “Activity Series” of Halogens
Higher Activity
Fluorine
Chlorine
Bromine
Iodine
If the lone halogen is
the paired
halogen, replacement
occur.
Lower Activity
2NaCl(s) + F2(g) 
MgCl2(s) + Br2(g) 
#3 Single Replacement Reactions Practice:
Higher • Lithium
activity • Potassium
• Calcium
• Sodium
• Magnesium
• Aluminum
• Zinc
• Chromium
• Iron
• Nickel
• Lead
• Hydrogen
• Bismuth
• Copper
• Mercury
• Silver
Lower
• Platinum
activity
• Gold
• Fe + CuSO4 
• Pb + KCl 
•
Al + HCl 
#4 - Double Replacement Reactions
•
things
• Reactants must be two
compounds.
each other.
• NaOH + FeCl3 
• positive ions change place
• NaOH + FeCl3 Fe+3 OH- + Na+1 Cl-1
=
Complete and balance:
• assume all of the following reactions
actually take place:
CaCl2 + NaOH 
CuCl2 + K2S 
KOH + Fe(NO3)3 
K2SO4 + BaF2 
How to recognize which type?
• Look at the reactants:
El + El = Combination
Cpd
= Decomposition
El + Cpd
= Single replacement
Cpd + Cpd = Double replacement
Practice Examples:
• H2 + O2 
• H2O 
• Zn + H2SO4 
• HgO 
• KBr + Cl2 
• AgNO3 + NaCl 
• Mg(OH)2 + H2SO3 
#5 – Combustion Reactions
• Combustion means “
”
• Normally, a cpd composed of only C, H,
(sometimes O) is reacted with oxygen –
called “burning”
•
combustion, products are
Combustion Reaction Examples:
• C4H12 + O2 
•
C6H12O6 +
•
C8H8 +
O2 
O2 
SUMMARY: An equation...
•
•
•
•
Describes a rxn
Must be
only balance by changing
special symbols to indicate physical state,
catalyst or energy required, etc.
Reactions
• 5 major types
• We can tell what type they are by looking at
• Single Replacement happens based on the
Series
Section 11.3
p. 342
Reactions in Aqueous Solution
NiCl2
Co(NO3)2
K2Cr2O7
K2CrO4
CuSO4
KMnO4
Net Ionic Equations
• Many reactions occur in water,
or
solution
• When dissolved in water, many ionic
cpds “dissociate”, or
,
into cations & anions
• Now write ionic equation
Net Ionic Equations
• Example (needs to be a double replacement reaction)
AgNO3 + NaCl  AgCl + NaNO3
1. this is the full balanced equation
2. next, write it as ionic equation by splitting
the cpds into their ions:
Ag1+ + NO31- + Na1+ + Cl1- 
AgCl(s) + Na1+ + NO31Solids do not split up.
Net Ionic Equations
3. Crossing out ions that did not change
(called spectator ions)
Ag1+ + Cl1-  AgCl (s)
This is the
Predicting the Precipitate
• Insoluble salt is a precipitate
• i.e. a solid
• General solubility rules are found:
a) Table 11.3, p. 344 in textbook
Solubility Rules
BaCl2 +
AgNO3 →
1. Break up into ions
2. Now write the net ionic equation.
NaCl + Ba(NO3)2 →
1. Break up into ions
2. Now write the net ionic equation.
PbCl2 (s) + Li2O 
1. Break up into ions
2. Now write the net ionic equation.