Transcript Ch. 7 & 8 Notes (Chemical Reactions) teacher 2013

•
Ch. 7 & 8 Notes -- Chemical Reactions
Chemical equations give information in two major areas:
Reactants
products
1. _____________
and ______________
of the reaction.
amount
2. Coefficients of a balanced chemical equation tell us the ______
of the substances involved.
Example of a Balanced Chemical Equation: 2H2 (g) + O2 (g)  2H2O (g)
Review: Reactants are on the ______
left side of the arrow, and the
right
yields
products are on the __________
side. The arrow means “________”,
or “reacts to produce” when read aloud.
•
•
•
From our example, hydrogen reacts with oxygen in a ___:___
2 1 ratio.
moles or
The coefficients represent either the number of _________
molecules present.
liters if the substances are
The coefficients can also represent _________
gases.
(1 mole = 6.02 x 1023 molecules, 1 mole of gas= 22.4 Liters)
Common Symbols used in Chemical Equations
+
→
(See Table 8.1)
= used to separate 2 reactants or 2 products from each other
= “yields” or “reacts to produce”
= _____________
reaction (like a rechargeable battery)
reversible
(s) (l) (g) (aq) = phase of matter: (solid, liquid, gas, or “aqueous”)
heat
= ___________
supplied to the reaction
MnO2
= a catalyst, (in this case, MnO2), is used to ________
speed ____
up
the reaction.
= _______
gas given off as a product
solid precipitate produced
= ______
Decoding Common Chemical Equation Symbols
Practice Problems: Describe the following reactions using complete
sentences.
a) NaHCO3 (s) + HCl (aq)  NaCl (aq) + H2O (l) + CO2
Solid sodium bicarbonate plus aqueous hydrochloric acid yields
aqueous sodium chloride plus water plus carbon dioxide gas.
b) H2SO4 (aq) + BaCl2 (aq)  HCl (aq) + BaSO4 (s)
Aqueous sulfuric acid plus aqueous barium chloride yields
aqueous hydrochloric acid plus solid barium sulfate.
c) Write a chemical equation from the following description: “Sodium
plus bromine, when heated, reacts to produce solid sodium bromide.”
Na(s) + Br2 (l)  NaBr (s)
Balancing Chemical Equations
Why do you have to balance a chemical equation?
• Law of Conservation of Matter (or Mass):
“Matter is neither
____________
nor _______________
in chemical reactions.”
created
destroyed
joined
• During a chemical reaction, atoms are either _________,
______________,
or rearranged. The _____________
and type of
separated
number
each atom stays the same.
How do you balance a chemical equation?
• __________________
are placed in front of the substances involved
Coefficients
in the chemical reaction to get the same number of atoms of each
element on both sides of the equation. This number will multiply the
number of atoms there are in a formula.
Rules for Balancing Chemical Equations
(1) Coefficients can only be placed ___
in _________
front
of a chemical
formula.
Practice Problems: How many atoms of each type are indicated in
the following compounds?
(a)
2 (NH4)3PO4
6
N= ___
24 P= ___
2 O= ___
8
H= ___
(b)
4 KC2H3O2
4 C= ___
8 H= ___
12 O= ___
8
K= ___
(c)
3 Ca(NO3)2
Ca= ___
3 N= ___
6 O= ___
18
Rules for Balancing Chemical Equations
(2) You cannot change a ________________!!
subscript
Example :
2 H2 +
O2  2 H2O
H2 O2
To balance oxygen, you cannot change water’s formula to_________!
(3) You cannot place the coefficient in the ______________
of a
middle
formula!!
Example :
2Al +
N2  2 AlN
To balance nitrogen, you cannot put a 2 in the middle to make _______.
Al2N
whole
# ratio.
(4) Reduce the coefficients to the simplest ____________
___
Example:
4H2 + 2O2  4H2O can be reduced to…
__H
2 2 + __O
1 2  __H
2 2O
Rules for Balancing Chemical Equations
(5) Get rid of any ____________!
Coefficients must be _________
fractions
whole #’s
• You can’t have a _______________
of a molecule or atom!
fraction
Example: 2 x ( 1H2 + ½O2  1H2O )
__
2 H2 + __O
1 2  __H
2 2O
changes to…
Balancing Equations: “Helpful Hints”
a) Balance elements that appear in more than one compound ________.
last
1
___(NH
4)2CO3 
2
___NH
3
+
1
___CO
2
+
1 2O
___H
ion “___________”
chunks
b) Balance _____
as though it were one item as long as
the ion stays together as a group on each side of the yields arrow.
2
3
___Al
+ ___CuSO
4 
1 2(SO4)3 +
___Al
3
___Cu
start _________
over and begin
c) If you can’t seem to get it balanced, _________
with a different element the next time, or put a “2” somewhere and
then try again.
2
2 2O
___Li
+ ___H

2
1 2
___LiOH
+ ___H
d) This is what I’ll constantly be telling you to do if you are stuck and
you need my help... “Pick an element to balance. How many are on
Fix
it
the left side? How many are on the right side? ________
____!”
2
___Fe(OH)
3 
1 2O3 + ___H
3 2O
___Fe
Balancing Equations: “Helpful Hints”
e) My goofy “balancing song” may help:
2 on the left and a ___on
3
3 on the
“If there’s a ___
the right, you put a ___
left and a ___
2 on the right, (makin’ money!)”
___Al
4
+ ___O
3 2  ___Al
2 2O3
f) If you see only C’s H’s and O’s, balance them in this order: C, H, O.
2 x (___C
H
1
)
+
___O
5/2 2 
___CO
2
2
+
___H
1 2O
___C
2 2H2 +
___O
5 2 
___CO
4
2
+
___H
2 2O
2
2
Five General Types of Reactions
1) ________________________:
Decomposition
•
one ______________
compound into simpler
A reaction that breaks apart ______
substances, (usually two elements or an element and a smaller
compound.)
+
General Form: _____
AX  ___
A + ___
X
H2 + _____
O2
Examples:
H O  _____
2
O2
KCl + _____
KClO3  _____
Remember that “HONClBrIF” elements are diatomic when alone!!
Categories of Decomposition (and Composition ) Reactions
a) carbonates  metallic oxide + CO2
CaO + _____
CO2
CaCO3  _____
b) chlorates  metallic chloride + O2
NaCl + _____
O2
NaClO3  _____
c) hydroxides  metallic oxide + H2O
MgO + _____
H2O
Mg(OH)2  _____
d) oxy acids  nonmetal oxide + H2O
SO3 + _____
H2O
H2SO4  _____
e) binary compounds  2 elements
Na + _____
Cl2
NaCl  _____
•
•
•
Every time you try to write the formula for a new compound, you
charges of the ions and ___________
cross
must look up the ___________
them
if they are different!!
Balance it _________
AFTER you get all the correct formulas written first!
Don’t forget about the HONClBrIF’s!
General Types of Reactions (Continued)
3) _____________
Single
Replacement:
•
one ______________
compound
one ____________
element
A reaction between ____
and ___
that produces a different _____________
compound
and ______________.
element
General Forms: ____
AX + __
Y  ____
AY
+
+ __
X
AX
B  ____
BX + __
A
____ + __
+
•
•
•
•
more
The element that is trying to replace the other must be ________
reactive
_______________
than the one it is replacing.
You must use the Activity Series to see if the reaction will happen.
Table 8.2
_________
Higher ___
up = more reactive
Elements from ____
Li to ____
Na can displace hydrogen in water to
form a metallic hydroxide and H2 gas.
General Types of Reactions (Continued)
2) _______________:
Composition
(sometimes called “Combination” or “Synthesis”)
•
•
two __________________,
substances
A reaction of _____
typically a metal and a
nonmetal to form ______
one ______________.
compound
It is the opposite of decomposition. (The same categories of
reactions from above apply, just in reverse.)
+
General Form:
Examples:
A + ___
X  _____
AX
___
Al
+
AlCl3
Cl2  _______
2 elements
 binary compound
PbO + H2O 
metallic + water 
oxide
Pb(OH)2
______
hydroxide
Activity Series
Single Replacement Reactions
Examples:
NaCl
NaF + _____
Cl2
+ F2  _____
FeCl2 +
KCl + _____
Fe
K  _____
HCl
+
H2
Zn  ZnCl
_____
2 + _____
HCl
+
no reaction
Au  _____
+ _____
Na 
H2
_____
NaOH + _____
Fe 
no reaction
_____
+ _____
AgNO3 + Cu 
CuNO
Ag
_____3 + _____
H2O +
H(OH)
H2O +
General Types of Reactions (Continued)
4) _______________
Double
Replacement: (sometimes called “Ionic”)
•
•
two ________________
compounds
A reaction between _____
that are dissolved in
water that produces _____
two ________________
compounds
, one of which is
________________.
insoluble
Water or a gas may be one of the two compounds being produced.
BX(s)
(aq)+ BY
(aq)  AY
(aq) + ____
General Form: AX
____
____
____
+
•
•
+
You must use the Solubility Chart to see which product is the
precipitate.
Solubility Chart Key
Examples: CaCl2 (aq)
I or _____=
sS
___
precipitate
3)2 (aq) + ________
AgCl
+ AgNO3 (aq)  Ca(NO
_________
NaCl (aq) + ________
H2O (l)
NaOH (aq) + HCl (aq)  ________
Double Replacement Reaction
General Types of Reactions (Continued)
5) _________________:
Combustion
•
•
•
A reaction between a Carbon/Hydrogen (and sometimes Oxygen)
O2
_________________
compound
with _____.
CO2 + ________
H2O
The products are always the same… ________
This reaction is too easy!! Don’t miss it!
General Form:
Examples:
C2H2
CO2 + ____
H2O
CxHy + O2  ____
+
C7H6O +
CO2 + _______
H2O
O2  _______
CO2 + _______
H2O
O2  _______
Writing Net Ionic Equations for Double Replacement Reactions
•
•
A “net ionic equation” only shows the _________
ions
that were used to
make the precipitate.
Some ions were always dissolved in water. These are called
“________________
ions”. (They don’t do anything, so we can
spectator
ignore them.)
Example: CaCl2 (aq) + 2AgNO3 (aq)  Ca(NO3)2 (aq) + 2AgCl (s)
Ionic Equation Written as Ions Dissolved in Water:
+2
−
+
Ca
2NO3 −(aq)  ___
Ca+2(aq) + _____
2NO3 −(aq) + _________
2AgCl (s)
___ (aq) + 2Cl
___ (aq) + 2Ag
___ (aq) + _____
•
Cancel out the spectator ions, and you are left with the Net Ionic
Equation!
−
+
2Cl
2Ag
2AgCl(s)
(aq)
(aq)
________ + _________
 __________
Writing Net Ionic Equations for Double Replacement Reactions
Practice Problem: Write the net ionic equation for the following reaction.
KNO3 (aq)+ _________
BaCO3
K2CO3 (aq) + Ba(NO3)2 (aq)  _________
−2
+2
CO
3
(aq)
+
Ba
BaCO3
(aq)
Net Ionic Equation = ____________________________________
Electron Transfer and Redox Reactions
An oxidation-reduction reaction, or redox reaction
involves the transfer of electrons from one atom to
another.
In a synthesis reaction, neutral elements become ions
(charged particles) and then form ionic bonds.
Complete these synthesis reactions:
2 Na + Cl2 
Mg + Cl2 
2 NaCl
MgCl2
(Na+
Cl-)
(Mg+2 Cl-)
Reactions where electrons are transferred and charges change
are called oxidation-reduction reactions. (redox)
To represent the change of a sodium atom into a sodium ion
we write:
Na  Na+ + eThis change is the result of the sodium atom losing an
electron. To represent this loss, we write the electron as a
separate product in the equation.
Oxidation is defined as the loss of electrons.
Since the electron must eventually be gained by a different
atom, we call this a half-reaction.
To represent the change of a chlorine atom into a chloride
ion we write:
Cl + e-  ClThis change is the result of the chlorine atom gaining an
electron. To represent this gain, we write the electron as a
separate reactant in the equation.
Reduction is defined as the gain of electrons.
Since chlorine exists as a diatomic molecule, the reduction
half-reaction for chlorine is actually:
Cl2 + 2 e-  2 Cl-
Here’s a way to remember the names of half-reactions:
“LEO the lion says GER”
Loss of
Electrons is
Oxidation
Gain of
Electrons is
Reduction
Trends you already know:
Metals generally lose electrons in synthesis reactions.
Complete and balance the oxidation half-reactions for these
metals.
K  K+ + eAl  Al+3 + 3 eNon-metals generally gain electrons in synthesis reactions.
Complete and balance the reduction half-reactions for these
non-metals.
2 e- + Br2  2 Br-
2 e- + S  S-2
For each of the following oxidation-reduction reactions,
identify which element is oxidized and which is reduced.
Ba
+
oxidized
O2
reduced
F2

BaF2
reduced
+
2 Sr
oxidized

2 SrO
Oxidizing and
Reducing Agents
2K(s) + Br2(g)  2 KBr(s)
• K  oxidized
– K is also called the reducing agent (electron donor).
– The reducing agent reduces something else.
• Br  reduced
– Br2 is also called the oxidizing agent (electron acceptor).
– The oxidizing agent oxidizes something else.
Voltaic Cells: Simple Batteries
The below picture is a voltaic cell using Zinc
and Copper
Anode: Oxidation
Zn  Zn+2 + 2e-
Cathode: Reduction
2e- + Cu+2  Cu
In each reaction, identify the atoms that are oxidized and
reduced, then label the oxidizing agent and the reducing agent.
1. Na + Cl2  NaCl
Oxidized = Na
Reduced = Cl2
Oxidizing Agent = Cl2
Reducing Agent = Na
2. Mn + S8  MnS
Oxidized = Mn
Reduced = S8
Oxidizing Agent = S8
Reducing Agent = Mn
Cleaning Supplies
Cleaning supplies are often oxidizing agents and reducing agents.
When mixed, cleaning supplies can make a dangerous gas that will
poison your family.
DON’T MIX THEM!!!
Balancing Redox Reactions
When balancing redox reactions you must consider the charge
as well as the atoms, since the total number of electrons must be
conserved in the reaction.
Below please break down the following into half-reactions to
balance:
Zn + Ag+ 
AlN+
Zn2+ + Ag
Mg  Al +
Mg3N2