Transcript Chemical Reactions - Johnston County Schools

Chemical Reactions
Types of Reactions
•
There are five types of chemical
reactions we will talk about:
1.
2.
3.
4.
5.
•
Synthesis reactions
_____________ reactions
Single displacement reactions
________________ reactions
Combustion reactions
You need to be able to identify the type
of reaction and predict the product(s)
Steps to Writing Reactions
•
Some steps for doing reactions
1.
2.
3.
Identify the type of reaction
Predict the product(s) using the type of
reaction as a model
Balance it
Don’t forget about the diatomic elements!
(HOFBrINCl) For example, Oxygen is O2 as an
element.
In a compound, it can’t be a diatomic element
because it’s not an element anymore, it’s a
compound!
Synthesis reactions
•
•
Synthesis reactions occur when two substances
(generally elements) combine and form a
compound. (Sometimes these are called
combination or addition reactions.)
reactant + reactant  1 product
Basically: A + B  AB
•
•
•
•
Example: 2H2 + O2  2H2O
Example: C + O2  CO2
Sodium and Chlorine Reaction.mov
Formation of AlBr3.MOV
Synthesis Reactions
•
Here is another example of a synthesis
reaction
Practice
•
•
•
•
Predict the products. Write and balance
the following synthesis reaction equations.
Sodium metal reacts with chlorine gas
Na(s) + Cl2(g) 
Solid Magnesium reacts with fluorine gas
Mg(s) + F2(g) 
Aluminum metal reacts with fluorine gas
Al(s) + F2(g) 
Decomposition Reactions
•
Decomposition reactions occur when a
compound breaks up into the elements or
in a few to simpler compounds
1 Reactant  Product + Product
In general: AB  A + B
Example: 2 H2O  2H2 + O2
Example: 2 HgO  2Hg + O2
•
Decomposition of HgO.mov
•
•
•
•
Decomposition Reactions
•
Another view of a decomposition reaction:
Decomposition Exceptions
•
Carbonates and chlorates are special case
decomposition reactions that do not go to
the elements.
•
Carbonates (CO32-) decompose to carbon
dioxide and a metal oxide
•
•
Example: CaCO3  CO2 + CaO
Chlorates (ClO3-) decompose to oxygen gas
and a metal chloride
•
Example: 2 Al(ClO3)3  2 AlCl3 + 9 O2
Decomposition of Water


Water cannot be decomposed by heating
(it boils, but does not change into a new
substance).
Instead, an electrical current must be
passed through an acidified sample of the
water to break it down into hydrogen and
oxygen gases.
Electrolysis of Water

Electrolysis of Water.MOV
Practice
•
•
•
Predict the products. Then, write and
balance the following decomposition
reaction equations:
Solid Lead (IV) oxide decomposes
PbO2(s) 
Aluminum nitride decomposes
AlN(s) 
Practice
Identify the type of reaction for each of the
following synthesis or decomposition
reactions, and write the balanced equation:
N2(g) + O2(g)  Nitrogen monoxide
BaCO3(s) 
Co(s)+ S(s)  (make Co be +3)
NH3(g) + H2CO3(aq) 
NI3(s) 
Single Replacement Reactions
•
•
•
Single Replacement Reactions occur when one
element replaces another in a compound.
A metal can replace a metal (+) OR
a nonmetal can replace a nonmetal (-).
element + compound product + product
A + BC  AC + B (if A is a metal) OR
A + BC  BA + C (if A is a nonmetal)
(remember the cation always goes first!)
When H2O splits into ions, it splits into
H+ and OH- (not H+ and O-2 !!)
Sodium and Potassium Plus Water.mov
Single Replacement Reactions
•
Another view:
Single Replacement Reactions
Write and balance the following single
replacement reaction equation:
• Zinc metal reacts with aqueous
hydrochloric acid
Zn(s) + 2 HCl(aq)  ZnCl2 + H2(g)
Note: Zinc replaces the hydrogen ion in the
reaction
•
Single Replacement Reactions
•
Sodium chloride solid reacts with fluorine gas
2 NaCl(s) + F2(g)  2 NaF(s) + Cl2(g)
Note that fluorine replaces chlorine in the compound
•
Aluminum metal reacts with aqueous copper
(II) nitrate
Al(s)+ Cu(NO3)2(aq)
Double Replacement Reactions
•
•
•
Double Replacement Reactions occur
when a metal replaces a metal in a compound
and a nonmetal replaces a nonmetal in a
compound
Compound + compound  product +
product
AB + CD  AD + CB
Double Replacement Reactions
•
•
•
Think about it like “foil”ing in algebra, first and
last ions go together + inside ions go together
Example:
AgNO3(aq) + NaCl(s)  AgCl(s) + NaNO3(aq)
Another example:
K2SO4(aq) + Ba(NO3)2(aq)  2 KNO3(aq) + BaSO4(s)
Practice
•
Predict the products. Balance the equation
5.
HCl(aq) + AgNO3(aq) 
CaCl2(aq) + Na3PO4(aq) 
Pb(NO3)2(aq) + BaCl2(aq) 
FeCl3(aq) + NaOH(aq) 
H2SO4(aq) + NaOH(aq) 
6.
KOH(aq) + CuSO4(aq) 
1.
2.
3.
4.
Combustion Reactions
•
•
Combustion reactions
occur when a hydrocarbon
reacts with oxygen gas.
This is also called
burning!!! In order to burn
something you need the 3
things in the “fire
triangle”:
1) A Fuel (hydrocarbon)
2) Oxygen to burn it with
3) Something to ignite the
reaction (spark)
Combustion Reactions
•
•
•
In general:
CxHy + O2  CO2 + H2O
Products in combustion are
ALWAYS carbon dioxide and
water. (although incomplete
burning does cause some byproducts like carbon monoxide)
Combustion is used to heat
homes and run automobiles
(octane, as in gasoline, is C8H18)
Combustion
Reactions
Edgar Allen Poe’s
drooping eyes and
mouth are potential
signs of CO
poisoning.
Combustion
•
Example
•
•
C5H12 + 8 O2  5 CO2 + 6 H2O
Write the products and balance the
following combustion reaction:
•
C10H22 + O2 
Combustion



Substances other than hydrocarbons can
also combust. However, you may not be
able to tell whether it’s combustion from
the chemical equation alone.
Remember that combustion must have O2
as a reactant and must release
(exothermic) heat and light energy.
Reactions with O2.mov
Mixed Practice
•
1.
2.
3.
4.
5.
State the type, predict the products, and
balance the following reactions:
BaCl2 + H2SO4 
C6H12 + O2 
Zn + CuSO4 
Cs + Br2 
FeCO3 
Total Ionic Equations





Once you write the molecular equation
(synthesis, decomposition, etc.), you should
check for reactants and products that are
soluble or insoluble.
We usually assume the reaction is in water
We can use a solubility table to tell us what
compounds dissolve in water.
If the compound is soluble (does dissolve in
water), then splits the compound into its
component ions
If the compound is insoluble (does NOT dissolve
in water), then it remains as a compound
Solubility Rules
Soluble Salts
1. The Na+, K+, and NH4+ ions form soluble salts. Thus, NaCl, KNO3, (NH4)2SO4, Na2S, and (NH4)2CO3
are soluble.
2. The nitrate (NO3-) ion forms soluble salts. Thus, Cu(NO3)2 and Fe(NO3)3 are soluble.
3. The chloride (Cl-), bromide (Br-), and iodide (I-) ions generally form soluble salts. Exceptions to this
rule include salts of the Pb2+, Hg22+, Ag+, and Cu+ ions. ZnCl2 is soluble, but CuBr is not.
4. The sulfate (SO42-) ion generally forms soluble salts. Exceptions include BaSO4, SrSO4, and PbSO4,
which are insoluble, and Ag2SO4, CaSO4, and Hg2SO4, which are slightly soluble.
Insoluble Salts
1. Sulfides (S2-) are usually insoluble. Exceptions include Na2S, K2S, (NH4)2S, MgS, CaS, SrS, and
BaS.
2. Oxides (O2-) are usually insoluble. Exceptions include Na2O, K2O, SrO, and BaO, which are soluble,
and CaO, which is slightly soluble.
3. Hydroxides (OH-) are usually insoluble. Exceptions include NaOH, KOH, Sr(OH)2, and Ba(OH)2,
which are soluble, and Ca(OH)2, which is slightly soluble.
4. Chromates (CrO42-) are usually insoluble. Exceptions include Na2CrO4, K2CrO4, (NH4)2CrO4, and
MgCrO4.
5. Phosphates (PO43-) and carbonates (CO32-) are usually insoluble. Exceptions include salts of the
Na+, K+, and NH4+ ions.
Solubilities Not on the Table!


Gases only slightly dissolve in water
Strong acids and bases dissolve in water




Hydrochloric, Hydrobromic, Hydroiodic, Nitric,
Sulfuric, Perchloric Acids
Group I hydroxides (should be on your chart anyway)
Water slightly dissolves (ionizes) in water! (H+
and OH-)
There are other tables and rules that cover more
compounds than your table, but for the most
part you can use the table in reference packet.
Total Ionic Equations
Molecular Equation:
K2CrO4 + Pb(NO3)2 
PbCrO4 + 2 KNO3
Soluble
Insoluble
Soluble
Soluble
Total Ionic Equation:
2 K+ + CrO4 -2 + Pb+2 + 2 NO3- 
PbCrO4 (s) + 2 K+ + 2 NO3-
Net Ionic Equations
These are the same as total ionic
equations, but you should cancel out ions
that appear on BOTH sides of the equation
Total Ionic Equation:
2 K+ + CrO4 -2 + Pb+2 + 2 NO3- 
PbCrO4 (s) + 2 K+ + 2 NO3Net Ionic Equation:
CrO4 -2 + Pb+2  PbCrO4 (s)

Net Ionic Equations

Try this one! Write the molecular, total ionic, and net
ionic equations for this reaction: Silver nitrate reacts
with Lead (II) Chloride in hot water.
Molecular:
Total Ionic:
Net Ionic:
“Oxidation-Reduction Reactions”
LEO SAYS GER
Oxidation and Reduction (Redox)
Early chemists saw “oxidation”
reactions only as the combination of
a material with oxygen to produce
an oxide.
• For example, when methane
burns in air, it oxidizes and forms
oxides of carbon and hydrogen.
Oxidation and Reduction
(Redox)
But, not all oxidation processes that use
oxygen involve burning:
•Elemental iron slowly oxidizes to
compounds such as iron (III) oxide,
commonly called “rust”
•Bleaching stains in fabrics
•Hydrogen peroxide also releases oxygen
when it decomposes
Oxidation and Reduction
(Redox)
A process called “reduction” is the opposite of
oxidation, and originally meant the loss of oxygen
from a compound
Oxidation and reduction always occur
simultaneously
The substance gaining oxygen (or losing
electrons) is oxidized, while the substance losing
oxygen (or gaining electrons) is reduced.
Transfer of Electrons
Today, many of these reactions may
not even involve oxygen
Redox currently says that electrons
are transferred between reactants
Mg
+
S→
Mg2+
+
S2-
(MgS)
•The magnesium atom (which has zero charge) changes to a
magnesium ion by losing 2 electrons, and is oxidized to Mg2+
•The sulfur atom (which has no charge) is changed to a
sulfide ion by gaining 2 electrons, and is reduced to S2-
Assigning Oxidation Numbers
0
1
0
1
2 Na  Cl 2  2 Na Cl
Each sodium atom loses one electron:
1
0
Na  Na  e

Each chlorine atom gains one electron:
0

1
Cl  e  Cl
LEO says GER :
Lose Electrons = Oxidation
1
0
Na  Na  e

Sodium is oxidized
Gain Electrons = Reduction
0

1
Cl  e  Cl
Chlorine is reduced
LEO says GER :
- Losing electrons is oxidation, and the
substance that loses the electrons is
called the reducing agent.
- Gaining electrons is reduction, and
the substance that gains the electrons is
called the oxidizing agent.
Mg is the
reducing
agent
Mg is oxidized: loses e-, becomes a Mg2+ ion
Mg(s) + S(s) → MgS(s)
S is the oxidizing agent
S is reduced: gains e- = S2- ion
It is easy to see the loss and gain of
electrons in ionic compounds, but what
about covalent compounds?
In water, oxygen is highly
electronegative, so:
the oxygen gains electrons (is
reduced and is the oxidizing agent),
and the hydrogen loses electrons (is
oxidized and is the reducing agent)
Not All Reactions are Redox Reactions
- Reactions in which there has been no
change in oxidation number are NOT
redox reactions.
Examples:
1 5 2
1
1
1
1
1 5 2
Ag N O3 (aq)  Na Cl (aq)  Ag Cl ( s)  Na N O3 (aq)
1 2 1
1
6 2
1
6 2
1
2
2 Na O H (aq)  H 2 S O 4 (aq)   Na 2 S O 4 (aq)  H 2 O(l )
Assigning Oxidation Numbers
• An “oxidation number” is a positive or
negative number assigned to an atom
to indicate its degree of oxidation or
reduction.
• Generally, a bonded atom’s oxidation
number is the charge it would have if
the electrons in the bond were
assigned to the atom of the more
electronegative element
Rules for Assigning Oxidation Numbers
1) The oxidation number of any
uncombined element is zero.
2) The oxidation number of a
monatomic ion equals its charge.
0
0
1
1
2 Na  Cl 2  2 Na Cl
Rules for Assigning Oxidation Numbers
3) The oxidation number of oxygen in
compounds is -2, except in
peroxides, such as H2O2 where it is -1.
4) The oxidation number of hydrogen in
compounds is +1, except in metal
hydrides, like NaH, where it is -1.
1
2
H2O
Rules for Assigning
Oxidation Numbers
5) The sum of the oxidation numbers of the
atoms in the compound must equal 0.
1
2
H2O
2(+1) + (-2) = 0
H
O
2
2 1
Ca(O H ) 2
(+2) + 2(-2) + 2(+1) = 0
Ca
O
H
Rules for Assigning Oxidation Numbers
6) The sum of the oxidation numbers in
the formula of a polyatomic ion is equal
to its ionic charge.
? 2
N O3

X + 3(-2) = -1
N
O
thus X = +5
? 2
S O4
2
X + 4(-2) = -2
S
O
thus X = +6
Reducing Agents and Oxidizing
Agents
•An increase in oxidation number = oxidation
• A decrease in oxidation number = reduction
1
0
Na  Na  e

Sodium is oxidized – it is the reducing agent
0

1
Cl  e  Cl
Chlorine is reduced – it is the oxidizing agent
Trends in Oxidation and
Reduction
Active metals:
Lose electrons easily
Are easily oxidized
Are strong reducing agents
Active nonmetals:
Gain electrons easily
Are easily reduced
Are strong oxidizing agents
Other Types of Reactions as
Redox?






Synthesis Reactions from elements are
Redox. Ex: 2Al (s) + 3Cl2 (g)  2AlCl3 (s)
Decomposition Reactions to elements are
Redox. Ex: 2H2O (l)  2H2 (g) + O2 (g)
All Single Replacement Reactions are Redox.
No Double Replacement Reactions are
Redox.
All Combusiton Reactions are Redox.
Dehydration of Sugar.MOV