Transcript Cl -1

Oxidation Numbers,
Reactions in Aqueous
Solutions, & Predicting
Products
I. Rules for Assigning
Oxidation Numbers
1. The oxidation number of any uncombined element is 0.
2. The oxidation number of a monatomic ion equals the charge on
the ion.
3. The more-electronegative element in a binary compound is
assigned the number equal to the charge it would have if it were
an ion.
4. The oxidation number of fluorine in a compound is always -1.
5. Oxygen has an oxidation number of -2 unless it is combined
with F (when it is +2), or it is in a peroxide (such as H2O2 or
Na2O2), when it is -1.
6. The oxidation state of hydrogen in most of its compounds is +1
unless it is combined with a metal, in which case it is -1.
Oxidation vs. Reduction

The sum of the oxidation states in a neutral
compound must equal zero and must be equal to
the overall charge in an ionic compound.
Assign oxidation states to each of the atoms in the following
compounds:
=
a. FeF2 Fe = +2, F = -1 b. H2O H = +1, O = -2 c. KMnO4 K+ = +1, Mn
1(+2) 2(-1)
d. C2H6
2(+1) 1(-2)
C = -3, H = +1
2(-3) 6(+1)
1(+1) +7
e. ICl5 I = +5, Cl = -1
+5
5(-1)
f. SO42-
0
+3 -2
7, O = 2
S = +6, O = -2
+6 4(-2) = -2
Assign oxidation states to each atom in the equation.
Fe2O3 + 2 Al  Al2O3 + 2 Fe Fe gains electrons.
+3 -2
4(-2)
It has been REDUCED
0
Al lost electrons. It has been OXIDIZED
II. REDOX Reactions
Remember LEO GER:
LOSE ELECTRONS OXIDATION
GAIN ELECTRONS REDUCTION
Something that is reduced is called an oxidizing agent.
Something that is oxidized is called a reducing agent.
For each reaction, identify that atoms that undergo reduction, or oxidation.
a.
2 H2
0
b.
Zn
0
c.
(g)
(s)
2 AgCl
+1 -1
+ O2
0
+ Cu2+
+2
(s)
(g)
 2 H2O (g)
oxidized: hydrogen (0 to +1)
reduced: oxygen (0 to -2)
(aq)
 Zn2+
+ Cu
+ H2
0
(g)
+1 -2
+2
(aq)
 2 H+
+1
0
(aq)
(s)
+ 2 Ag
0
oxidized: zinc (0 to +2)
reduced: copper (+2 to 0)
(s)
+ 2 Cl-
(aq)
-1 oxidized: hydrogen (0 to +1)
reduced: silver (+1 to 0)
d. 2 MnO4-(aq) + 16 H+(aq) + 5 C2O42-(aq)  2 Mn2+(aq) + 10 CO2(g) + 8 H2O(l)
+7 -2
+1
+3 -2
+2
+4 -2
oxidized: carbon (+3 to +4)
reduced: manganese (+7 to +2)
+1 -2
Learning Check
Assign oxidation states to each of the atoms in the following
compounds:
a. SO2
b. S
c. SO3
Assign oxidation states to each atom in the equation. Determine
which element was oxidized and which was reduced.
2 Al + 3 CuCl2  2 AlCl3 + 3 Cu
III. Dissociation
When a soluble ionic salt dissolves in water the ions
separate and a hydration shell is formed around each ion
(Dissociation)
Ex. Ba(NO3)2 in water –
Ba(NO3)2 (aq)  Ba+2(aq) + 2 NO3-1(aq)
Aluminum acetate
Al(C2H3O2)3 (aq)  Al+3(aq)+ 3 C2H3O2-1(aq)
Aluminum carbonate
Al2(CO3)3 (s)  Al2(CO3)3 (s) or No RXN
Draw a beaker of dissociated sodium chloride;
a beaker of dissociated Aluminum nitrate;
and a beaker of silver chloride.
Na+1
NO3-1
Na+1
NO3-1
Cl-1
Cl-1
Cl-1
Cl-1
= sodium ion
= chloride ion
Al+3
Ag+1
Al+3
Na+1
Na+1
NO3-1
NO3-1
NO3-1
Al+3
NO3-1
NO3-1
Cl-1 Ag Cl-1
+1
Cl-1 Ag+1 Cl-1 Ag+1
= aluminum ion
Ag+1
= silver ion
= nitrate ion
Cl-1
= chloride ion
Learning check


Write a dissociation equation for aluminum
sulfate
Draw a beaker of dissociated aluminum
sulfate
IV. Predicting Products
A. Single Displacement reactions: an
element and a compound combine to form
a new element and compound.
* Use the activity series.
*If you don’t know the
General Equation:
charge use +2
A + YB  Y + AB (Cation)
Example:
B + AZ  Z + AB (Anion)
Na+1
Pb+2 C2H3O2-1
Sodium + Lead (II) acetate
2 Na(s) + Pb(C2H3O2)2(aq)  Pb(s) + 2 NaC2H3O2 (aq)
Single Replacement Reactions:
Activity Series



Active metal elements can replace less active
metals, active nonmetal elements can replace
less active nonmetals.
Use the Activity series (snoopy sheet) to
determine whether or not the reaction will
occur.
Driving force is the transfer of electrons.
Learning check
Try these single displacement reactions:
3. copper + silver nitrate

4. bromine + sodium chloride
Lab – Metal Activity and
Reactivity
B. Double Displacement reactions: two
compounds combine to produce two
Driving force =
different compounds - Acid-Base
liquid
and Precipitation Reactions. Driving force =
solid
* Use solubility rules.
AB + YZ  AZ + YB
General Equation:
C2H3O2-1
Pb+2
Na+1
Cl-1
Example: Lead (II) acetate + sodium chloride
Pb(C2H3O2)2 (aq) + 2 NaCl (aq)  PbCl2 (s) + 2 NaC2H3O2 (aq)
Learning check
Try these double displacement reactions:

sodium sulfate + lead (II) nitrate

sulfuric acid + potassium hydroxide
Complete Ionic and Net Ionic
Equations
Molecular Equation: shows the complete formula of all
reactants and products
Ex:
Ag+1 NO3-1
Ba+2 Cl-1
Silver nitrate + barium chloride
2 AgNO3 (aq) +
BaCl2 (aq) 
2 AgCl
(s)
+
Ba(NO3)2 (aq)
Complete Ionic Equation: represents aqueous
compounds as ions
Ex: 2 Ag+1(aq) + 2NO3-1(aq) + Ba+2(aq) + 2Cl-1(aq)  2AgCl(s) + Ba+2(aq) +
2NO3-1(aq)
Net Ionic Equation: includes only those components
directly involved in the reaction. Ions present on both
sides on the equation and do not participate directly in
the reaction are called Spectator Ions
Ex: 2Ag+1(aq) + 2Cl-1(aq)  2AgCl(s)
Ex: Write the molecular, complete ionic and
net ionic equations for the following
reaction.
Na+1
Zn+2
NO3-1
Sodium + Zinc Nitrate
Molecular – 2 Na(s) + Zn(NO3)2(aq)  Zn(s) + 2NaNO3(aq)
Complete 2 Na(s) + Zn+2(aq) + 2 NO3-1(aq)  Zn(s) + 2 Na+1(aq) + 2 NO3-1(aq)
Net – 2 Na(s) + Zn+2(aq)  Zn(s) + 2 Na+1(aq)
C. Decomposition reactions: a single compound is
broken down into more than one product. There are six
different types.
A is metal/cation, B is nonmetal/anion
1. Decomposition of a binary compound into its elements.
* Usually requires heat or electricity. Δ = heat or
electricity
General Equation:
AB  A + B
Example: Sodium Chloride
2 NaCl(aq)  2 Na(s) + Cl2(g)
2. Decomposition of a base into a metal-oxide and water.
Base = Compound
General Equation:
AOH  AO + H2O
Example:
that contains
hydroxide (OH-1)
Sodium Hydroxide
Metal-oxide = solid
2 NaOH(aq)  Na2O(s) + H2O (l)
A is metal/cation, B is nonmetal/anion
3. Decomposition of a ternary acid into a
nonmetal-oxide and water.
nonmetal-oxide = gas
General Equation: HBO  BO + H2O
Example: sulfuric acid
H2SO4(aq)  SO3(g) + H2O (l)
4. Decomposition of a metallic carbonate into a
metal-oxide and carbon dioxide.
General Equation: ACO3  AO + CO2
Example: Sodium Carbonate
Na2CO3(aq)  Na2O (s) + CO2(g)
A is metal/cation, B is nonmetal/anion
5. Decomposition of a metallic chlorate into a
metal-chloride and oxygen gas.
General Equation: AClO3  ACl + O2
Example: Sodium Chlorate
2 NaClO3(aq)  2 NaCl (aq) + 3 O2(g)
6. Decomposition of a tertiary salt into a metaloxide and a non-metal oxide
General Equation: ABO  AO+ BO
Example: Sodium Phosphate
2 Na3PO4(aq) 3Na2O(s)+ P2O5(g)
Special Situations




Whenever H2CO3, H2SO3, or NH4OH is a
product it will decompose immediately as
follows:
H2CO3  H2O + CO2
H2SO3  H2O + SO2
NH4OH  H2O + NH3
Learning check

Try these decomposition reactions:
5. Calcium chlorate
6. Phosphoric acid
7. Barium hydroxide
8. Tin (IV) carbonate
Synthesis reactions: two substances combine to
form one product. There are four different types.
A is metal/cation, B is nonmetal/anion
1.
Two elements combine to form a binary compound.
General Equation:
A + B  AB
Example: Sodium + Chlorine
2 Na(s) + Cl2(g)  2 NaCl(aq)
2. Combining a metal-oxide and water to produce a base.
General Equation:
AO + H2O  AOH
Example: Barium oxide + water
BaO(s) + H2O(l)  Ba(OH)2(aq)
A is metal/cation, B is nonmetal/anion
3.
Combining a nonmetal–oxide and water to
produce a tertiary acid.
General Equation: BO + H2O  HBO
Example: dinitrogen pentoxide + water
N2O5(g) + H2O(l) 
2 HNO3(aq)
4.
Combining a metal-oxide and a nonmetaloxide to produce a tertiary salt.
General Equation: AO + BO  ABO
Example: Barium oxide + dinitrogen pentoxide
BaO(s) + N2O5(g)  Ba(NO3)2(aq)
Learning check

Try these synthesis reactions:
9. Water + magnesium oxide
10. Water + dinitrogen trioxide
11. Bromine + sodium
Combustion reactions: Certain organic compounds
(Hydrocarbons – compound containing Carbon and Hydrogen or
Carbon, Hydrogen and Oxygen) burn to produce specific
products. There are two types.
1.
Complete combustion – combining a hydrocarbon with
excess oxygen to produce carbon dioxide and water.
* if the equation does not indicate limited oxygen assume
complete combustion
General Equation: C H + O  CO + H O
x
Example:
x
2
CH4(g) + 2 O2(g) 
2
2
CO2(g) + 2 H2O(g)
2.
Incomplete combustion - combining a hydrocarbon with
limited oxygen to produce carbon monoxide and water.
General Equation: C H + limited O  CO + H O
x
Example: 2 CH
4(g)
x
2
2
+ limited 3 O2(g)  2 CO(g) + 4 H2O(g)
Learning check

Try these combustion reactions:
12. C8H18 + oxygen
13. C2H2 + oxygen
V. Reaction Rates


Reaction rate depends on the collisions
between reacting particles.
Successful collisions occur if the particles...

collide
with each other

have the correct orientation

have enough kinetic energy
to break bonds

To speed up the rate of the reaction:

Increase
surface area (smaller particles or
dissolve in water)

Increase concentration (add more reactant)

Increase temperature (add heat source)

Add catalyze/enzyme
VI. Heat in Reactions

Exothermic reactions release heat



Heat is a product
Feels hot
Endothermic reactions absorb heat


Heat is a reactant
Feels cold