Download Oxidation - Reduction Chemistry

Oxidation - Reduction Reactions
Oxidation - Reduction
Chemistry
Oxidation - reduction (redox) reactions are chemical processes
that involve a transfer of electrons between substances
-- this can be a complete transfer to form ionic bonds
or a partial transfer to form covalent bonds
In all redox reactions:
• one substance loses electrons -- this substance is oxidized
• one substance gains electrons -- this substance is reduced
There are lots of processes in the natural world (and in the
laboratory) that involve redox reactions
• e.g., corrosion, batteries, photosynthesis/respiration, etc.
Reaction between zinc and sulfuric acid
Sulfuric acid (solution of
Zn strip
H+
Reaction between zinc and sulfuric acid
2-
and SO4 ions)
Overall reaction
H2 bubbles
Zn(s) + H2SO4(aq)
ZnSO4(aq) + H2(g)
Zinc loses electrons
Overall ionic reaction
• zinc is oxidized
Zn(s)
Zn2+(aq) + 2e-
-- all dissolved ions are explicitly shown
Zn(s) + 2 H+(aq) + SO42–(aq)
Zn2+(aq) + SO42–(aq) + H2(g)
Hydrogen gains electrons
Net ionic reaction
• hydrogen is reduced
2 H+(aq) + 2eElectrons are transferred from zinc to hydrogen
H2(g)
-- includes only substances that undergo change
-- ions that are present but do not react (spectator ions) are not shown
Zn(s) + 2 H+(aq)
Zn2+(aq) + H2(g)
Oxidation - Reduction Reactions
Oxidation - reduction (redox) reactions are chemical processes
that involve a transfer of electrons between substances
• Oxidation occurs when a substance loses electrons
• Reduction occurs when a substance gains electrons
In a redox reaction, oxidation and reduction occur simultaneously
-- one cannot occur in the absence of the other
Reaction between Cu and AgNO3
Cu(s)
Ag+(aq)
Ag(s)
initial
final
Cu2+(aq)
NO3–(aq)
NO3–(aq)
Oxidation of Cu: Cu(s) ! Cu2+(aq) + 2eReduction of Ag+: 2 Ag+(aq) + 2e- ! 2 Ag(s)
Electrons are transferred from Cu atoms to Ag+ ions in solution
Cu loses electrons (oxidation)
initial
Cu(s)
Ag+(aq)
Cu(s) + 2 Ag+(aq)
final
Ag(s)
Cu2+(aq)
Cu2+(aq) + 2 Ag(s)
Ag+ gains electrons (reduction)
Voltaic cells
Overall Reaction: Cu(s) + 2 AgNO3(aq) ! 2 Ag(s) + Cu(NO3)2(aq)
Overall Ionic Equation:
Cu(s) + 2 Ag+(aq) + 2 NO3–(aq) ! 2 Ag(s) + Cu2+(aq) + 2 NO3–(aq)
Net Ionic Equation: Cu(s) + 2 Ag+(aq) ! 2 Ag(s) + Cu2+(aq)
Zinc-copper voltaic cell
A voltaic cell is a device that produces an electric current from
a spontaneous redox reaction
• the
oxidation reaction and the reduction reaction are
physically separated and connected with a wire
• electrons transferred during the redox reaction must pass
through the wire, producing an electric current
• the electric current is used to perform work
- e.g., lighting a bulb, running an electric motor, etc.
Chemical potential energy (energy stored in chemical bonds)
is converted to electricity that is used to perform work
Oxidation of Zn: Zn(s) ! Zn2+(aq) + 2eReduction of Cu2+: Cu2+(aq) + 2e- ! Cu(s)
Net Ionic: Zn(s) + Cu2+(aq) ! Cu(s) + Zn2+(aq)
Electrons are transferred from Zn atoms to Cu2+ ions in solution
Dry cell batteries
Oxidation number
A dry cell battery is a small, efficient voltaic cell that
contains a non-liquid electrolyte
metal cap (+)
The oxidation number of an atom is an integer value that
represents the number of electrons gained, lost, or
unequally shared by that atom
Alkaline-type
dry cell battery
carbon rod
(positive electrode)
Electrolyte is
NaOH or KOH
zinc case
(negative electrode)
Electrolyte is
source of OHand H2O for
redox reactions
manganese (IV) oxide
moist paste of
NaOH or KOH
(electrolyte)
metal bottom (–)
• an oxidation number of zero indicates that the atom has
the same number of electrons assigned to it as there are
in the free, neutral atom
• a positive oxidation number indicates that the atom has
fewer electrons assigned to it than in the neutral atom
• a negative oxidation number indicates that the atom has
more electrons assigned to in than in the neutral atom
Oxidation of Zn: Zn(s) + 2 OH-(aq) ! ZnO(s) + H2O(l) + 2eReduction of Mn4+: 2 MnO2(s) + H2O(l) + 2e- ! Mn2O3(s) + 2 OH-(aq)
Oxidation number
An element in its free state (uncombined with other elements)
has an oxidation number of zero
•• Ba
barium (Ba)
•• ••
••Cl •• Cl ••
•• ••
chlorine (Cl2)
The oxidation number of an element that has gained or
lost electrons to form an ion is that same as its positive or
negative charge
+2
Ba2+
barium ion
(oxidation number: +2)
Clchloride ion
(oxidation number: –1)
-1
BaCl2
In an ionic compound,
the ions retain their
oxidation number
Oxidation
numbers of
common
monoatomic
ions
Oxidation number
In covalent compounds (shared electrons), oxidation
numbers are assigned by an arbitrary system based on
relative electronegativities of the atoms
H •• H
hydrogen (H2)
•• ••
•• Cl •• Cl ••
•• ••
chlorine (Cl2)
In diatomic molecules containing only one element (nonpolar covalent bonding), the bonding pair of electrons is
shared equally between the atoms (" electronegativity = 0)
• each atom is assigned an oxidation number of zero
Oxidation number
In covalent compounds containing different elements, the
bonding atoms are shared unequally between the atoms
-- the higher the electronegativity of the atom, the greater
its affinity for the shared electrons
In these types of compounds, oxidation numbers are
determined by assigning both bonding electrons to the
most electronegative atom
Example: Water
+1
Electronegativity
Oxygen: 3.5
Hydrogen: 2.1
–2
••
H •• O •• H
••
both bonding electrons
assigned to oxygen
+1
both bonding electrons
assigned to oxygen
Many elements have multiple
oxidation numbers
Many elements have multiple
oxidation numbers
It depends on the types of compounds they form
It depends on the types of compounds they form
N oxidation
number
N2
N2O
NO
N2O3
NO2
N2O5
NO3–
0
+1
+2
+3
+4
+5
+5
Note: The oxidation number of oxygen is –2 in all of these compounds
Elemental copper (Cu)
Copper (II) sulfate
(CuSO4)
Cu oxidation state: 0
Cu oxidation state: +2
Rules for assigning oxidation numbers
1. All elements in their free state (uncombined with other elements) have
an oxidation number of zero (e.g., Na, Cu, Mg, H2, O2, Cl2, etc.)
2. H is +1, except in metal hydrides, where it is -1 (e.g., NaH, CaH2)
3. O is -2, except in peroxides, where it is -1, and in OF2, where it is +2
4. The metallic element in an ionic compound has a positive oxidation
number
5. In covalent compounds, the most electronegative element is assigned
a negative oxidation number
6. The sum of the oxidation numbers of the elements in a neutral
compound is zero
Rules for determining the oxidation number
of an element within a compound
Step 1: Write the oxidation number of each known atom
below the atom in the formula
Step 2: Multiply each oxidation number by the number of
atoms of that element in the compound
Step 3: Assign oxidation numbers for the other atoms in
the compound in order to make the sum of the
oxidation numbers equal to zero
7. The sum of the oxidation numbers of the elements in a polyatomic ion
is equal to the charge of the ion
Example: Determine the oxidation number of
carbon in carbon dioxide
CO2
-2
2(-2) = -4
4 + C + (-4) = 0 + 4
Example: Determine the oxidation number of
sulfur in sulfuric acid
H2SO4
+1
2(+1) = +2
-2
4(-2) = -8
(–2) + 8 +2 + S + (-8) = 0 (–2) + 8
C = +4
S = +6
(oxidation number for carbon)
(oxidation number for sulfur)
Step 1: Write the oxidation number of each known atom below the atom in
the formula
Step 1: Write the oxidation number of each known atom below the atom in
the formula
Step 2: Multiply each oxidation number by the number of atoms of that
element in the compound
Step 2: Multiply each oxidation number by the number of atoms of that
element in the compound
Step 3: Assign oxidation numbers for the other atoms in the compound in
order to make the sum of the oxidation numbers equal to zero
Step 3: Assign oxidation numbers for the other atoms in the compound in
order to make the sum of the oxidation numbers equal to zero
Example: Determine the oxidation number of
chromium in Cr2O72-
Example: Determine the oxidation number of
potassium and nitrogen in KNO3
Cr2O72-
KNO3
-2
K+
NO3–
Recognize that KNO3 is an ionic compound between K+ and NO3-
7(-2) = -14
2Cr + (-14) = -2 (the charge on the ion)
Cr = +6 (oxidation number for chromium)
Step 1: Write the oxidation number of each known atom below the atom in
the formula
The oxidation number of potassium in K+ is +1 (the charge on the ion)
For nitrogen:
NO3–
-2
3(-2) = -6
Step 2: Multiply each oxidation number by the number of atoms of that
element in the compound
N + (-6) = -1 (the charge on the ion)
Step 3: Assign oxidation numbers for the other atoms in the compound in
order to make the sum of the oxidation numbers equal to zero
N = +5 (oxidation number for nitrogen)
Oxidation - Reduction Reactions
Oxidation - Reduction Reactions
Oxidation - reduction (redox) reactions are chemical
processes that involve a transfer of electrons between
substances
Oxidation - reduction (redox) reactions are chemical
processes that involve a transfer of electrons between
substances
-- this can be a complete transfer to form ionic bonds
or a partial transfer to form covalent bonds
-- this can be a complete transfer to form ionic bonds
or a partial transfer to form covalent bonds
L ose
E lectrons
O xidation
G ain
E lectrons
R eduction
In redox reactions, the oxidation numbers of the
elements involved in the reaction change
• oxidation of an element (loss of electrons) results in
an increase in its oxidation number
• reduction of an element (gain of electrons) results in a
decrease in its oxidation number
Reaction can be rewritten to
emphasize electron transfer
Reaction between zinc and sulfuric acid
Zn(s) + H2SO4(aq)
ZnSO4(aq) + H2(g)
Zn + 2 H+ + SO42-
Zn2+ + SO42- + H2
0
+2
0
+2
+1
0
+1
0
•zinc loses electrons
•zinc loses electrons
•the oxidation number of Zn increases
•zinc is oxidized
•the oxidation number of Zn increases
•zinc is oxidized
•hydrogen gains electrons
•hydrogen gains electrons
•the oxidation number of H decreases
•hydrogen is reduced
•the oxidation number of H decreases
•hydrogen is reduced
Electrons are transferred from zinc to hydrogen
Electrons are transferred from zinc to hydrogen
Reaction between zinc and sulfuric acid
Oxidizing and reducing agents
Oxidizing agent: The reactant that causes another
substance to be oxidized
– i.e., the reactant that causes an increase in
the oxidation state of another substance
The oxidizing agent is reduced in a redox reaction
Reducing agent: The reactant that causes another
substance to be reduced
– i.e., the reactant that causes a decrease in
the oxidation state of another substance
The reducing agent is oxidized in a redox reaction
Zn(s) + H2SO4(aq)
ZnSO4(aq) + H2(g)
0
+2
+1
0
•
•
•
•
zinc loses electrons
the oxidation number of Zn increases
zinc is oxidized
zinc is the reducing agent
(it causes hydrogen to be reduced)
•
•
•
•
hydrogen gains electrons
the oxidation number of H decreases
hydrogen is reduced
sulfuric acid is the oxidizing agent
(it causes zinc to be oxidized)
Example: Is the following a redox reaction?
Example: Is the following a redox reaction?
Neutralization reaction between hydrochloric acid and potassium hydroxide
Thermite reaction
HCl (aq) + KOH (aq)
+1
-1
Element
+1
H2O (l) + KCl (aq)
-2 +1
+1
Oxidation number Oxidation number
before reaction
after reaction
H
+1
+1
O
–2
–2
K
+1
+1
Cl
–1
–1
-2
+1
-1
All oxidation
numbers unchanged
No redox reactions
occurred
Homework assignment
Chapter 6 Problems:
6.76, 6.77, 6.78, 6.79, 6.84, 6.87, 6.88, 6.89, 6.90,
6.91, 6.92
2 Al (s) + Fe2O3 (s)
0
Element
+3
Al2O3 (l) + 2 Fe (l)
-2
+3
-2
Oxidation number Oxidation number
before reaction
after reaction
Al
0
+3
Fe
+3
0
O
–2
–2
0
Which element
is oxidized?
Al
Which element
is reduced?
Fe