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Chapter 3
Chemical Reactions
Jeffrey Mack
California State University,
Sacramento
Chemical Reactions
Reactants: Zn + I2
Product: ZnI2
Chemical Reactions
Evidence of a chemical reaction:
• Gas Evolution
• Temperature Change
• Color Change
• Precipitation (insoluble species forms)
In general, a reaction involves a rearrangement
or change in oxidation state of atoms from
reactants to products.
Chemical Equations
Chemical Equations show:
• the reactants and products in a reaction.
• the relative amounts in a reaction.
Example:
4 Al(s) + 3 O2(g)
2 Al2O3(s)
• The numbers in the front are called
stoichiometric coefficients
• The letters (s), (g), (l) and (aq) are the
physical states of compounds.
Reaction of Phosphorus with Cl2
Notice the stoichiometric coefficients and the physical states of
the reactants and products.
Reaction of Iron with Cl2
Notice the stoichiometric coefficients and the physical states of
the reactants and products.
Chemical Equations
4 Al(s) + 3 O2(g)
2 Al2O3(s)
This equation states that:
4 Al atoms + 3 O2 molecules
react to form 2 formula units
of Al2O3
or...
4 moles of Al + 3 moles of
O2 react to form 2 moles of
Al2O3
Chemical Equations
Law of the
Conservation of Matter
• Because the same
number of atoms are
present in a reaction at
the beginning and at
the end, the amount of
matter in a system
does not change.
2HgO(s)
2 Hg(l) + O2(g)
Chemical Equations
• Since matter is
conserved in a chemical
reaction, chemical
equations must be
balanced for mass!
• In other words, there
must be same number of
atoms of the each kind
on both sides of the
equatoin.
Lavoisier, 1788
Balancing Chemical Reactions
Steps in balancing a chemical reaction using coefficients:
1. Write the equation using the formulas of the reactants
and products. Include the physical states (s, l, g, aq
etc…)
2. Balance the compound with the most elements in the
formula first using integers as coefficients.
3. Balance elements on their own last.
4. Check to see that the sum of each individual elements
are equal on each side of the equation.
5. If the coefficients can be simplified by dividing though
with a whole number, do so.
Balancing Chemical Equations:
Example
balance last
C2H6 +
2 C’s & 6 H’s
O2
CO2
+
2 O’s
1 C & 2 O’s
H2O
2 H’s & 1 O
Balancing Chemical Equations:
Example
balance last
C2H6 +
2 C’s & 6 H’s
O2
CO2
+
2 O’s
1 C & 2 O’s
H2O
2 H’s & 1 O
balance H first
___C2H6
+
O2
This side will always have
an even # of O-atoms
CO2
+
3 H2O
___
This side has an odd # of
O-atoms
Balancing Chemical Equations:
Example
balance last
C2H6 +
2 C’s & 6 H’s
O2
CO2
+
2 O’s
1 C & 2 O’s
H2O
2 H’s & 1 O
balance H first
2 2H6
___C
+
O2
CO2
+
3 H2O
___
Balancing Chemical Equations:
Example
balance last
C2H6 +
2 C’s & 6 H’s
O2
CO2
+
2 O’s
1 C & 2 O’s
H2O
2 H’s & 1 O
balance H first
2 2H6
___C
+
O2
CO2
3 H2O
___
+
balance C next
2C2H6 +
O2
4 CO2
___
+
6H2O
Balancing Chemical Equations:
Example
balance last
C2H6 +
2 C’s & 6 H’s
O2
CO2
+
2 O’s
1 C & 2 O’s
H2O
2 H’s & 1 O
balance H first
2 2H6
___C
+
O2
CO2
3 H2O
___
+
balance C next
2C2H6 +
O2
4 CO2
___
+
6H2O
balance O
2C2H6 +
7 O2
____
4CO2 +
6H2O
Balancing Chemical Equations:
Example
balance last
C2H6 +
2 C’s & 6 H’s
O2
CO2
+
2 O’s
1 C & 2 O’s
H2O
2 H’s & 1 O
balance H first
2 2H6
___C
+
O2
CO2
3 H2O
___
+
balance C next
2C2H6 +
O2
4 CO2
___
+
6H2O
balance O
2C2H6 +
7 O2
____
4 C’s 12 H’s 14 O’s
4CO2 +
6H2O
4 C’s 12 H’s 14 O’s
Balancing Equations
___ Al(s) + ___ Br2(l)
___ Al2Br6(s)
Balancing Equations: Practice
___C3H8(g) + ___ O2(g)
___ CO2(g) + _____ H2O(g)
___B4H10(g) + ___ O2(g)
___ B2O3(g) + ___ H2O(g)
Balancing Equations: Practice
• Solid magnesium hydroxide reacts with hydrochloric
acid to form aqueous magnesium chloride and
water.
• Write the balanced chemical equation for this
reaction.
Balancing Equations: Practice
• Solid magnesium hydroxide reacts with hydrochloric
acid to form aqueous magnesium chloride and
water.
• Write the balanced chemical equation for this
reaction.
_ Mg(OH)2(s) + _ HCl(aq)
Balancing Equations: Practice
• Solid magnesium hydroxide reacts with hydrochloric
acid to form aqueous magnesium chloride and
water.
• Write the balanced chemical equation for this
reaction.
_ Mg(OH)2(s) + _ HCl(aq)
_ MgCl2(aq) + _ H2O(l)
Balancing Equations: Practice
• Solid magnesium hydroxide reacts with hydrochloric
acid to form aqueous magnesium chloride and
water.
• Write the balanced chemical equation for this
reaction.
_ Mg(OH)2(s) + _ HCl(aq)
_ MgCl2(aq) + _ H2O(l)
Balance with a coefficient of ―2‖ in front of both HCl
and water.
Balancing Equations: Practice
• Solid magnesium hydroxide reacts with hydrochloric
acid to form aqueous magnesium chloride and
water.
• Write the balanced chemical equation for this
reaction.
_ Mg(OH)2(s) + _ HCl(aq)
_ MgCl2(aq) + _ H2O(l)
Balance with a coefficient of ―2‖ in front of both HCl
and water.
Mg, Cl, O and H are now balanced.
Chemical Equations: Review
• What Scientific Principles are used in the process of
balancing chemical equations?
• What symbols are used in chemical equations:
gasses:
_____
liquids:
_____
solids:
_____
aqueous species in solution:
_____
• What is the difference between P4 and 4P in an
eq.?
• In balancing a chemical equation, why are the
reactant and product subscripts not changed?
Chemical Equilibrium
When writing chemical reactions one starts with:
Reactants
products
N2(g) + 3H2(g)
2NH3(g)
Some reactions can also run in reverse:
2NH3(g)
N2(g) + 3H2(g)
Under these conditions, the reaction can be written:
3H2 (g) N2 (g)
2NH3 (g)
Double arrows indicate ―Equilibrium‖.
Chemical Equilibrium
Once equilibrium is achieved, reaction continues, but there
is no net change in amounts of products or reactants.
Classifying Compounds
• Salts (ionic compounds): Composed of a
metal and non metal element(s).
• Acids: Arrhenius definition
Produce H+(aq) in water
Examples: HCl, HNO3, HC2H3O2
• Bases: Arrhenius definition
Produce OH (aq) in water
Examples: NaOH, Ba(OH)2, NH3
Classifying Compounds
• Molecular Compounds:
• Covalently bonded atoms, not acids, bases or
salts.
• Compounds like alcohols (C2H5OH) or table
sugar (C6H12O6)
• These never break up into ions.
Classifying Compounds
• Classify the following as ionic, molecular,
acid or base.
Compound
Na2SO4
Ba(OH)2
H3PO4
CH4
P2O5
NH3
HCN
Type
Classifying Compounds
• Classify the following as ionic, molecular,
acid or base.
Compound
Na2SO4
Ba(OH)2
H3PO4
CH4
P2O5
NH3
HCN
Type
ionic
base
acid
molecular
molecular
base
acid
Reactions in Aqueous Solutions
Aqueous Solutions: Water as the solvent
Solution =
solute
That which is dissolved
(lesser amount)
+
solvent
That which is dissolves
(greater amount)
There are three types of aqueous solutions:
Those with Strong Electrolytes
Those with Weak Electrolytes
& those with non-Electrolytes
Reactions in Aqueous Solutions
Many reactions involve ionic compounds,
especially reactions in water — aqueous
solutions.
KMnO4 in water
K+(aq) + MnO4-(aq)
Ionic Compounds (CuCl2) in Water
Strong Electrolyte
When ions are present in water,
the solutions conduct
electricity!
Ions in solution are called
ELECTROLYTES
Examples of Strong Electrolytes:
HCl (aq), CuCl2(aq) and NaCl
(aq) are strong electrolytes.
These dissociate completely (or
nearly so) into ions.
Strong Electrolytes conduct
electricity well.
Strong Electrolytes
HCl(aq), CuCl2(aq) and NaCl(aq) are strong
electrolytes.
These dissociate completely (or nearly so) into
ions.
Weak Electrolytes
Acetic acid ionizes only to a small
extent, it is a weak electrolyte.
CH3CO2H(aq)
CH3CO2 (aq)
H (aq)
Weak electrolytes exist in solution
under equilibrium conditions.
The small concentration of ions
conducts electricity poorly.
Weak electrolytes exit primarily in
their molecular form in water.
Weak Electrolytes
Weak electrolytic solutions are characterized by
equilibrium conditions in solution:
When acetic acid dissociates, it only partially
ionizes.
HC2H3O2 (aq)
+
H (aq) + C2H3O2 (aq)
95%
5%
The majority species in solution is acetic acid in its
molecular form.
When writing a weak electrolyte in solution, one
NEVER breaks it up into the corresponding ions!
HC2H3O2 (aq)
+
H (aq) + C2H3O2 (aq)
×
Weak Electrolytes
Acetic acid ionizes only to a small extent, so it
is a weak electrolyte.
CH3CO2H(aq)
CH3CO2-(aq) + H+(aq)
Non-Electrolytes
Some compounds dissolve in
water but do not conduct
electricity.
They are non-electrolytes.
Examples include:
• sugar
• ethanol
• ethylene glycol
Non-electrolytes do not
dissociate into ions!
Species in Solution: Electrolytes
Strong electrolytes: Characterized by ions only (cations &
anions) in solution (water).
Conduct electricity well
Weak electrolytes:
Characterized by ions (cations & anions)
& molecules in solution.
Conduct electricity poorly
Non-electrolytes:
Characterized by molecules in solution.
Do not conduct electricity
Solutes in Aqueous Solutions
Solubility Rules
How do we know if a compound will be soluble in
water?
For molecular compounds, the molecule must be
polar.
We will discuss polarity later, for now I will tell you
whether or not a molecular compound is polar…
For ionic compounds, the compound solubility is
governed by a set of SOLUBILITY RULES!
You must learn the basic rules on your own!!!
Water Solubility of Ionic
Compounds
If one ion from the ―Soluble
Compound‖ list is present in a
compound, then the compound is
water soluble.
Types of Reactions in a Solution
Precipitation Reactions: A reaction where an
insoluble solid (precipitate) forms and drops out
of the solution.
Acid–base Neutralization: A reaction in which an
acid reacts with a base to yield water plus a salt.
Gas forming Reactions: A reaction where an
insoluble gas is formed.
Reduction and Oxidation Reactions (RedOx): A
reaction where electrons are transferred from
one reactant to another.
EXCHANGE: Precipitation Reactions
EXCHANGE
Gas-Forming
Reactions
REACTIONS
REDOX REACTIONS
EXCHANGE
Acid-Base
Reactions
Chemical Reactions in Water
EXCHANGE REACTIONS
Pb(NO3) 2(aq) + 2 KI(aq)
PbI2(s) + 2 KNO3 (aq)
The anions exchange
places between cations.
A precipitate forms if one of
the products in insoluble.
Precipitation Reactions
The ―driving force‖ is the formation of
an insoluble solid called a precipitate.
Pb(NO3)2(aq) + 2 KI(aq)
2 KNO3(aq) + PbI2(s)
BaCl2(aq) + Na2SO4(aq)
BaSO4(s) + 2 NaCl(aq)
Precipitates are determined from the
solubility rules.
Precipitation Reactions
Which species is the precipitate?
Pb(NO3)2(?) + 2KI(?)
2KNO3(?) + PbI2(?)
Precipitation reactions
Which species is the precipitate?
Pb(NO3)2(?) + 2KI(?)
2KNO3(?) + PbI2(?)
From the solubility rules:
All nitrate salts are soluble, therefore:
Pb(NO3)2(aq) + 2KI(?)
2KNO3(aq) + PbI2(?)
Precipitation Reactions
Which species is the precipitate?
Pb(NO3)2(?) + 2KI(?)
2KNO3(?) + PbI2(?)
From the solubility rules:
All nitrate salts are soluble, therefore:
Pb(NO3)2(aq) + 2KI(?)
2KNO3(aq) + PbI2(?)
All potassium salts are soluble, therefore:
Pb(NO3)2(aq) + 2KI(aq)
2KNO3(aq) + PbI2(?)
Precipitation Reactions
Which species is the precipitate?
Pb(NO3)2(?) + 2KI(?)
2KNO3(?) + PbI2(?)
From the solubility rules:
All nitrate salts are soluble, therefore:
Pb(NO3)2(aq) + 2KI(?)
2KNO3(aq) + PbI2(?)
All potassium salts are soluble, therefore:
Pb(NO3)2(aq) + 2KI(aq)
By the solubility rules:
Pb(NO3)2(aq) + 2KI(aq)
2KNO3(aq) + PbI2(?)
PbI2 is the ppt.
2KNO3(aq) + PbI2(s)
Net Ionic Equations
Molecular Equation: all species listed as formula units or in
molecular form.
reactants
products
• Note all states of each reactant or product by: (s), (l), (g) or
(aq)
Ionic Equation: All soluble (aq) species present are listed as
ions.
• Leave all (s), (l) or (g) species as is. They do not dissociate
into ions
Net Ionic Equation:
• From the ionic equation, cancel out any species that appear
on either side of the equation.
• These are known as the ―spectator ions‖ and they are
never part of a net ionic equation!
Writing Net Ionic Equations
Molecular Equation:
Pb(NO3)2(aq) + 2KI(aq)
2KNO3(aq) + PbI2(s)
Writing Net Ionic Equations
Molecular Equation:
Pb(NO3)2(aq) + 2KI(aq)
2KNO3(aq) + PbI2(s)
Total Ionic Equation:
Pb2+ (aq) + 2NO3– (aq) + 2K+(aq) + 2I–(aq)
2K+(aq) + 2NO3– (aq) +
PbI2(s)
Writing Net Ionic Equations
Molecular Equation:
Pb(NO3)2(aq) + 2KI(aq)
2KNO3(aq) + PbI2(s)
Total Ionic Equation:
Pb2+ (aq) + 2NO3– (aq) + 2K+(aq) + 2I–(aq)
Never break up
any (s), (l) or (g)
or molecular
(aq) species!
2K+(aq) + 2NO3– (aq) +
PbI2(s)
Writing Net Ionic Equations
Molecular Equation:
Pb(NO3)2(aq) + 2KI(aq)
2KNO3(aq) + PbI2(s)
Total Ionic Equation:
Pb2+ (aq) + 2NO3– (aq) + 2K+(aq) + 2I–(aq)
Never break up
any (s), (l) or (g)
or molecular
(aq) species!
2K+(aq) + 2NO3– (aq) +
PbI2(s)
Cancel out the spectator ions to yield the net ionic equation:
Pb2+ (aq) + 2I–(aq)
PbI2(s)
Acids & Bases
Arrhenius Definition:
• An acid is any substance that increases the
H+(aq) concentration in an aqueous solution.
HX(aq)
H+(aq) + X–(aq)
• A base is any substance that increases the
OH–(aq) concentration in an aqueous
solution.
MOH(aq)
M+(aq) + OH–(aq)
Acids and Bases
Brönsted-Lowry:
• An acid is any substance that donates H+(aq)
to another species in an aqueous solution.
HX(aq) + H2O(l) H3O+(aq) + X–(aq)
H3O+(aq) = H+(aq)
• A base is any substance that accepts an
H+(aq) in an aqueous solution.
H+(aq) + NH3(aq) NH4+(aq)
Acids
Strong Acids
Examples: Strong acids are almost completely ionized in
water. (strong electrolytes)
HX (aq) (X = Cl, Br & I)
hydro ___ ic acid
HNO3 (aq)
nitric acid
HClO4 (aq)
perchloric acid
H2SO4 (aq)*
sulfuric acid
* Only the 1st H is strong, sulfuric acid dissociates via:
H2SO4 (aq)
H+ (aq) + HSO4– (aq)
Acids
An acid: H3O+ in water
Weak Acids
Examples: Weak Acids are incompletely ionized in water.
(weak electrolytes) Weak acids are governed by
dynamic equilibrium.
HC2H3O2 (aq)
acetic acid (vinegar)
nitrous acid
HNO2 (aq)
hydrosulfuric acid
H2S (aq)
hydrogen sulfate ion
HSO4–(aq)
Weak acids are always written in their molecular form.
See you text and home work for more examples.
Strong Bases
Bases: A base is a substance that produces OH– (aq) ions in
water by dissociation in water:
NaOH(s)
H2O( )
Na (aq) ΟΗ aq
Strong bases are almost completely ionized in aqueous
solution. (Strong electrolytes)
Examples: Hydroxides of Group 1 (MOH(aq) where M = Li,
Na, K ect…) and Ca, Sr, Ba.*
*Ca(OH)2, Sr(OH)2 & Ba(OH)2 are slightly soluble, but that
which dissolves is present as ions only.
Bases
Base: OH- in water
NaOH(aq)
NaOH is a
strong base
Na+(aq) + OH-(aq)
Weak Bases
Weak Bases:
NH3 acts as a base by reacting with water:
NH3(aq) + H2O(l)
NH4+(aq) + OH –(aq)
Ammonia can also accept H+ from an acid:
NH3(aq) + H+(aq)
NH4+(aq)
Ammonia, NH3
Reactions of Acids & Bases:
Acid-Base Neutralization
Acid + Base
Salt + Water (usually)
HA (aq) + MOH(aq)
MA(aq) + HOH(l)
Strong acid - Strong base neutralization: HBr(aq)/KOH(aq)
Molecular Equation:
HBr(aq) + KOH(aq)
KBr (aq) + H2O(l)
Total Ionic Equation:
/
H+ (aq) + Br– (aq)+ K+(aq) + OH– (aq)
/
K+(aq) + Br– (aq)
+ H2O(l)
/
/
Net Ionic equation:
H+ (aq) + OH– (aq)
H2O (l)
Acid-Base Reactions
• The ―driving force‖ is the formation of water.
NaOH(aq) + HCl(aq)
NaCl(aq) + H2O(liq)
• Net ionic equation
OH-(aq) + H3O+(aq)
2 H2O(l)
• This applies to ALL reactions
of STRONG acids and bases.
Reactions of Acids & Bases:
Acid-Base Neutralization
Reactions of weak acids and strong bases:
Molecular Equation:
HC2H3O2(aq) + NaOH(aq)
NaC2H3O2(aq) + H2O(l)
Total Ionic Equation:
/
HC2H3O2(aq) + Na+(aq) + OH–(aq)
/
Na+(aq) + C2H3O2–(aq) + H2O(l)
Leave in
molecular
form
Net Ionic: HC2H3O2(aq) + OH–(aq)
C2H3O2–(aq) + H2O(l)
Non-Metal Acids
Nonmetal oxides can form acids in
aqueous solutions:
Examples:
CO2(aq) + H2O(s)
H2CO3(aq)
SO3(aq) + H2O(s)
H2SO4(aq)
Both gases come from the burning
of fossil fuels.
Bases
Metal oxides form bases
in aqueous solution
CaO(s) + H2O(l)
Ca(OH)2(aq)
CaO in water. Indicator
shows solution is basic.
Gas-Forming Reactions
Gas-Forming Reactions
Metal carbonate salts react with acids to the corresponding
metal salt, water and carbon dioxide gas.
2HCl(aq) + CaCO3(s)
CaCl2(aq) + H2CO3(aq)
decomposes
H2O(l) + CO2(g)
Similarly:
HCl(aq) + NaHCO3(s)
acid
base
NaCl(aq) + H2O(l) + CO2(g)
salt
water
Neutralization!!!
Gas-Forming Reactions
Group I metals: Na, K, Cs etc.. react vigorously
with water
2K(s) + 2H2O(l)
2KOH(aq)+ H2(g)
Metals & acid:
Some metals react vigorously with acid solutions:
Zn(s) + 2H+(aq)
Zn2+(aq) + H2(g)
Gas-Forming Reactions
CaCO3(s) + H2SO4(aq)
2 CaSO4(s) + H2CO3(aq)
Carbonic acid is unstable and forms CO2 & H2O
H2CO3(aq)
CO2 + water
(The antacid tablet contains citric acid + NaHCO3)
Oxidation-Reduction Reactions
Thermite reaction:
Fe2O3(s) + 2Al(s)
2Fe(s) + Al2O3(s)
Oxidation-Reduction Reactions
REDOX = reduction & oxidation
O2(g) + 2 H2(g)
2 H2O(l)
Oxidation-Reduction Reactions
• Oxidation involves a reactant atom or compound losing
electrons.
• Reduction involves a reactant atom or substance gaining
electrons.
• Neither process can occur alone… that is, there must be
an exchange of electrons in the process.
• The substance that is oxidized is the reducing agent
• The substance that is reduced is the oxidizing agent
oxidized
reduced
Mg(s) + 2H+(aq)
reducing
agent
oxidizing
agent
Mg2+(aq) + H2(g)
Oxidation Numbers
• Chemists use oxidation numbers to account for the
transfer of electrons in a RedOx reaction.
• Oxidation numbers are the actual or apparent
charge on atom when alone or combined in a
compound.
1. The atoms of pure elements always have an
oxidation number of zero.
Examples:
Mg(s)
Hg(l)
I2(s)
O2(g)
All have an
oxidation number
of zero (0)
Oxidation Numbers
2. If an atom is charged, then the charge is the
oxidation numbers .
Examples:
Ion
Oxidation Number
Mg2+(aq)
+2
Cl (aq)
1
Sn4+(s)
+4
Hg22 (aq)
+2/2 = +1 for each Hg atom
Oxidation Numbers
3. In a compound, fluorine always has an oxidation
numbers of 1.
4. Oxygen most often has an oxidation number of 2.
» *When combined with fluorine, oxygen has a positive O.N.
» *In peroxide, the O.N. is 1.
5. In compounds, Cl, Br & I are 1 (Except with F and
O present)
6. In compounds, H is +1, except as a hydride
(H : 1)
Oxidation Numbers
Examples:
compound
Oxidation Numbers
HF(g)
H = +1
F= 1
H2O(l)
H = +1
O= 2
OF2(g)
O = +2
F= 1
Na2O2(s)
Na = +1
O= 1
HCl(g)
H = +1
Cl = 1
NaH(l)
Na = +1
H= 1
Oxidation Numbers
Most common oxidation numbers:
Oxidation Numbers
7. For neutral compounds, the sum of the oxidation numbers
equals zero.
For a poly atomic ion, the sum equals the charge.
Examples:
+2 + 2 × (−1) = 0
MgCl2
3 + 4 × (+1) = +1
NH4
Oxidation Numbers
Determine the oxidation number of iron in the
following compound:
? + 3 ( 1) = 0
Fe(OH)3
Iron must have an oxidation number of +3!
Recognizing a Redox Reaction
In a RedOx reaction, the species oxidized and the
species reduced are identified by the changes in
oxidation numbers :
Oxidation numbers:
+1
0
2Ag+ (aq) + Cu(s) ® 2Ag(s) + Cu2+ (aq)
0
+2
Oxidation numbers:
Since silver goes from +1 to zero, it is reduced.
Since copper goes from zero to +2, it is oxidized.
The reaction is balanced for both mass and charge.
Practice:
Identify the species that is Oxidized and
Reduced by assigning oxidation numbers in the
following reaction.
3CH4 (g) Cr2O72 (aq) 8H (aq)
3
3CH3OH(l) 2Cr (aq) 4H2O(l)
Answer:
Practice:
Identify the species that is Oxidized and
Reduced by assigning oxidation numbers in the
following reaction.
3CH4 (g) Cr2O72 (aq) 8H (aq)
3
3CH3OH(l) 2Cr (aq) 4H2O(l)
Answer:
• The carbon in methane (CH4) is oxidized ( 4 to 2)
Practice:
Identify the species that is Oxidized and
Reduced by assigning oxidation numbers in the
following reaction.
3CH4 (g) Cr2O72 (aq) 8H (aq)
3
3CH3OH(l) 2Cr (aq) 4H2O(l)
Answer:
• The carbon in methane (CH4) is oxidized ( 4 to 2)
• Chromium in dichromate is reduced (+6 to +3)
Redox Reactions
Oxidation-Reduction Reactions
• Iron gains 3 electrons
(+3 to 0) oxidation
number change. It is
Reduced.
• Carbon loses 2
electrons (+2 to +4) it
is Oxidized.
Redox Reactions
REDOX = reduction & oxidation
Corrosion of aluminum
2 Al(s) + 3 Cu2+(aq)
2 Al3+(aq) + 3 Cu(s)
Redox Reactions
Cu(s) + 2 Ag+(aq)
Cu2+(aq) + 2 Ag(s)
In all reactions if a
species is oxidized then
another species must
also been reduced
Redox Reactions
Cu(s) + 2 Ag+(aq)
Cu2+(aq) + 2 Ag(s)
Electron Transfer in a Redox
Reaction
e
e
2Ag+(aq) + Cu(s)
•
•
•
•
Cu2+(aq) + 2Ag(s)
Two electrons leave copper.
The silver ions accept them.
The copper metal is oxidized to copper (II) ion.
The silver ion is reduced to solid silver metal.
Redox Reactions in Our World
Batteries
Corrosion
Fuels
Manufacturing metals
Examples of Redox Reactions
Metal + halogen
2 Al + 3 Br2 Al2Br6
Examples of Redox Reactions
Nonmetal (P) + Oxygen
Metal (Mg) + Oxygen
P4O10
MgO
Examples of Redox Reactions
Metal + acid
Mg + HCl
Mg = reducing agent
H+ = oxidizing agent
Metal + acid
Cu + HNO3
Cu = reducing agent
HNO3 = oxidizing agent
Reviewing What You’ve Learned
• You have the following items available to you:
Deionized water, pH paper, test tubes various
metal nitrate salts, common acid and base
solutions.
• Suggest a simple test or set of tests for
identifying the unknown substances. Use
proper terminology and write balanced
chemical equations where applicable.
• Justify your answers thoroughly.
Reviewing What You’ve Learned
• How would you determine whether or not a
test tube containing a clear colorless solution
is water or sulfuric acid?
• Given a white powder that my be silver
chloride or sodium chloride.
• Whether a compound is silver nitrate or
sodium nitrate.3