Transcript Chapter 5 A Closer Look at Chemical Equations

Chapter 5
A Closer Look at
Chemical Equations
Single Replacement Equations
Decomposition
Gas-Forming Reactions
Neutralization Reactions
Precipitation Reactions
Oxidation-Reduction Reactions
Single Replacement
Reactions
• Single replacement reactions are reactions
that involve an element replacing one part
of a compound
• The products include the displacement
element and a new compound.
• An element can only replace another
element that is less active than itself.
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General Activity Series
for Metals
•
•
•
•
•
•
•
•
•
•
•
•
Li (most active)
Ca
Na
Mg
Al
Zn
Fe
Pb
[H2]
Cu
Ag
Pt (least active)
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Can metal 1 replace metal 2?
Metal 1
Cu
Metal 2
Al
Cu
Ag
Mg
Na
Pb
Zn
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Yes
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No
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General Activity Series
for Nonmetals
•
•
•
•
F2 (most active)
Cl2
Br2
I2 (least active
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Can nonmetal 1 replace nonmetal 2?
Nonmetal 1
Nonmetal 2
Br
Cl
Cl
F
Br
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Yes
No
I
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• Active metals replace less active metals
from their compounds in aqueous solutions.
• Example: Magnesium turnings are added
to a solution of iron (III) chloride.
3 Mg(s) + 2FeCl3 --- > 3 MgCl2 + 2 Fe(s)
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Can metal 1 replace metal 2 in
compound?
Metal 1
Al
Compound
CuCl2
Pb
AgNO3
Mg
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Yes
No
NaNO3
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• Active metals replace hydrogen in
water
• Example: Sodium is added to water.
2Na(s) + 2 HOH --- > 2NaOH(aq) + H2
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(g)
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List the products
Active
Metal
Li
Water
HOH
Ca
HOH
K
HOH
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Product 1 Product 2
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• Active metals replace hydrogen in
acids.
• Example: Lithium is added to
hydrochloric acid
• 2 Li(s) + 2HCl(aq) --- > H2(g) 2 LiCl(aq)
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List the products
Active
Metal
Zn
Acid
HBr
Ca
HCl
Mg
H2SO4
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Product 1
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Product 2
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Are these reactions possible?
Metal
Al
Acid
HCl
Ag
HCl
Cu
HCl
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Product 1
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Product 2
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• Active nonmetals replace less active
nonmetals from their compounds in
aqueous solutions.
• Example: Chlorine gas is bubbled into
a solution of potassium iodide
• Cl2(g) + 2 KI(aq) --- > I2(g) + 2 KCl(aq)
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• If a less reactive element is
combined with a more reactive
element in compound form, there will
be no resulting reaction.
• Example: Chlorine gas is bubbled into
a solution of potassium fluoride
• Cl2(g) + KF(aq) --- > No reaction
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Are these reactions possible?
Nonmetal 1
Br2
Cl2
I2
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Compound
NaCl
Yes
No
NaI
NaCl
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Homework Wkst.
Activity Series of
Metals
• 1. Cu(s) + H2O(l) --- > No reaction
• 2. Br2(l) + 2NaI(s) --- > I2(g) + 2NaBr
• 3. Al(s) + 3AgNO3(aq) --- > Al(NO3)3 + 3Ag
• 4. Zn(s) + H2SO4(aq) ---- > ZnSO4 + H2(g)
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5.
3F2(g) + 2AlCl3(aq) --- > 3Cl2(g) + 2AlF3
6.
Mg(s) + Pb(C2H3O2)2(aq) -- > Mg(C2H3O2)2 + Pb(s)
7.
I2(g) + NaCl(aq) --- > No reaction
8.
Ca(s) _+ 2HNO2 ---- > Ca(NO2)2 + H2
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9. 2Li(s) + 2 H2O(l) ---- > 2LiOH + H2(g)
10. FeCl3(aq) + Pt(s) ---- > No reaction
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Decomposition
Reactions
P.Q. I – Activity Series
1.
Which is the more active metal?
Ca or Ag
2.
Which is the more active nonmetal?
Cl2 or Br2
3.
Which reaction isn’t possible?
a. 2Na + Cl2 ---- > 2 NaCl
b. KBr + F2 ---- > 2KF + Br2
c. Cu + HCl ---- > CuCl2 + H2
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Decomposition reactions occur when a
single reactant is broken down into
two or more products.
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Metallic Carbonates
• Decompose into metallic oxides and
carbon dioxide
• Example: A sample of magnesium
carbonate is heated.
• MgCO3 --- > MgO + CO2
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Metallic Carbonates
Metallic
Carbonate
K2CO3
Product 1
Product 2
CaCO3
Na2CO3
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Metallic chlorates decompose into
metallic chlorides and oxygen.
Example. A sample of magnesium
chlorate is heated.
Mg(ClO3)2(s) --- > MgCl2(s) + 3 O2(g)
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Ammonium carbonate decomposes into
ammonia, water and carbon dioxide.
(NH4)2CO3--- >2 NH3(g) + H2O(l) + CO2(g)
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Sulfurous acid decomposes into sulfer
dioxide and water.
H2SO3 ---- > H2O + SO2(g)
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Carbonic acid decomposes into carbon
dioxide and water.
H2CO3 ---- > H2O + CO2(g)
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A binary compound may break down to
produce two elements.
2 NaCl ---- > 2 Na + Cl2
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Hydrogen peroxide decomposes into
water and oxygen.
2H2O2 --- > 2 H2O + O2
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Ammonium hydroxide decomposes into
ammonia and water.
NH4OH --- > NH3 + HOH
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Synthesis Reactions
Synthesis reactions occur when two
or more reactants combine to form a
single product.
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A metal combines with a nonmetal to
form a binary salt.
Example: solid sodium oxide is added
to water
6 Li(s) + N2(g) --- > 2Li3N(aq)
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Type 2
Metallic oxides and water form bases
(metallic hydroxides).
Solid sodium oxide is added to water
Na2O(s) + HOH(l) --- > 2 NaOH(aq)
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Type 3
Nonmetallic oxides and water form acids.
The nonmetal retains its oxidation number.
Example: Carbon dioxide is bubbled into
water
CO2(g) + H2O(l) ---- > H2CO3(aq)
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Type 4.
Metallic oxides and nonmetallic oxides
form salts.
Example: Solid sodium oxide is added
to carbon dioxide
Na2O(s) + CO2(g) ---- > Na2CO3(s)
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Homework Wkst.
Synthesis/Decomposition
1.
CaCO3(s) --- > CaO(s) + CO2(g)
2. SO2(g) + H2O(l) -- > H2SO3(aq)
3. K2O(s) + CO2(aq) --- > K2CO3
4. 2H2O2(l) ---- > 2H2O(l) + O2(g)
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5. Li2O(s) + HOH(l) --- > 2 LiOH(aq)
6. 2AlCl3(l) --- > 2 Al(l)
+
3 Cl2(g)
7. 2Na(s) + I2(g) --- > 2NaI(s)
8. H2CO3 (aq) --- > CO2(g) + H2O(l)
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9. 2KClO3(s) --- > 2KCl(s) + 3 O2(g)
10. MgO(s) + SO3(g) --- > MgSO4(s)
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Double Replacement
Reactions
Formation of a Gas
Neutralization Reactions
Formation of a Precipitate
Formation of A Gas
Common Gases in
Metathesis Reactions
Gas 1. H2S
Any sulfide (salt of S2-) plus any acid form
H2S(g) and a salt.
Example: iron (II) sulfide + hydrochloric
acid
FeS(s) + 2HCl(aq) -- > FeCl2(aq) + H2S(g)
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Gas 2. Carbon Dioxide
Any carbonate salt (salt of CO32-) plus
any acid form CO2(g) , H2O and a salt.
Example: Potassium carbonate plus nitric acid
K2CO3(s) + 2 HNO3(aq) -- > CO2(g) + HOH(l) + 2KNO3(aq)
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Gas 3. SO2
Any sulfite (salt of SO32-) plus any acid form SO2(g),
HOH and a salt
Example:
sodium sulfite plus hydrochloric acid
Na2SO3(s) + 2HCl(aq) --- > 2NaCl(aq) + SO2(g) + HOH(l)
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Gas 4. NH3
Any ammonium salt (salt of NH4+) plus any
soluble strong hydroxide react upon
heating to form NH3(g), HOH and a salt.
Example: ammonium chloride plus sodium
hydroxide
NH4Cl(aq) + NaOH -- > NH3(g) + HOH(l) + NaCl(aq)
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Homework Wkst.
Formation of a Gas
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H3C6H5O7 + NaHCO3
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1.
(NH4)2SO4(aq)+2KOH(aq) - > 2NH3(g)+ 2HOH(l) +K2SO4(aq)
2. (NH4)2S(aq) + 2HCl(aq) -- > H2S(g) + 2 NH4Cl(aq)
3. CoCl2(aq) + 2AgNO3(aq) -- > 2AgCl(s) + Co(NO3)2(aq)
4. CaCO3(a) + H2SO4(aq)-- > HOH(l) + CO2(g) + CaSO4(s)
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5.
K2SO3(aq) + 2HBr(aq) -- > HOH(l) + SO2(g) + 2KBr(aq)
6.
K2S(aq) + 2 HNO3(aq) --- > 2KNO3(aq) + H2S(g)
7. 2NH4I(aq +MgSO4(aq)--- > MgI2(aq) + (NH4)2SO4(aq)T
8. Ti(CO3)2(s) + 4HCl(aq)-- >TiCl4(aq) + CO2(g) + HOH (l)
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9.
CaSO4 + 2HC2H3O2(aq) - >Ca(C2H3O2)2 + HOH + SO2(g)
10.
Sr(OH)2(aq) + (NH4)2S(aq) - >SrS(aq) +2 NH3(g)+ 2HOH(l)
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Neutralization
Reactions
Acids & Bases
Strong Acids and Bases
Acids
HClO4
HClO3
HCl
HBr
HI
HNO3
H2SO4
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Bases
Group IA
Ba2+
Ca2+
Sr2+
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• Acids ionize
[(H3O+) = hydronium ion]
HCl(g) + H2O(l) --- > H3O+(aq) + Cl-
or HCl(aq) ----- > H+(aq) + Cl-(aq)
• Bases dissociate
NaOH(s) ----- > Na+(aq) + OH-(aq)
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Electrolytes &
Nonelectrolytes
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Ionization of Monoprotic Acids
(One ionizable hydrogen ion)
HBr
_______________________
HI
_______________________
HNO3
_______________________
HClO4
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_______________________
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Ionization of Polyprotic Acids
(Two or more ionizable hydrogen ions)
H2SO4
__________________________
H3PO4
H2SO3
__________________________
__________________________
__________________________
__________________________
H2C2O4
__________________________
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Acids react with bases to produce salts and
water.
A salt consists of a cation from a base and
an anion from an acid.
H2S(g) + 2KOH(aq) --- > K2S(aq) + 2 HOH(l)
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Later you will see that the
concentration (molarity) will
determine the type of salt formed
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Formation of a
Precipitate
Precipitate
An insoluble substance (solid) formed by the
reaction of two aqueous substance.
Result of ions bonding together so strongly
that the solvent (water) cannot pull them
apart.
Will settle out from the solution
Results in the removal of ions from the
solution
* Nothing is completely insoluble in water
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Ionic Equations
• Soluble salts, strong acids and strong
bases are written as separate ions.
• Insoluble salts, suspensions, solids, weak
acids, weak bases, gases, water and
organic compounds are always written as
individual molecules.
• Solubility rules must be memorized.
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Solubility Rules
Soluble Compound
All Group IA (alkali metal cations)
NH4+
Inorganic acids
Low molecular mass organic acids
All nitrates
All chlorates
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Water Soluble
Soluble Compounds with Exceptions
Exceptions
Acetates
Ag+
Chlorides
Ag+, Hg22+, Pb2+
Bromides
Ag+, Pb2+, Hg22+, Hg+
Iodides
Ag+, Pb2+, Hg22+, Hg+
Sulfates
Ag+, Pb2+, Hg22+,
Ba2+, Ca2+, Sr2+
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Mainly Water Insoluble
Carbonates
Chromates
Hydroxides
Phosphates
Sulfites
Sulfides
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Exceptions
Group IA elements & NH4+
Group IA elements, NH4+,
Ca2+ & Sr2+
Group IA elements, NH4+
Ba 2+ , Ca2+ & Sr2+
Group IA elements & NH4+
Group IA elements & NH4+
Group IA elements & NH4+
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Soluble or Insoluble
Compound
Soluble
Insoluble
Al(OH)3
Mn(NO3)2
PbS
NiCl2
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Writing Ionic Equations
The only common substances that should be
written as ions in ionic equations are:
soluble salts
strong acids
strong bases
weak base – ammonium hydroxide
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• Weak acids and weak bases don’t ionize or
dissociate fully. Write as molecular
compounds.
Examples
Acetic acid
hydrofluoric acid
phosphoric acid
nitrous acid
ammonia
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HC2H3O2
HF
H3PO4
HNO3
NH3
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Dissociating Water Soluble
Compounds
HCl
H+
Cl-
CaCO3
Ca2+
CO32-
(NH4)2S
2 NH4+
S2-
Ba(NO3)2
Ba2+
2 NO3-
K2CrO4
2 K+
CrO42-
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Net Overall Reaction
Step 1: Write molecular equation first.
Products? Reactants exchange
partners.
Step 2: Indicate whether the
substance is aq, s, l or g.
Step 3: Write the ionic equation
Step 4: Write the net overall equation
(all spectators eliminated)
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Example 1
Calcium nitrate and rubidium chloride are mixed together.
1.
Write skeletal equation
2+ + 2+ + CaNO3 + RbCl ---- > CaCl + RbNO3
2.
Balance formulas and determine whether they are soluble or insoluble in
water
Ca(NO3)2(aq) + RbCl(aq) --- > CaCl2(aq) + RbNO3(aq)
3. Balance equation
Ca(NO3)2(aq) + 2RbCl(aq) --- > CaCl2(aq) + 2 RbNO3(aq)
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Ca(NO3)2(aq) + 2RbCl(aq) --- > CaCl2(aq) + 2 RbNO3(aq)
4.
Write Ionic equation:
Ca2+ + 2 NO3- + 2 Rb+ + 2 Cl- --- > Ca2+ + 2 Cl- + 2 Rb+ + 2 NO3
5.
Net Overall Reaction (net ionic equation)
No reaction (everything cancels out)
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Example 2
Calcium hydroxide is added to magnesium chloride
1.
Write skeletal equation
2+ 2+ 2+ 2+ CaOH + MgCl --- > CaCl + MgOH
2.
Balance formulas and determine whether they are soluble or
insoluble in water
Ca(OH)2(aq) + MgCl2(aq) --- > CaCl2(aq) + Mg(OH)2(s)
3. Write balanced equation
Ca(OH)2(aq) + MgCl2(aq) --- > CaCl2(aq) + Mg(OH)2(s)
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Ca(OH)2 + MgCl2 --- > CaCl2 + Mg(OH)2
4.
Write ionic equation.
Ca2+(aq) + 2OH-(aq) + Mg2+(aq) + 2Cl- (aq) --- > Ca2+(aq) + Mg(OH)2(s)
5.
Net Overall equation. (net ionic equaton)
Mg2+(aq) + 2 OH-
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(aq)
---- > Mg(OH)2(s)
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Silver Nitrate & Sodium Chloride
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Homework Wkst.
Formation of a
Precipitate
1. 2AgNO3(aq) + Na2CO3(aq) -- >Ag2CO3(s) +2 NaNO3(aq)
2Ag+(aq) + CO32-(aq) --- > Ag2CO3(s)
2.
2KOH + (NH4)3PO4 --- > 3NH4OH + K3PO4
No reaction
3. NiSO4 + Ba(NO3)2 -- > Ni(NO3)2 + BaSO4(s)
NO Reaction
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4. NaOH + SrCl2 --- > 2NaCl + Sr(OH)2(s)
Sr2+(aq) + 2 OH-(aq) --- > Sr(OH)2(s)
5.
Hg2(NO3)2 + 2NaBr -->2NaNO3(aq) + Hg2Br2(s)
Hg22+(aq) + 2 Br-(aq) --- > Hg2Br2(s)
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6. Ag2SO4(aq)+ MgI2(aq) - > 2AgI(s) +MgSO4(aq)
2Ag+(aq) + 2I-(aq) --- > 2AgI(s)
7. Same as #6
8. BaS(aq) + CuSO4(aq) -- > BaSO4(s) + CuS
Ba2+(aq) + SO42(aq) --- > BaSO4(s)
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(aq)
82
10. (NH4)2SO4 + ZnCl2 --> 2NH4Cl
No Reaction
(aq)
+ ZnS(aq)
11.
AlCl3(aq) + 3NaOH(aq) --- > 3NaCl(aq) +Al(OH)3
Al3+(aq) + 3OH-(aq) -- > Al(OH)3(s)
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12.
3Na2SO4(aq) +2FeBr(3) -- >6NaBr(aq) + Fe2(SO4)3(s)
2 Fe3+(aq) + 3SO42-(aq) --- > 2FeSO4(s)
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Separating Cations
Using Solubility Rules
What reagent solution might you use to separate the cations in
the following mixtures, that is, with one ion appearing in
solution and the other in a precipitate?
Example 1: BaCl2 and CaCl2
barium chloride and calcium chloride
Use a reagent containing the SO42- ion. (Na2SO4)
Barium sulfate would form.
Barium sulfate would precipitate out of solution because it is
insoluble.
Calcium chloride would remain in solution.
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Example 2.
Lead (II) sulfate and copper (II) nitrate
PbSO4
+ Cu(NO3)2
Add water. Copper (II) nitrate will dissolve.
Lead (II) sulfate will settle to the bottom.
It is insoluble in water.
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Homework
Chp. 5: 14 & 15
14.
a)
b)
c)
d)
e)
f)
Pb2+ (aq) + 2Br- (aq) ---- > PbBr2(s)
No reaction occurs
Fe3+ (aq) + 3 OH- (aq) --- > Fe(OH)3(s)
Ca2+(aq) + CO32- (aq) --- > CaCO3(s)
Ba 2+(aq) + SO42- (aq) --- > BaSO4(s)
No reaction
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15.
a)
b)
c)
d)
e)
f)
Neutralization: salt and water produced
No reaction occurs
Gas evolution: gas and salt produced
Gas evolution: gas and water produced
Redox: lost and gain of electrons
No reaction occurs
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Oxidation-Reduction
Reactions
(Redox)
• Electrons are transferred from one
atom to another
• Reaction is known as an oxidationreduction reaction. (Redox for
short)
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Oxidation
Oxidation (oxidize) is the loss of electrons by a substance.
LEO
Lose electrons – oxidation
Oxidation number increases.
Nao --- > Na+ + 1 e
-
Sodium has been oxidized.
It is called the reducing agent. (electron donor)
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Reduction
Reduction involved in the gain of electrons.
GER
Gain electrons – reduction.
Oxidation number decreases.
Clo + 1 e- ---- > Cl-1
Chlorine has been reduced.
It is called the oxidizing agent (electron acceptor).
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Identifying Redox
Reactons
o
o
2+ 22 Mg(s) + O2(g) --- > 2 MgO(s)
Mg has gone from O to 2+. It has lost 2
electrons.
Oxygen has gone from 0 to 2-. It has gained
2 electrons.
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Half-Reactions
• Place the lost electrons on the product
side.
• Place the gained electrons on the reactant
side.
o
2+
Mg --- > Mg + 2eOxidation
o
22O2 + 4e- ---- > 2 OReduction
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Cu + HNO3
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Exercise I:
Identifying Redox
Equations.