Lecture 8

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Lecture 8: Ionic Chemical Reactions (Ch 10)
Suggested HW:
1, 9, 35, 39, 41, 45,
51, 61
1
Recap
• To date, we have learned about:
– Chemical elements and their classifications
– Valence electron configurations and their effect on
chemical properties and reactivity
– Predicting ionic charges
• Now, we will begin to learn about chemical reactions
involving ionic compounds
Recap
Na+
Cl-
• Electrostatic interactions
(coulombic attractions)
between NaCl molecules holds
them together in a lattice
• This is the strongest type of
intermolecular force, which is
why all ionic compounds are
solid at room temperature and
have extremely high melting
points.
About Polyatomic Ions
• Polyatomic ions are covalent molecules that possess
charge and behave as normal ions in solution.
• When a salt containing a polyatomic ion is dissolved in
water, the polyatomic ions themselves DO NOT break
apart. They are simply separated from the counter-ion.
• Example: Phosphate (PO43-)
Na3PO4 (s)
Sodium
Phosphate
H2O (L)
3Na+ (aq) + PO43-(aq)
Sodium
cations
Phosphate
anion
KNOW YOUR POLYATOMIC IONS !!!!!
Charge
-1
Name
Structure
OH-
Carbonate
CO32-
Cyanide
CN-
Oxalate
C2O42-
Bicarbonate
HCO3-
Sulfate
SO42-
Acetate
CH3COO-
Sulfite
SO32-
Nitrate
NO3-
Nitrite
NO2-
Perchlorate
ClO4-
Name
Structure
Hydroxide
Charge
-2
Charge
Charge
-3
Name
Structure
Phosphate
PO43-
+1
Name
Structure
Ammonium NH4+
5
Group Examples
• Write out chemical formulas for the following:
– Calcium nitrate
– Sodium sulfate
– Ammonium hydroxide
– Iron (III) acetate
– Tungsten (VI) cyanide
– Aluminum perchlorate
6
Ionic Reactions
Chemical reactions involving ionic compounds can be
classified as one of the following:
1.
2.
3.
4.
combination reactions
decomposition reactions
single replacement reactions
double replacement reactions
Combination Reactions
• In a combination reaction, multiple reactants combine to
form a single product
– The reaction may occur between two elements
– Or between an element and a compound
– Or between two compounds
3Li(s) + P(g)  Li3P(s)
Ca(s) + Cl2(g)  CaCl2(s)
SO3(g) + H2O(l)  H2SO4(aq)
MgO(s) + CO2(g)  MgCO3 (s)
Decomposition Reactions
2HgO(s)
2Hg(l) + O2(g)
2KClO3(s)
2KCl(s) + 3O2(g)
Single Replacement Reactions
In a single replacement reaction,
 When one metal replaces another, this is also called a
transmetallation reaction.
Zn(s) + 2AgCl (aq)  ZnCl2(aq) + 2Ag(s)
Fe(s) + CuSO4(aq)  FeSO4(aq) + Cu(s)
Single Replacement
• Single replacement reactions occur because one metal is less
stable (more active) in its solid state than the other.
• In the reaction below, Zn displaces Ag because Zn is more
reactive:
Zn(s) + 2AgCl (aq)
ZnCl2(aq) + 2Ag(s)
• More active metals prefer to exist as aqueous ions instead of
neutral atoms
• A more active metal will displace a less active metal. The
opposite will NOT occur. An activity series is provided on pg.
325 of the text.
Group Work
As shown in the table, Li, which has the
highest activity, is the most likely to
react. It will displace any other metal.
Potassium will displace any metal
except Li. And so on.
Predict the products.
Li (s) + Ca(ClO4)2 (aq)
Na (s) + ZnSO4(aq)
K (s) + LiCl (aq)
Activity Series
Single Replacement Reactions Involving Acids
• When a metal reacts with a binary acid (HX), the metal
replaces the hydrogen atom to yield an ionic compound
and hydrogen gas.
Zn(s) + 2HCl (aq)
ZnCl2(aq) + H2(g)
Redox Reactions
• Single replacement reactions are examples of red-ox
(reduction-oxidation) reactions
• A reduction process corresponds to a process in which the
oxidation state (charge) of an element/ion becomes more
negative during the course of a reaction
• In an oxidation process, the oxidation state of an element/ion
becomes more positive during a reaction
Redox Reactions
• Consider the following single replacement reaction:
Zn(s) + Cu SO4 (aq)
On the reactant side,
we have elemental Zn.
The charge on any
pure element is 0
On the reactant side,
we have a Cu2+ ion.
Zn SO4 (aq) + Cu (s)
On the product side,
we have a Zn2+ ion.
Since the charge of Zn
has gone from 0 to 2+,
Zn has undergone an
oxidation. Zn loses 2
electrons. Where did
they go???
On the product side,
we have elemental
Cu, so Cu has
undergone a
reduction from 2+ to
0 by taking electrons
from Zn.
Oxidizing and Reducing Agents
Zn(s) + Cu SO4 (aq)
Zn SO4 (aq) + Cu (s)
• We have identified the reduction and oxidation processes in
the reaction above
Zn0  Zn2+ + 2eRED-OX
REACTIONS
Cu2+ + 2e-  Cu0
• Because Zn gets oxidized, it is the reducing agent. In other
words, the oxidation of Zn causes the reduction of Cu2+
• Because Cu2+ gets reduced, it is the oxidizing agent. Zn is
oxidized because Cu2+ takes electrons away from more active
Zn.
Oxidizing and Reducing Agents
Zn(s) + Cu SO4 (aq)
Zn SO4 (aq) + Cu (s)
(C.I.R.L) Rust Formation
Reduced
4Fe(s) + 3O2(g)
Oxidized
2Fe2O3(s)
Double Replacement Reactions
In a double replacement result,
 two salts react, and the anions exchange places
AgNO3(aq) + NaCl(aq)
AgCl(s) + NaNO3(aq)
ZnS(s) + 2HCl(aq)
ZnCl2(aq) + H2S(g)
Examples
• Balance the following double replacement reactions
A. CaBr2 (aq) + K2CO3(aq)
B. NH4Cl (aq) + MgSO4 (aq)
Most Double Replacement Reactions Yield Precipitates
• An easy way to identify a chemical reaction is if there is a
change in phase.
• In a Precipitation Reaction, an insoluble (solid, does not
dissolve) ionic product is formed.
• In the figure to the left, Na2S
(aq) and Cd(NO3)2 (aq)
undergo double replacement
to form CdS and NaNO3 .
• CdS is insoluble
Net Ionic Equations
• It is proper practice to use NET IONIC EQUATIONS when there is
a change in phase
• Ex. Na2S(aq) + Cd(NO3)2(aq)
2NaNO3(aq) + CdS(s)
• Since we know that ionic solutions dissociate in water, we can
rewrite the equation above in ionic form:
2Na+(aq) + S2-(aq) + Cd2+(aq) + 2NO3-(aq)
CdS(s) + 2Na+(aq) + 2NO3-(aq)
The ions in red undergo a chemical reaction, as indicated by
the change in phase.
The remaining ions are called
SPECTATOR IONS because they are not involved in the
reaction in any way.
Net Ionic Equations
Na+(aq) + S2-(aq) + Cd2+(aq) + NO3-(aq)
CdS(s) + Na+(aq) + NO3-(aq)
• The spectators ions cancel out. The remaining reactants and
products comprise the net ionic equation.
Cd2+(aq) + S2-(aq)
CdS(s)
NET IONIC EQUATION
• In order to write a net ionic equation, you must know which
ionic compounds are insoluble. The solubility rules enable this.
Solubility Rules
1. All group 1 and ammonium salts are soluble!
2. All nitrates, acetates, and perchlorates are soluble
3. With the exception of all anions mentioned in #2, Ag, Pb, and
Hg(I) salts are all insoluble
4. With the exception of those cations mentioned in #1,
carbonates, sulfides, oxides, and phosphates are insoluble
5. With the exception of Ca, Sr, and Ba, and those cations
mentioned in #1, all hydroxides are insoluble
6. All sulfates are soluble EXCEPT for Ca, Sr, Ba, and those
cations mentioned in #3.
Examples
• Use solubility rules to predict the products of the following
double replacement reactions. Write the net ionic
reaction. If there is no reaction, write ‘no reaction’:
MgBr2 (aq) + K2CO3 (aq)
NaCH3COO (aq) + CaBr2 (aq)

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