Solution Chemistry
Net Ionics
And
RedOx Reactions
Major Difference:
• soluble ionic compounds exist as IONS
Three Types of Chemical Reactions
• Complete Chemical Reaction
• Complete Ionic Reaction
• Net Ionic Reaction
Complete Chemical Reaction:
• shows the chemical process in the non-ionized
form
• uses parenthetical notation to indicate the states
of matter
Examples
1) 2 NaOH (aq) + CuCl2 (aq)  2 NaCl (aq) +
Cu(OH)2 (s)
2) HBr (aq) + NaOH (aq)  NaBr (aq) + H2O (l)
(acid)
(base)
3) CaCO3 (s) +
(carbonate)
2 HCl (aq)  CO2 (g) + H2O (l) + CaCl2 (aq)
(acid)
Complete Ionic Reaction:
• shows aqueous solutions in the ionic form
(THEY ARE BROKEN APART)
▫ Coefficients are distributed to each ion
▫ Ion subscripts become coefficients
• non-soluble substance, non-ionic substances,
gases, liquids and precipitates are written in
standard form (THEY ARE KEPT TOGETHER)
• Ex:
▫ 2 NaOH (aq)  2 Na+ + 2 OH▫ CuCl2 (aq)  Cu+2 + 2 Cl▫ H2O (l)  No change
▫ CO2 (g)  No change
▫ CaCO3 (s)  No Change
Not Aqueous
1) 2 NaOH (aq) + CuCl2 (aq)  2 NaCl (aq) +
Cu(OH)2 (s)
Becomes:
1) 2 Na+
+
2 OH-
1) 2 NaOH (aq) + CuCl2 (aq)  2 NaCl (aq) +
Cu(OH)2 (s)
Becomes:
1) 2 Na+
+
2 OH- + Cu2+ + 2 Cl-
1) 2 NaOH (aq) + CuCl2 (aq)  2 NaCl (aq) +
Cu(OH)2 (s)
Becomes:
1) 2 Na+
Cl-
+
2 OH- + Cu2+ + 2 Cl-  2 Na+ + 2
1) 2 NaOH (aq) + CuCl2 (aq)  2 NaCl (aq) +
Cu(OH)2 (s)
Becomes:
1) 2 Na+
+ 2 OH- + Cu2+ + 2 Cl-  2 Na+ + 2
Cl- + Cu(OH)2
2) HBr (aq) + NaOH (aq)  NaBr (aq) + H2O (l)
Becomes
2) H+ + Br-+ Na+ + OH-  Na+ + Br- + H2O
3) CaCO3 (s) + 2 HCl (aq)  CO2 (g) + H2O
(l) + CaCl2 (aq)
Becomes:
3) CaCO3 + 2 H+ + 2 Cl-  CO2 + H2O + Ca+2 + 2 Cl-
Net Ionic Reactions
• shows the parts of the reaction that actually
undergo change
• ignores ions that remain unchanged on both
sides of the reaction
• called SPECTATOR IONS because they are not
part of the process
• the most simplified equation
Examples
1) 2 Na+
+ 2 OH- + Cu2+ + 2 Cl-  2 Na+ + 2
Cl- + Cu(OH)2
Examples
1) 2 Na+ + 2 OH- + Cu2+ + 2 Cl-  2 Na+ +
2 Cl- + Cu(OH)2
Examples
Net Ionic:
2 OH- + Cu2+  Cu(OH)2
2) H+ + Br-+ Na+ + OH-  Na+ + Br- + H2O
2) H+ + Br-+ Na+ + OH-  Na+ + Br- + H2O
Net Ionic:
• H+ + OH-  H2O
3) CaCO3 + 2 H+ + 2 Cl-  CO2 + H2O + Ca+2 +
2 Cl-
3) CaCO3 + 2 H+ + 2 Cl-  CO2 + H2O + Ca+2 +
2 Cl-
Net Ionic:
CaCO3 + 2 H+  CO2 + H2O + Ca+2
REDUCTION and OXIDATION REACTIONS
Chemical reactions that show a change of charge in the atoms of
the substances.
Reduction: substances that gain electrons (charge is reduced)
Oxidation: substances that lose electrons (charge is increased)
• LEO goes GER
• Loss of Electrons equals Oxidation
• Gain of Electrons equals Reduction
• Oxidizing Agent: the substance that removes electrons (gets
reduced)
• Reducing Agent: the substance that gives away electrons (gets
oxidized)
Examples
• 2 K + Br2  2 KBr
▫
▫
▫
▫
▫
Initial charge of K and Br2 is neutral
K becomes K+1 and Br becomes Br-1
K lost electrons to Br
K is oxidized (reducing agent)
Br is reduced (oxidizing agent)
• N2 + 3 H2  2 NH3
N2 = no charge
H2 = no charge
NH3 = N-3 and H+1
Oxidizing agent = N
Reducing agent = H
• 2 Al + 6 HCl  2 AlCl3 + 3 H2
Reactant Side:
Al = no charge
H = +1
Cl = -1
Oxidized = Al
Reduced = H
Product Side:
Al = +3
H = no charge
Cl = -1
Determining Oxidation Numbers
1. Free elements are assigned an oxidation state of
zero.
• Ex. Metals (K, Fe, Al), Diatomics
2. The sum of the oxidation states of all that atoms
in a species must be equal to the net charge on
the species.
• Ex. ClO3-  Cl+5 + (O-2)3 = -1
3. The alkali metals (Li, Na, K, Rb, and Cs) in
compounds are always assigned an oxidation
state of +1.
• Ex. KCl, HBr
4. Fluorine in compounds is always assigned an
oxidation state of -1.
• Ex. HF
5. The alkaline earth metals (Be, Mg, Ca, Sr, Ba,
and Ra) and also Zn and Cd in compounds are
always assigned an oxidation state of +2.
• Ex. CaO
6. Hydrogen in compounds with other non-metals
that are more electronegative is assigned an
oxidation state of +1.
• Ex. HCl, NH3
7. Hydrogen in compounds with metals is
assigned an oxidation number of -1.
• Ex. KH
8. Oxygen in compounds is assigned an oxidation
state of -2.
• Ex. H2O, CaO
9. Halogen in compounds is assigned an oxidation
state of -1.
• Ex. HF
Examples
1) PO4 -3
2) CaF2
3) Zn
4) CH4
5) ClO4-1
6) ClO2-1
7) O2
8) N2O3
• What is an activity series, and how is it
used?
•
From: http://antoine.frostburg.edu/chem/senese/101/redox/faq/activity-series.shtml
• An activity series is a list of substances ranked
in order of relative reactivity. For example,
magnesium metal can knock hydrogen ions out
of solution, so it is considered more reactive
than elemental hydrogen:
• Mg(s) + 2 H+(aq)  H2(g) + Mg2+(aq)
• Zinc can also displace hydrogen ions from
solution:
• Zn(s) + 2 H+(aq)  H2(g) + Zn2+(aq)
• so zinc is also more active than hydrogen. But
magnesium metal can remove zinc ions from
solution:
• Mg(s) + Zn2+(aq)  Zn(s) + Mg2+(aq)
The most active metals are at the top of the table; the least active
are at the bottom.
• For a reaction to occur the more active element
will reduce the other element. This usually
occurs when the more active element is in the
neutral (elemental) state and the less reactive
element is an ion.
• Each metal is able to displace the elements
below it from solution
• OR, each metal can reduce the cations of metals
below it to their elemental forms.
• Predicting Products Using the Activity Series:
Ex:
zinc metal in a copper(II) sulfate solution
copper placed into a zinc sulfate solution