# 09. Net Ionic Equations tutorial.doc by rtu13707

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```									09. Net Ionic Equations tutorial.doc

Introduction to Net Ionic Equations
This tutorial will give you an algorithm for writing net ionic equations for aqueous reactions in
general chemistry. Knowledge of the solubility rules is necessary to complete this tutorial. A list of
solubility rules is provided at the end of the tutorial for review. Net ionic equations are necessary in
chemistry for several reasons:
1. They show only those species undergoing chemical change – “spectator” ions are removed.
2. They must be used in any equilibrium calculation.
3. They allow chemical reactions to be combined properly.
4. Used when balancing half-reactions in oxidation-reduction chemistry.
To take a chemical reaction in molecular form and rewrite as a net ionic equation is straight forward
if you follow the steps given below. We will cover each step in detail using a common gas forming
reaction from chemistry 1A.

1. Write the correct reactant and products – DO NOT BALANCE!
The first step is to write the correct chemical formulas for the reactants and products, DO NOT
BALANCE. The reaction will be balanced once the net-ionic equation is complete. As an example
reaction, we will look at the aqueous reaction of sodium carbonate with acetic acid.

Na2CO3 + HC2H3O2 ! NaC2H3O2 + H2O + CO2                                 (skeleton equation)

The above reaction is called the skeleton equation; it contains only the correct chemical formulas of
each reactant and product. Do not balance this equation!

2. Add phase symbols to EVERY reactant and product.
This next step is crucial! You need to add the correct phase symbols to each reactant and product.
You must rely on the solubility rules to assign the phase symbols. The phase symbols are
• (s) solid - for insoluble salts. Check the solubility rules!
• (l) liquid - for pure liquids like water.
• (g) gas – for gaseous reactants and products.
• (aq) aqueous – for any compound, ionic or covalent, that is soluble in water. Check the
solubility rules!

For the above reaction lets assign phases for each compound.
1. Na2CO3: sodium salt. All sodium salts are soluble. Phase is (aq).
2. HC2H3O2: acetic acid. Low molar mass organic acid. Should be soluble. Phase is (aq).
3. NaC2H3O2: sodium salt. All sodium salts are soluble. Phase is (aq).
4. H2O: water as a product or reactant. We will treat it as pure. Phase is (l).
5. CO2: Low molar mass covalent compound. Gas at room temperature. Phase is (g).

Rewrite the equation with the phase symbols:

Na2CO3(aq) + HC2H3O2(aq) ! NaC2H3O2(aq) + H2O(l) + CO2(g) (molecular equation)

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09. Net Ionic Equations tutorial.doc

The above reaction is called the molecular equation. All reactant and products are present with the
correct phases.

3. Find the STRONG ELECTROLYTES in the molecular equation
In this step you highlight only the strong electrolytes in the molecular equation. Only the strong
electrolytes will be rewritten as ions in solution. Weak or non-electrolytes will be ignored and not
written as ions. To find the strong electrolytes LOOK ONLY at (AQ) REACTANTS and
PRODUCTS. Why? To be a strong electrolyte a compound must be soluble! What (aq) species are
strong electrolytes in water? ONLY TWO: 1) SOLUBLE SALTS (including the strong bases) AND
2) STRONG ACIDS.

The strong acids are: HCl(aq), HBr(aq), HI(aq), HNO3(aq), HClO4(aq), HClO3(aq), H2SO4(aq)
Note: for H2SO4(aq) write as H+(aq) and HSO4–(aq) NOT H+(aq) and SO42–(aq)

STRONG ELECTROLYTES

Phase is (aq)?

Yes                            No
Possible Strong Electrolyte     NOT a Strong Electrolyte

Is Chemical a Soluble Salt?

Yes                               No
STRONG ELECTROLYTE             Is Chemical a Strong Acid?

Yes                              No
STRONG ELECTROLYTE             NOT a Strong Electrolyte

Lets apply the flowchart to each of our aqueous chemicals in our example reaction.
1. Na2CO3(aq): Soluble salt = STRONG ELECTROLYTE
2. HC2H3O2(aq): Weak acid = NOT a STRONG ELECTROLYTE
3. NaC2H3O2(aq): Soluble salt = STRONG ELECTROLYTE

We have two strong electrolytes in the molecular equation, Na2CO3(aq) and NaC2H3O2(aq).

4. Write the STRONG ELECTROLYTES as ions
Now we rewrite the strong electrolytes as ions in solution. Each ion retains the (aq) phase. All other
species are written as they were! When writing the strong electrolytes as ions, DO NOT ADD A
COEFFIENT FOR THE NUMBER OF IONS FROM THE COMPOUND, ONLY ONE IS
NEEDED FOR EACH ION. We will balance latter.

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09. Net Ionic Equations tutorial.doc

Na+(aq) + CO32–(aq) + HC2H3O2(aq) ! Na+(aq) + C2H3O2–(aq) + H2O(l) + CO2(g)
(ionic equation)

This is termed the ionic equation, each strong electrolyte is written as its constituent ions.

5. Cancel “spectator ions”
We now remove any ions from the equation that appear on both sides of the equation. These are
called “spectator ions”. They are not involved in the chemistry and simply stay dissolved in solution
while the chemistry takes place. For our example equation the only spectator ion is sodium. It has
been removed.

CO32–(aq) + HC2H3O2(aq) ! C2H3O2–(aq) + H2O(l) + CO2(g)                     (net-ionic equation)

This equation is the unbalanced net-ionic equation! Only the reactive species are shown in their
proper physical state: ion, solid, liquid, gas, or (aq) for soluble weak and non electrolytes.

6. Balance the net-ionic equation
The last step is to balance the net-ionic equation. When balancing, make sure BOTH MASS AND
CHARGE are balanced.

CO32–(aq) + 2 HC2H3O2(aq) ! 2 C2H3O2–(aq) + H2O(l) + CO2(g)
(balanced net-ionic equation)

Other Examples
We will look at a few more examples. In each case we write the reaction (skeleton), add phase
symbols (molecular), highlight the strong electrolytes (blue), write the ionic equation, then the net-
ionic and finally the balanced net-ionic equation. Make sure you understand how each step was
completed before moving on!

1.   Na3PO4 + Cu(NO3)2 ! NaNO3 + Cu3(PO4)2                                      (skeleton)
2.   Na3PO4(aq) + Cu(NO3)2(aq) ! NaNO3(aq) + Cu3(PO4)2(s)                     (molecular)
3.   Na3PO4(aq) + Cu(NO3)2(aq) ! NaNO3(aq) + Cu3(PO4)2(s)            (strong electrolytes)
4.   Na (aq) + PO4 (aq) + Cu (aq) + NO3 (aq) ! Na (aq) + NO3 (aq) + Cu3(PO4)2(s) (ionic)
+         3+         2+          –       +          –

5.   PO43+(aq) + Cu2+(aq) ! Cu3(PO4)2(s)                                       (net-ionic)
6.   2 PO4 (aq) + 3 Cu (aq) ! Cu3(PO4)2(s)
3+           2+
(balanced net-ionic)

1.   Al + H2SO4! Al2(SO4)3 + H2                                                     (skeleton)
2.   Al(s) + H2SO4(aq) ! Al2(SO4)3(aq) + H2(g)                                    (molecular)
3.   Al(s) + H2SO4(aq) ! Al2(SO4)3(aq) + H2(g)                           (strong electrolytes)
4.   Al(s) + H (aq) + HSO4 (aq) ! Al (aq) + SO4 (aq) + H2(g)
+            –         3+         2–
(ionic)
5.   Al(s) + H (aq) + HSO4 (aq) ! Al (aq) + SO4 (aq) + H2(g) (net-ionic – no spectators!)
+            –         3+         2–

6.   2 Al(s) + 3 H+(aq) + 3 HSO4–(aq) ! 2 Al3+(aq) + 3 SO42–(aq) + 3 H2(g)(balanced net-ionic)

Daley                                                 3                                            10/9/09
09. Net Ionic Equations tutorial.doc

Summary
In this tutorial we went over the 6 steps necessary to write a net ionic equation for any aqueous
reaction. You must practice and pay close attention to the solubility rules to gain mastery of this
subject. When balancing, make sure the charge is balanced as well!

Self Test
Write net ionic equations for the following reactions. Answers are on the following page.

A. Cr2(SO4)3 + (NH4)2CO3 ! Cr(CO3)2 + (NH4)2 SO4

B. NH3 + HCl ! NH4Cl

C. NaOH + H2SO4 ! Na2SO4 + H2O

D. CaCl2 + NaHCO3 ! CaCO3 +NaCl + HCl

E. FeO + HNO3 ! Fe(NO3)2 + H2O

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09. Net Ionic Equations tutorial.doc

A. Cr2(SO4)3 + (NH4)2CO3 ! CrCO3 + (NH4)2 SO4

1.   Cr2(SO4)3 + (NH4)2CO3 ! Cr2(CO3)3 + (NH4)2SO4
2.   Cr2(SO4)3(aq) + (NH4)2CO3(aq) ! Cr2(CO3)3(s) + (NH4)2 SO4(aq)
3.   Cr2(SO4)3(aq) + (NH4)2CO3(aq) ! Cr2(CO3)3(s) + (NH4)2 SO4(aq)
4.   Cr 3+(aq) + SO42–(aq) + NH4+(aq) + CO32–(aq) ! Cr2(CO3)3(s) + NH4+(aq) + SO42–(aq)
5.   Cr 3+(aq) + CO32–(aq) ! Cr2(CO3)3(s)
6.   2 Cr 3+(aq) + 3 CO32–(aq) ! Cr2(CO3)3(s)

B. NH3 + HCl ! NH4Cl

1.   NH3 + HCl ! NH4Cl
2.   NH3(aq) + HCl(aq) ! NH4Cl(aq)
3.   NH3(aq) + HCl(aq) ! NH4Cl(aq)
4.   NH3(aq) + H+(aq) + Cl–(aq) ! NH4+(aq) + Cl– (aq)
5.   NH3(aq) + H+(aq) ! NH4+(aq)
6.   NH3(aq) + H+(aq) ! NH4+(aq)

C. NaOH + H2SO4 ! Na2SO4 + H2O

1.   NaOH + H2SO4 ! Na2SO4 + H2O
2.   NaOH(aq) + H2SO4(aq) ! Na2SO4(aq) + H2O(l)
3.   NaOH(aq) + H2SO4(aq) ! Na2SO4(aq) + H2O(l)
4.   Na+(aq) + OH–(aq) + H+(aq) + HSO4–(aq) ! Na+(aq) + SO42–(aq) + H2O(l)
5.   OH–(aq) + H+(aq) + HSO4–(aq) ! SO42–(aq) + H2O(l)
6.   2 OH–(aq) + H+(aq) + HSO4–(aq) ! SO42–(aq) + 2 H2O(l)

D. CaCl2 + NaHCO3 ! CaCO3 +NaCl + HCl

1.   CaCl2 + NaHCO3 ! CaCO3 +NaCl + HCl
2.   CaCl2(aq) + NaHCO3(aq) ! CaCO3(s) +NaCl(aq) + HCl(aq)
3.   CaCl2(aq) + NaHCO3(aq) ! CaCO3(s) +NaCl(aq) + HCl(aq)
4.   Ca2+(aq) + Cl–(aq) + Na+(aq) + HCO3–(aq) ! CaCO3(s) + Na+(aq)+ Cl–(aq) + H+(aq)
5.   Ca2+(aq) + HCO3–(aq) ! CaCO3(s) + H+(aq)
6.   Ca2+(aq) + HCO3–(aq) ! CaCO3(s) + H+(aq)

E. FeO + HNO3 ! Fe(NO3)2 + H2O

1.   FeO + HNO3 ! Fe(NO3)2 + H2O
2.   FeO(s) + HNO3(aq) ! Fe(NO3)2(aq) + H2O(l)
3.   FeO(s) + HNO3(aq) ! Fe(NO3)2(aq) + H2O(l)
4.   FeO(s) + H+(aq) + NO3–(aq) ! Fe2+(aq) + NO3– (aq) + H2O(l)
5.   FeO(s) + H+(aq) ! Fe2+(aq) + H2O(l)
6.   FeO(s) + 2 H+(aq) ! Fe2+(aq) + H2O(l)

Daley                                             5                                         10/9/09
Sol Rules.doc

1A                                                                                                                             8A
1                                                                                                                              2
1        H                                                                                                                              He
1.008 2A                                                                             3A       4A       5A       6A       7A    4.003
3     4                                                                           5        6        7        8        9        10
2        Li    Be                                                                          B        C        N        O        F        Ne
6.941 9.012                                                                       10.81 12.01 14.01 16.00 19.00 20.18
11     12                                                                         13     14    15    16    17    18
3        Na   Mg                                                                           Al       Si       P        S        Cl       Ar
22.99 24.31 3B       4B     5B     6B     7B           8B            1B     2B    26.98 28.09 30.97 32.07 35.45 39.95
19     20    21     22     23     24     25     26    27      28    29     30     31     32    33    34    35    36
4        K     Ca   Sc     Ti    V      Cr     Mn     Fe    Co      Ni    Cu     Zn        Ga       Ge       As       Se       Br       Kr
39.10 40.08 44.96 47.90 50.94 52.00 54.94 55.85 58.93 58.70 63.55 65.39 69.72 72.59 74.92 78.96 79.90 83.80
37     38    39    40    41    42    43    44    45    46    47    48    49    50    51    52    53    54
5        Rb    Sr    Y     Zr    Nb     Mo     Tc     Ru    Rh      Pd    Ag     Cd        In       Sn       Sb       Te        I       Xe
85.47 87.62 88.91 91.22 92.91 95.94 (98)        101.1 102.9 106.4 107.9 112.4 114.8 118.7 121.8 127.6 126.9 131.3
55     56    57    72    73    74    75         76     77    78    79    80    81    82    83    84    85    86
6        Cs    Ba   La     Hf    Ta     W      Re     Os     Ir     Pt    Au     Hg        Tl       Pb       Bi       Po       At       Rn
132.9 137.3 138.9 178.5 180.9 183.9 186.2 190.2 192.2 195.1 197.0 200.6 204.4 207.2 209.0 (209) (210)                         (222)
87     88    89    104   105   106   107   108   109   110   111
7        Fr    Ra   Ac     Rf    Ha     Sg     Ns     Hs    Mt
(223) (226)    (227) (261) (262) (266) (262)    (265) (266)

58     59     60     61     62    63      64    65     66     67       68       69       70       71
6    Ce     Pr     Nd     Pm     Sm    Eu      Gd    Tb     Dy        Ho       Er    Tm          Yb       Lu
140.1 140.9 144.2 (145) 150.4 152.0 157.3 158.9 162.5 164.9 167.3 168.9 173.0 175.0
90     91    92    93    94    95    96    97    98    99    100   101   102   103
7    Th     Pa     U      Np     Pu   Am      Cm     Bk     Cf        Es    Fm       Md          No    Lw
232.0 (231) 238.0 (244)    (242) (243) (247) (247) (251)     (252) (257) (258) (259) (260)

Selected solubility rules for ionic compounds in water to REMEMBER:
(The text has a more complete list, but some of these are not listed in the text.)

1. Soluble with no exceptions: Compounds containing a Group 1A metal cation, the
ammonium ion (NH4+) or the nitrate ion (NO3–).

2.       Soluble with some exceptions:
a) Compounds containing the chloride, bromide or iodide ion (Cl–, Br–, or I–). Insoluble
exceptions include the Ag+ and Pb2+ halides.
b) Compounds containing the sulfate ion (SO42–). Insoluble exceptions include BaSO4, PbSO4,
and Ag 2SO4.
c) Compounds containing the acetate ion (C2H3O2–). An insoluble exception is AgC2H3O2.

3.       Insoluble with some exceptions:
a) Compounds containing the carbonate ion (CO32–). Soluble exceptions are Group 1A and
NH4+ carbonates.
b) Compounds containing the phosphate ion (PO43–). Soluble exceptions are Group 1A and
NH4+ phosphates.
c) Compounds containing the hydroxide ion (OH–). Some soluble exceptions are Group 1A
hydroxides and Ba(OH)2. (Note: NH4OH is also soluble, but this salt cannot be isolated in
pure form. NH4OH only exists in aqueous form as NH4+(aq) and OH–(aq) in small amounts in
aqueous ammonia solutions.)
d) Compounds containing the sulfide ion (S2–). Some soluble exceptions are Group 1A sulfides
and the sulfides of Ba2+, Sr2+, Ca2+, and NH4+.
e) Compounds containing the oxide ion (O2–). Soluble exceptions are Group 1A oxides.