Writing ionic equations
You have now had some practice at writing basic balanced
equations. These are the language of chemistry and you
need to ensure that you become as fluent as possible at
doing this. Equations tell us exactly what is happening in a
reaction. However, as you will see, they sometimes include
species that do nothing. We can leave these species out to
just show the essentials of what is going on. The result, an
ionic equation (although it doesn’t just contain ions) only
shows species that are fully involved; spectator ions (those
that do nothing) are omitted. Once you are accomplished
with standard balanced equations, these are only a small
step further. Promise!
What does the term species mean here?
In the following, to be sure you look at how EVERYTHING
changes, be sure you write the state symbols (s, l, g or aq
as appropriate).
1. Write a chemical equation for the precipitation reaction
between silver nitrate (AgNO3) solution with sodium chloride
(NaCl) solution, forming silver chloride (as a solid precipitate)
and one other product solution.
What is really happening here? To see this write the equation out
again, breaking everything down into ions. What has actually
changed? How could you summarise this?
2. Write a chemical equation for the displacement reaction
between sodium metal and magnesium sulphate (MgSO4)
solution, forming magnesium metal and one other product.
What is really happening here? Write the equation out again, breaking
everything down into ions. What has actually changed? How could you
summarise this?
Answers to 1 and 2 below
What is really happening when an ionic solid
dissolves?
Watch the demonstration below carefully and listen /
contribute to the discussion that follows.
General rules for writing state symbols
 Solid (s) for all nearly elements (except obvious gases:
O2, H2, N2, F2 and Cl2 and two liquid elements [Br2, Hg])
and for all insoluble compounds. But also (s) for ANY
ionic compound if not dissolved.
 Liquid (l) for Bromine and Mercury (only elements – see
above), plus many short-medium chain length organic
compounds. Including all alcohols, carboxylic acids etc.
 Gas (g) for obvious gaseous elements (see above), plus
very short chain alkanes and alkenes e.g. methane and
other known gases, e.g. carbon dioxide
 Solution (aq) for all acids and all soluble compounds if
used in solution.
General rules for writing state symbols
 Solid (s) for all nearly elements (except obvious gases:
O2, H2, N2, F2 and Cl2 and two liquid elements [Br2, Hg])
and for all insoluble compounds. But also (s) for ANY
ionic compound if not dissolved.
 Liquid (l) for Bromine and Mercury (only elements – see
above), plus many short-medium chain length organic
compounds. Including all alcohols, carboxylic acids etc.
 Gas (g) for obvious gaseous elements (see above), plus
very short chain alkanes and alkenes e.g. methane and
other known gases, e.g. carbon dioxide
 Solution (aq) for all acids and all soluble compounds if
used in solution.
Example 2 final answer:
Writing ionic equations
You have now had some practice at writing basic balanced equations. These are the language at
chemistry and you need to ensure that you become as fluent as possible at doing this. Equations
tell us exactly what is happening in a reaction. However, as you will see, they sometimes include
species that do nothing. We can leave these species out to just show the essentials of what is
going on. The result, an ionic equation (although it doesn’t just contain ions) only shows
species that are fully involved; spectator ions (those that do nothing) are omitted. Once you are
accomplished with standard balanced equations, these are only a small step further. Promise!
All the basic rules that you already know and use for balancing equations
apply here. There are four additional points of emphasis:
1. State symbols are even more important, and must be included.
2. Write down any elements, covalent compounds and insoluble ionic
compounds exactly as before but show soluble ionic compounds as
separated into their constituent ions including the relevant charges1.
3. The overall equation must balance for electronic charge as well as for
numbers and identities of atoms. If you have got the atoms right, the charges
will follow automatically, but you must check.
4. Once you have the full equation balanced, it needs to be simplified: cancel
out everything that appears unchanged on either side (of the arrow) and
write the simplified equation underneath without these SPECTATOR
IONS: this is the ionic equation.
To practice:
all answers hand-written below
[experimental examples below]
1. Magnesium displaces hydrogen from sulphuric acid to leave a solution of magnesium sulphate.
Represent the reaction with a balanced symbol equation and then convert to an ionic equation.
2. If aqueous calcium iodide is mixed with aqueous silver nitrate, a yellow precipitate of insoluble silver
iodide forms and a solution of calcium nitrate is also formed. Represent the reaction with a
balanced symbol equation and then convert to an ionic equation.
3. Aqueous sodium thiosulphate (Na2S2O3; contains ions Na+ and S2O32-) reacts with hydrochloric acid
to form solid sulphur, sulphur dioxide gas, sodium chloride and water. Represent the reaction with a
balanced symbol equation and then convert to an ionic equation.
4. If aqueous potassium iodide is mixed with aqueous lead (II) nitrate, a yellow precipitate of lead (II)
iodide forms (along with one other product, which you should be able to work out – see 2 above)
Represent the reaction with a balanced symbol equation and then convert to an ionic equation.
5. Magnesium displaces copper from copper sulphate solution, forming soluble magnesium sulphate.
Represent the reaction with a balanced symbol equation and then convert to an ionic equation.
6. Sodium freely oxidises in air, forming solid sodium oxide. Write the balanced symbol and then the
ionic equation. What do you notice in this case?
Remember: ionic equations do not just contain ions. They contain everything
that undergoes a change in the reaction. The only species omitted from the
overall equation are the spectator ions: these are superfluous, as they undergo no
change in the reaction.
This reflects the fact that when a when a soluble ionic compound is dissolved, the ions will fully dissociate
(separate) from one another spread through the solvent individually. For example a solution made by
dissolving sodium chloride in water consists of countless Na+ ions each surrounded by water molecules and
Cl- ions similarly surrounded by water molecules. There are no pairings of NaCl swirling through the
solvent together. See DAIGRAMMATIC NOTES SHEET. How do you know if something is
soluble (and so will dissociate) or not? Simple, see the provided SOLUBILITY RULES.
1
Solubility Rules These are greatly simplified (and you will meet some
exceptions and more specific trends else later in the course) but will do for
now:
 all common sodium, potassium and ammonium salts are soluble
 all nitrates are soluble

common chlorides are soluble, except silver chloride. Other halides (fluorides, bromides and
iodides follow the same basic pattern)

common sulfates are soluble, except those of barium and calcium
 common carbonates are insoluble, except those of sodium, potassium and ammonium
 group 1 hydroxides are soluble but those of transition metals are generally not.
 Most transition metal oxides are not soluble
 Common acids are soluble (and dissociate into H+ ions and corresponding anions e.g.
Cl-, SO42- for hydrochloric acid HCl and sulphuric acid H2SO4 respectively)
Solubility Rules These are greatly simplified (and you will meet some
exceptions and more specific trends else later in the course) but will do for
now:
 all common sodium, potassium and ammonium salts are soluble
 all nitrates are soluble

common chlorides are soluble, except silver chloride. Other halides (fluorides, bromides and
iodides follow the same basic pattern)

common sulfates are soluble, except those of barium and calcium
 common carbonates are insoluble, except those of sodium, potassium and ammonium
 group 1 hydroxides are soluble but those of transition metals are generally not.
 Most transition metal oxides are not soluble
 Common acids are soluble (and dissociate into H+ ions and corresponding anions e.g.
Cl-, SO42- for hydrochloric acid HCl and sulphuric acid H2SO4 respectively)
Solubility Rules These are greatly simplified (and you will meet some
exceptions and more specific trends else later in the course) but will do for
now:
 all common sodium, potassium and ammonium salts are soluble
 all nitrates are soluble

common chlorides are soluble, except silver chloride. Other halides (fluorides, bromides and
iodides follow the same basic pattern)

common sulfates are soluble, except those of barium and calcium
 common carbonates are insoluble, except those of sodium, potassium and ammonium
 group 1 hydroxides are soluble but those of transition metals are generally not.
 Most transition metal oxides are not soluble
 Common acids are soluble (and dissociate into H+ ions and corresponding anions e.g.
Cl-, SO42- for hydrochloric acid HCl and sulphuric acid H2SO4 respectively)
(warm-up for 2nd Lesson)
CaCl2
And a couple of easier examples
NaCl
H2
HCl
AlCl3
Solid aluminium metal reacts with
hydrochloric acid solution to form hydrogen
gas and soluble aluminium chloride.
Write a balanced equation and then obtain the ionic
equation from this.
Aluminium + hydrochloric acid
aluminium chloride + hydrogen
Solid sodium metal reacts with calcium
chloride solution to form solid calcium metal
and soluble sodium chloride.
Write a balanced equation and then obtain the ionic
equation from this.
Sodium + calcium chloride
sodium chloride + calcium
(2nd Lesson)
Experimental examples of ionic equations
1. Mix some magnesium metal turnings with hydrochloric acid.
1. Record your observations in as much detail as possible overleaf.
2. Write a fully balanced equation for the reaction, and then convert that into an
ionic equation.
3. Look as closely as possible at your equation. You should be able to see how it
intimately links to your observations from 1. If it doesn’t, either your
observations lack the required detail and / or your equation is wrong!
2. Reacting silver nitrate with soluble sodium halides.
 You will be provided with a dropper bottle of silver nitrate, and containers
1 – 3. These are aqueous solutions of sodium chloride, sodium bromide and
sodium iodide. Your task is to use the silver nitrate to deduce which number
refers to which halide. The basic reaction is:
Silver Nitrate + Sodium halide
Silver halide + Sodium nitrate
Important fact: all silver halides are insoluble
Record your observations below:
Solution
Observation on adding silver nitrate
Identify of original
sodium halide solution
1
2
3
1. Add the silver nitrate to the solutions.
2. Record each observation as you go along.
3. Do a little research to identify each solution from your observation.
CODE BELOW
4. Once you have identified each solution, write a balanced equation for each
reaction, and then derive the respective ionic equation for each.
5. What do you notice about the ionic equation for each case?
6. Use these three ionic equations to explain exactly what a precipitation reaction
is.
7. You have identified each halide solution by looking at the colour of the precipitate
made when silver nitrate solution was added. But those three colours are quite
similar to one another. Explain how you could use ammonia solution as a backup test to confirm your answer in 3: what would you do and what would you see for
each solution? You will need to go an look this up.
Answers:
Question 3
1 = sodium bromide solution
2 = sodium iodide solution
3 = sodium chloride solution
Why does it not matter what the cation
is for any of these solutions?
See student answers below
6 Each ionic equation has the form:
Ag+ (aq) + X- (aq)
AgX (s) where X- is any halide
This shows that a precipitation reaction is simply when solid is
formed from two different solutions. The solid is called a precipitate.
7
Identifying halide ions:
Ion
What happens when silver
nitrate (aq) is added?
What happens to the AgX
precipitate when ammonia
solution is added?
Chloride
White precipitate
obtained
Bromide
Grey-white or creamwhite precipitate
obtained
Yellow precipitate
obtained
Dissolves in both dilute
ammonia and concentrated
ammonia
Only dissolves in concentrated
ammonia (insoluble in dilute
ammonia)
Insoluble (does not dissolve) in
any form of ammonia (dilute or
concentrated)
Iodide