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.
  2. 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?
  1. 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 charges[1].
  3. The overall equation must balance forelectronic 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.

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)And a couple of easier examples

CaCl2NaCl H2HCl 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 solutionfrom your observation.

CODE BELOW

  1. Once you have identified each solution, write a balanced equationfor each reaction, and then derive the respective ionic equationfor each.
  2. What do you notice about the ionic equation for each case?
  3. Use these three ionic equations to explain exactly what a precipitation reaction is.
  4. 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 solutionas a back-up 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 = sodiumiodide solution

3 = sodiumchloridesolution

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 reactionis 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 / Dissolves in both dilute ammonia and concentrated ammonia
Bromide / Grey-white or cream-white precipitate obtained / Only dissolves in concentrated ammonia (insoluble in dilute ammonia)
Iodide / Yellow precipitate obtained / Insoluble (does not dissolve) in any form of ammonia (dilute or concentrated)

[1]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.